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Apollo 11 | Citations | Citations
In some of the following sources, times are shown in the format hours:minutes:seconds (e.g. 109:24:15), referring to the mission's Ground Elapsed Time (GET), based on the official launch time of July 16, 1969, 13:32:00 UTC (000:00:00 GET). |
Apollo 11 | Sources | Sources
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Apollo 11 | External links | External links
"Apollo 11 transcripts" at Spacelog
Apollo 11 in real time
Apollo 11 Press Conference filmed by KPRC-TV at Texas Archive of the Moving Image
Apollo 11 and 13 Checklists at The Museum of Flight Digital Collections.
Apollo 11, 12, and 14 Traverses, at the Lunar and Planetary Institute |
Apollo 11 | Multimedia | Multimedia
Remastered videos of the original landing.
Dynamic timeline of lunar excursion. Lunar Reconnaissance Orbiter Camera
The Eagle Has Landed: The Flight of Apollo 11 (1969) (transcript) from US National Archives (via YouTube)
Apollo 11 Restored EVA Part 1 (1hour of restored footage)
Apollo 11: As They Photographed It (Augmented Reality)—The New York Times, Interactive, July 18, 2019
"Coverage of the Flight of Apollo 11" provided by Todd Kosovich for RadioTapes.com. Radio station recordings (airchecks) covering the flight of Apollo 11.
Category:1969 on the Moon
Category:Buzz Aldrin
Category:Apollo program missions
Category:Neil Armstrong
Category:Articles containing video clips
Category:Michael Collins (astronaut)
Category:Crewed missions to the Moon
Category:Soft landings on the Moon
Category:Spacecraft launched by Saturn rockets
Category:Successful space missions |
Apollo 11 | Table of Content | Short description, Background, Personnel, Prime crew, Backup crew, Support crew, Capsule communicators, Flight directors, Other key personnel, Preparations, Insignia, Call signs, Mementos, Site selection, First-step decision, Pre-launch, Mission, Launch and flight to lunar orbit, Lunar descent, Landing, Lunar surface operations, Lunar ascent, ''Columbia'' in lunar orbit, Return, Splashdown and quarantine, Celebrations, Legacy, Cultural significance, Spacecraft, Moon rocks, Experiment results, Moonwalk camera, Lunar Module ''Eagle'' memorabilia, Anniversary events, <span id="40th anniversary events"></span>40th anniversary, 50th anniversary, Films and documentaries, See also, References, Notes, Citations, Sources, External links, Multimedia |
Apollo 8 | Short description | Apollo 8 (December 21–27, 1968) was the first crewed spacecraft to leave Earth's gravitational sphere of influence, and the first human spaceflight to reach the Moon. The crew orbited the Moon ten times without landing and then returned to Earth. The three astronauts—Frank Borman, James Lovell, and William Anders—were the first humans to see and photograph the far side of the Moon and an Earthrise.
Apollo 8 launched on December 21, 1968, and was the second crewed spaceflight mission flown in the United States Apollo space program (the first, Apollo7, stayed in Earth orbit). Apollo8 was the third flight and the first crewed launch of the Saturn V rocket. It was the first human spaceflight from the Kennedy Space Center, adjacent to Cape Kennedy Air Force Station in Florida.
Originally planned as the second crewed Apollo Lunar Module and command module test, to be flown in an elliptical medium Earth orbit in early 1969, the mission profile was changed in August 1968 to a more ambitious command-module-only lunar orbital flight to be flown in December, as the lunar module was not yet ready to make its first flight. Astronaut Jim McDivitt's crew, who were training to fly the first lunar module flight in low Earth orbit, became the crew for the Apollo9 mission, and Borman's crew were moved to the Apollo8 mission. This left Borman's crew with two to three months' less training and preparation time than originally planned, and replaced the planned lunar module training with translunar navigation training.
Apollo 8 took 68 hours to travel to the Moon. The crew orbited the Moon ten times over the course of twenty hours, during which they made a Christmas Eve television broadcast where they read the first ten verses from the Book of Genesis. At the time, the broadcast was the most watched TV program ever. Apollo8's successful mission paved the way for Apollo 10 and, with Apollo11 in July 1969, the fulfillment of U.S. president John F. Kennedy's goal of landing a man on the Moon before the end of the decade. The Apollo8 astronauts returned to Earth on December 27, 1968, when their spacecraft splashed down in the northern Pacific Ocean. The crew members were named Time magazine's "Men of the Year" for 1968 upon their return. |
Apollo 8 | Background | Background
In the late 1950s and early 1960s, the United States was engaged in the Cold War, a geopolitical rivalry with the Soviet Union. On October 4, 1957, the Soviet Union launched Sputnik 1, the first artificial satellite. This unexpected success stoked fears and imaginations around the world. It not only demonstrated that the Soviet Union had the capability to deliver nuclear weapons over intercontinental distances, it challenged American claims of military, economic, and technological superiority. The launch precipitated the Sputnik crisis and triggered the Space Race.
President John F. Kennedy believed that not only was it in the national interest of the United States to be superior to other nations, but that the perception of American power was at least as important as the actuality. It was therefore intolerable to him for the Soviet Union to be more advanced in the field of space exploration. He was determined that the United States should compete, and sought a challenge that maximized its chances of winning.
The Soviet Union had heavier-lifting carrier rockets, which meant Kennedy needed to choose a goal that was beyond the capacity of the existing generation of rocketry, one where the US and Soviet Union would be starting from a position of equality—something spectacular, even if it could not be justified on military, economic, or scientific grounds. After consulting with his experts and advisors, he chose such a project: to land a man on the Moon and return him to the Earth. This project already had a name: Project Apollo.
An early and crucial decision was the adoption of lunar orbit rendezvous, under which a specialized spacecraft would land on the lunar surface. The Apollo spacecraft therefore had three primary components: a command module (CM) with a cabin for the three astronauts, and the only part that would return to Earth; a service module (SM) to provide the command module with propulsion, electrical power, oxygen, and water; and a two-stage lunar module (LM), which comprised a descent stage for landing on the Moon and an ascent stage to return the astronauts to lunar orbit. This configuration could be launched by the Saturn V rocket that was then under development. |
Apollo 8 | Framework | Framework |
Apollo 8 | Prime crew | Prime crew
The initial crew assignment of Frank Borman as Commander, Michael Collins as Command Module Pilot (CMP) and William Anders as Lunar Module Pilot (LMP) for the third crewed Apollo flight was officially announced on November 20, 1967. Collins was replaced by Jim Lovell in July 1968, after suffering a cervical disc herniation that required surgery to repair. This crew was unique among pre-Space Shuttle era missions in that the commander was not the most experienced member of the crew: Lovell had flown twice before, on Gemini VII and Gemini XII. This would also be the first case of a commander of a previous mission (Lovell, Gemini XII) flying as a non-commander. This was also the first mission to reunite crewmates from a previous mission (Lovell and Borman, Gemini VII).
As of June 2024, James Lovell is the last surviving Apollo 8 astronaut. Frank Borman and William Anders died on November 7, 2023, and on June 7, 2024, respectively. |
Apollo 8 | Backup crew | Backup crew
The backup crew assignment of Neil Armstrong as Commander, Lovell as CMP, and Buzz Aldrin as LMP for the third crewed Apollo flight was officially announced at the same time as the prime crew. When Lovell was reassigned to the prime crew, Aldrin was moved to CMP, and Fred Haise was brought in as backup LMP. Armstrong would later command Apollo11, with Aldrin as LMP and Collins as CMP. Haise served on the backup crew of Apollo11 as LMP and flew on Apollo13 as LMP. |
Apollo 8 | Support personnel | Support personnel
During Projects Mercury and Gemini, each mission had a prime and a backup crew. For Apollo, a third crew of astronauts was added, known as the support crew. The support crew maintained the flight plan, checklists, and mission ground rules, and ensured that the prime and backup crews were apprised of any changes. The support crew developed procedures in the simulators, especially those for emergency situations, so that the prime and backup crews could practice and master them in their simulator training. For Apollo8, the support crew consisted of Ken Mattingly, Vance Brand, and Gerald Carr.
The capsule communicator (CAPCOM) was an astronaut at the Mission Control Center in Houston, Texas, who was the only person who communicated directly with the flight crew. For Apollo8, the CAPCOMs were Michael Collins, Gerald Carr, Ken Mattingly, Neil Armstrong, Buzz Aldrin, Vance Brand, and Fred Haise.
The mission control teams rotated in three shifts, each led by a flight director. The directors for Apollo8 were Clifford E. Charlesworth (Green team), Glynn Lunney (Black team), and Milton Windler (Maroon team). |
Apollo 8 | Mission insignia and callsign | Mission insignia and callsign
thumb|Apollo 8 space-flown silver Robbins medallion
The triangular shape of the insignia refers to the shape of the Apollo CM. It shows a red figure8 looping around the Earth and Moon to reflect both the mission number and the circumlunar nature of the mission. On the bottom of the8 are the names of the three astronauts. The initial design of the insignia was developed by Jim Lovell, who reportedly sketched it while riding in the back seat of a T-38 flight from California to Houston shortly after learning of Apollo8's re-designation as a lunar-orbital mission.
The crew wanted to name their spacecraft, but NASA did not allow it. The crew would have likely chosen Columbiad, the name of the giant cannon that launches a space vehicle in Jules Verne's 1865 novel From the Earth to the Moon. The Apollo11 CM was named Columbia in part for that reason. |
Apollo 8 | Preparations | Preparations |
Apollo 8 | Mission schedule | Mission schedule
On September 20, 1967, NASA adopted a seven-step plan for Apollo missions, with the final step being a Moon landing. Apollo4 and Apollo6 were "A" missions, tests of the SaturnV launch vehicle using an uncrewed Block I production model of the command and service module (CSM) in Earth orbit. Apollo5 was a "B" mission, a test of the LM in Earth orbit. Apollo7, scheduled for October 1968, would be a "C" mission, a crewed Earth-orbit flight of the CSM. Further missions depended on the readiness of the LM. It had been decided as early as May 1967 that there would be at least four additional missions. Apollo8 was planned as the "D" mission, a test of the LM in a low Earth orbit in December 1968 by James McDivitt, David Scott, and Russell Schweickart, while Borman's crew would fly the "E" mission, a more rigorous LM test in an elliptical medium Earth orbit as Apollo9, in early 1969. The "F" Mission would test the CSM and LM in lunar orbit, and the "G" mission would be the finale, the Moon landing.
thumb|left|upright|The first stage of AS-503 being erected in the Vehicle Assembly Building (VAB) on February 1, 1968
Production of the LM fell behind schedule, and when Apollo8's LM-3 arrived at the Kennedy Space Center (KSC) in June 1968, more than a hundred significant defects were discovered, leading Bob Gilruth, the director of the Manned Spacecraft Center (MSC), and others to conclude that there was no prospect of LM-3 being ready to fly in 1968. Indeed, it was possible that delivery would slip to February or March 1969. Following the original seven-step plan would have meant delaying the "D" and subsequent missions, and endangering the program's goal of a lunar landing before the end of 1969. George Low, the Manager of the Apollo Spacecraft Program Office, proposed a solution in August 1968 to keep the program on track despite the LM delay. Since the next CSM (designated as "CSM-103") would be ready three months before LM-3, a CSM-only mission could be flown in December 1968. Instead of repeating the "C" mission flight of Apollo7, this CSM could be sent all the way to the Moon, with the possibility of entering a lunar orbit and returning to Earth. The new mission would also allow NASA to test lunar landing procedures that would otherwise have had to wait until Apollo10, the scheduled "F" mission. This also meant that the medium Earth orbit "E" mission could be dispensed with. The net result was that only the "D" mission had to be delayed, and the plan for lunar landing in mid-1969 could remain on timeline.
On August 9, 1968, Low discussed the idea with Gilruth, Flight Director Chris Kraft, and the Director of Flight Crew Operations, Donald Slayton. They then flew to the Marshall Space Flight Center (MSFC) in Huntsville, Alabama, where they met with KSC Director Kurt Debus, Apollo Program Director Samuel C. Phillips, Rocco Petrone, and Wernher von Braun. Jerry Wittenstein, deputy chief of flight mechanics, presented trajectories for the new mission.family history Kraft considered the proposal feasible from a flight control standpoint; Debus and Petrone agreed that the next Saturn V, AS-503, could be made ready by December 1; and von Braun was confident the pogo oscillation problems that had afflicted Apollo6 had been fixed. Almost every senior manager at NASA agreed with this new mission, citing confidence in both the hardware and the personnel, along with the potential for a circumlunar flight providing a significant morale boost. The only person who needed some convincing was James E. Webb, the NASA administrator. Backed by the full support of his agency, Webb authorized the mission. Apollo8 was officially changed from a "D" mission to a "C-Prime" lunar-orbit mission.
With the change in mission for Apollo 8, Slayton asked McDivitt if he still wanted to fly it. McDivitt turned it down; his crew had spent a great deal of time preparing to test the LM, and that was what he still wanted to do. Slayton then decided to swap the prime and backup crews of the Dand Emissions. This swap also meant a swap of spacecraft, requiring Borman's crew to use CSM-103, while McDivitt's crew would use CSM-104, since CM-104 could not be made ready by December. David Scott was not happy about giving up CM-103, the testing of which he had closely supervised, for CM-104, although the two were almost identical, and Anders was less than enthusiastic about being an LMP on a flight with no LM. Instead, Apollo8 would carry the LM test article, a boilerplate model that would simulate the correct weight and balance of LM-3.
Added pressure on the Apollo program to make its 1969 landing goal was provided by the Soviet Union's Zond5 mission, which flew some living creatures, including Russian tortoises, in a cislunar loop around the Moon and returned them to Earth on September 21. There was speculation within NASA and the press that they might be preparing to launch cosmonauts on a similar circumlunar mission before the end of 1968. Compounding these concerns, American reconnaissance satellites observed a mockup N1 being rolled to the pad at Baikonur on November 25, 1967.
thumb|right|Erection and mating of spacecraft 103 to Launch Vehicle AS-503 in the VAB for the Apollo8 mission
The Apollo 8 crew, now living in the crew quarters at Kennedy Space Center, received a visit from Charles Lindbergh and his wife, Anne Morrow Lindbergh, the night before the launch. They talked about how, before his 1927 flight, Lindbergh had used a piece of string to measure the distance from New York City to Paris on a globe and from that calculated the fuel needed for the flight. The total he had carried was a tenth of the amount that the Saturn V would burn every second. The next day, the Lindberghs watched the launch of Apollo8 from a nearby dune. |
Apollo 8 | Saturn V redesign | Saturn V redesign
The Saturn V rocket used by Apollo8 was designated AS-503, or the "03rd" model of the SaturnV ("5") rocket to be used in the Apollo-Saturn ("AS") program. When it was erected in the Vehicle Assembly Building on December 20, 1967, it was thought that the rocket would be used for an uncrewed Earth-orbit test flight carrying a boilerplate command and service module. Apollo6 had suffered several major problems during its April 1968 flight, including severe pogo oscillation during its first stage, two second-stage engine failures, and a third stage that failed to reignite in orbit. Without assurances that these problems had been rectified, NASA administrators could not justify risking a crewed mission until additional uncrewed test flights proved the Saturn V was ready.
Teams from the MSFC went to work on the problems. Of primary concern was the pogo oscillation, which would not only hamper engine performance, but could exert significant g-forces on a crew. A task force of contractors, NASA agency representatives, and MSFC researchers concluded that the engines vibrated at a frequency similar to the frequency at which the spacecraft itself vibrated, causing a resonance effect that induced oscillations in the rocket. A system that used helium gas to absorb some of these vibrations was installed.
thumb|upright=1.4|left|Apollo 8 atop SaturnV being rolled out to Pad 39A atop the crawler-transporter
Of equal importance was the failure of three engines during flight. Researchers quickly determined that a leaking hydrogen fuel line ruptured when exposed to vacuum, causing a loss of fuel pressure in engine two. When an automatic shutoff attempted to close the liquid hydrogen valve and shut down engine two, it had accidentally shut down engine three's liquid oxygen due to a miswired connection. As a result, engine three failed within one second of engine two's shutdown. Further investigation revealed the same problem for the third-stage engine—a faulty igniter line. The team modified the igniter lines and fuel conduits, hoping to avoid similar problems on future launches.
The teams tested their solutions in August 1968 at the MSFC. A Saturn stage IC was equipped with shock-absorbing devices to demonstrate the team's solution to the problem of pogo oscillation, while a Saturn Stage II was retrofitted with modified fuel lines to demonstrate their resistance to leaks and ruptures in vacuum conditions. Once NASA administrators were convinced that the problems had been solved, they gave their approval for a crewed mission using AS-503.
The Apollo 8 spacecraft was placed on top of the rocket on September 21, and the rocket made the slow journey to the launch pad atop one of NASA's two massive crawler-transporters on October9. Testing continued all through December until the day before launch, including various levels of readiness testing from December5 through 11. Final testing of modifications to address the problems of pogo oscillation, ruptured fuel lines, and bad igniter lines took place on December 18, three days before the scheduled launch. |
Apollo 8 | Mission | Mission |
Apollo 8 | Parameter summary | Parameter summary
thumb|upright=3.0|Mission profile
As the first crewed spacecraft to orbit more than one celestial body, Apollo8's profile had two different sets of orbital parameters, separated by a translunar injection maneuver. Apollo lunar missions would begin with a nominal circular Earth parking orbit. Apollo8 was launched into an initial orbit with an apogee of and a perigee of , with an inclination of 32.51° to the Equator, and an orbital period of 88.19 minutes. Propellant venting increased the apogee by over the 2hours, 44 minutes, and 30 seconds spent in the parking orbit.
This was followed by a trans-lunar injection (TLI) burn of the S-IVB third stage for 318 seconds, accelerating the command and service module and LM test article from an orbital velocity of to the injection velocity of which set a record for the highest speed, relative to Earth, that humans had ever traveled. This speed was slightly less than the Earth's escape velocity of , but put Apollo8 into an elongated elliptical Earth orbit, close enough to the Moon to be captured by the Moon's gravity.
The standard lunar orbit for Apollo missions was planned as a nominal circular orbit above the Moon's surface. Initial lunar orbit insertion was an ellipse with a perilune of and an apolune of , at an inclination of 12° from the lunar equator. This was then circularized at , with an orbital period of 128.7 minutes. The effect of lunar mass concentrations ("mascons") on the orbit was found to be greater than initially predicted; over the course of the ten lunar orbits lasting twenty hours, the orbital distance was perturbated to .
Apollo 8 achieved a maximum distance from Earth of . |
Apollo 8 | Launch and trans-lunar injection | Launch and trans-lunar injection
thumb|Apollo 8 launch
Apollo 8 was launched at 12:51:00 UTC (07:51:00 Eastern Standard Time) on December 21, 1968, using the Saturn V's three stages to achieve Earth orbit. The S-IC first stage landed in the Atlantic Ocean at , and the S-II second stage landed at . The S-IVB third stage injected the craft into Earth orbit and remained attached to perform the TLI burn that would put the spacecraft on a trajectory to the Moon.
Once the vehicle reached Earth orbit, both the crew and Houston flight controllers spent the next 2hours and 38 minutes checking that the spacecraft was in proper working order and ready for TLI. The proper operation of the S-IVB third stage of the rocket was crucial, and in the last uncrewed test, it had failed to reignite for this burn. Collins was the first CAPCOM on duty, and at 2hours, 27 minutes and 22 seconds after launch he radioed, "Apollo8. You are Go for TLI." This communication meant that Mission Control had given official permission for Apollo8 to go to the Moon. The S-IVB engine ignited on time and performed the TLI burn perfectly. Over the next five minutes, the spacecraft's speed increased from .
After the S-IVB had placed the mission on course for the Moon, the command and service modules (CSM), the remaining Apollo8 spacecraft, separated from it. The crew then rotated the spacecraft to take photographs of the spent stage and then practiced flying in formation with it. As the crew rotated the spacecraft, they had their first views of the Earth as they moved away from it—this marked the first time humans had viewed the whole Earth at once. Borman became worried that the S-IVB was staying too close to the CSM and suggested to Mission Control that the crew perform a separation maneuver. Mission Control first suggested pointing the spacecraft towards Earth and using the small reaction control system (RCS) thrusters on the service module (SM) to add to their velocity away from the Earth, but Borman did not want to lose sight of the S-IVB. After discussion, the crew and Mission Control decided to burn in the Earth direction to increase speed, but at instead. The time needed to prepare and perform the additional burn put the crew an hour behind their onboard tasks.
thumb|Apollo 8 S-IVB rocket stage shortly after separation. The LM test article, a circular boilerplate model of the LM, is visible with four triangular legs connecting it to the stage.
Five hours after launch, Mission Control sent a command to the S-IVB to vent its remaining fuel, changing its trajectory. The S-IVB, with the test article attached, posed no further hazard to Apollo8, passing the orbit of the Moon and going into a solar orbit with an inclination of 23.47° from the Earth's equatorial plane, and an orbital period of 340.80 days. It became a derelict object, and will continue to orbit the Sun for many years, if not retrieved.
The Apollo 8 crew were the first humans to pass through the Van Allen radiation belts, which extend up to from Earth. Scientists predicted that passing through the belts quickly at the spacecraft's high speed would cause a radiation dosage of no more than a chest X-ray, or 1milligray (mGy; during a year, the average human receives a dose of 2to 3mGy from background radiation). To record the actual radiation dosages, each crew member wore a Personal Radiation Dosimeter that transmitted data to Earth, as well as three passive film dosimeters that showed the cumulative radiation experienced by the crew. By the end of the mission, the crew members experienced an average radiation dose of 1.6 mGy. Sec. 2, Ch. 3. |
Apollo 8 | Lunar trajectory | Lunar trajectory
Lovell's main job as Command Module Pilot was as navigator. Although Mission Control normally performed all the navigation calculations, it was necessary to have a crew member adept at navigation so that the crew could return to Earth in case communication with Mission Control was lost. Lovell navigated by star sightings using a sextant built into the spacecraft, measuring the angle between a star and the Earth's (or the Moon's) horizon. This task was made difficult by a large cloud of debris around the spacecraft, which made it hard to distinguish the stars.
By seven hours into the mission, the crew was about 1hour and 40 minutes behind flight plan because of the problems in moving away from the S-IVB and Lovell's obscured star sightings. The crew placed the spacecraft into Passive Thermal Control (PTC), also called "barbecue roll", in which the spacecraft rotated about once per hour around its long axis to ensure even heat distribution across the surface of the spacecraft. In direct sunlight, parts of the spacecraft's outer surface could be heated to over , while the parts in shadow would be . These temperatures could cause the heat shield to crack and propellant lines to burst. Because it was impossible to get a perfect roll, the spacecraft swept out a cone as it rotated. The crew had to make minor adjustments every half hour as the cone pattern got larger and larger.
thumb|left|The first image taken by humans of the whole Earth, probably photographed by William Anders. (time tag: 003:42:55) South America is visible in the lower half.
The first mid-course correction came eleven hours into the flight. The crew had been awake for more than 16 hours. Before launch, NASA had decided at least one crew member should be awake at all times to deal with problems that might arise. Borman started the first sleep shift but found sleeping difficult because of the constant radio chatter and mechanical noises. Testing on the ground had shown that the service propulsion system (SPS) engine had a small chance of exploding when burned for long periods unless its combustion chamber was "coated" first by burning the engine for a short period. This first correction burn was only 2.4 seconds and added about velocity prograde (in the direction of travel). This change was less than the planned , because of a bubble of helium in the oxidizer lines, which caused unexpectedly low propellant pressure. The crew had to use the small RCS thrusters to make up the shortfall. Two later planned mid-course corrections were canceled because the Apollo8 trajectory was found to be perfect.
About an hour after starting his sleep shift, Borman obtained permission from ground control to take a Seconal sleeping pill. The pill had little effect. Borman eventually fell asleep, and then awoke feeling ill. He vomited twice and had a bout of diarrhea; this left the spacecraft full of small globules of vomit and feces, which the crew cleaned up as well as they could. Borman initially did not want everyone to know about his medical problems, but Lovell and Anders wanted to inform Mission Control. The crew decided to use the Data Storage Equipment (DSE), which could tape voice recordings and telemetry and dump them to Mission Control at high speed. After recording a description of Borman's illness they asked Mission Control to check the recording, stating that they "would like an evaluation of the voice comments".
The Apollo 8 crew and Mission Control medical personnel held a conference using an unoccupied second-floor control room (there were two identical control rooms in Houston, on the second and third floors, only one of which was used during a mission). The conference participants concluded that there was little to worry about and that Borman's illness was either a 24-hour flu, as Borman thought, or a reaction to the sleeping pill. Researchers now believe that he was suffering from space adaptation syndrome, which affects about a third of astronauts during their first day in space as their vestibular system adapts to weightlessness. Space adaptation syndrome had not occurred on previous spacecraft (Mercury and Gemini), because those astronauts could not move freely in the small cabins of those spacecraft. The increased cabin space in the Apollo command module afforded astronauts greater freedom of movement, contributing to symptoms of space sickness for Borman and, later, astronaut Rusty Schweickart during Apollo9.
thumb|right|Still from film of the crew taken while they were in orbit around the Moon. Frank Borman is in the center.
The cruise phase was a relatively uneventful part of the flight, except for the crew's checking that the spacecraft was in working order and that they were on course. During this time, NASA scheduled a television broadcast at 31 hours after launch. The Apollo8 crew used a camera that broadcast in black-and-white only, using a Vidicon tube. The camera had two lenses, a very wide-angle (160°) lens, and a telephoto (9°) lens.
During this first broadcast, the crew gave a tour of the spacecraft and attempted to show how the Earth appeared from space. However, difficulties aiming the narrow-angle lens without the aid of a monitor to show what it was looking at made showing the Earth impossible. Additionally, without proper filters, the Earth image became saturated by any bright source. In the end, all the crew could show the people watching back on Earth was a bright blob. After broadcasting for 17 minutes, the rotation of the spacecraft took the high-gain antenna out of view of the receiving stations on Earth and they ended the transmission with Lovell wishing his mother a happy birthday.
By this time, the crew had completely abandoned the planned sleep shifts. Lovell went to sleep hours into the flight – three-and-a-half hours before he had planned to. A short while later, Anders also went to sleep after taking a sleeping pill. The crew was unable to see the Moon for much of the outward cruise. Two factors made the Moon almost impossible to see from inside the spacecraft: three of the five windows fogging up due to out-gassed oils from the silicone sealant, and the attitude required for passive thermal control. It was not until the crew had gone behind the Moon that they would be able to see it for the first time.
Apollo 8 made a second television broadcast at 55 hours into the flight. This time, the crew rigged up filters meant for the still cameras so they could acquire images of the Earth through the telephoto lens. Although difficult to aim, as they had to maneuver the entire spacecraft, the crew was able to broadcast back to Earth the first television pictures of the Earth. The crew spent the transmission describing the Earth, what was visible, and the colors they could see. The transmission lasted 23 minutes. |
Apollo 8 | Lunar sphere of influence | Lunar sphere of influence
thumb|left|This photograph of the Moon was taken from Apollo8 at a point above 70 degrees east longitude.
At about 55 hours and 40 minutes into the flight, and 13 hours before entering lunar orbit, the crew of Apollo8 became the first humans to enter the gravitational sphere of influence of another celestial body. In other words, the effect of the Moon's gravitational force on Apollo8 became stronger than that of the Earth. At the time it happened, Apollo8 was from the Moon and had a speed of relative to the Moon. This historic moment was of little interest to the crew, since they were still calculating their trajectory with respect to the launch pad at Kennedy Space Center. They would continue to do so until they performed their last mid-course correction, switching to a reference frame based on ideal orientation for the second engine burn they would make in lunar orbit.
The last major event before Lunar Orbit Insertion (LOI) was a second mid-course correction. It was in retrograde (against the direction of travel) and slowed the spacecraft down by , effectively reducing the closest distance at which the spacecraft would pass the Moon. At exactly 61 hours after launch, about from the Moon, the crew burned the RCS for 11 seconds. They would now pass from the lunar surface.
At 64 hours into the flight, the crew began to prepare for Lunar Orbit Insertion1 (LOI-1). This maneuver had to be performed perfectly, and due to orbital mechanics had to be on the far side of the Moon, out of contact with the Earth. After Mission Control was polled for a "go/no go" decision, the crew was told at 68 hours that they were Go and "riding the best bird we can find". Lovell replied, "We'll see you on the other side", and for the first time in history, humans travelled behind the Moon and out of radio contact with the Earth. Frances "Poppy" Northcutt, who was the first woman in NASA's mission control and helped calculate the return to Earth trajectory for this mission, recounts what it was like when Apollo 8 went behind the Moon for the first time in an interview: "That was a very nerve-racking period on the team I was on, and I think it was a very nerve-racking period in general because of this thing with losing signal. You've got this big mystery going on there on the backside of the Moon. You do not know what's happening and there's not a darn thing anybody here can do about it until we hear from them."
With ten minutes remaining before LOI-1, the crew began one last check of the spacecraft systems and made sure that every switch was in its correct position. At that time, they finally got their first glimpses of the Moon. They had been flying over the unlit side, and it was Lovell who saw the first shafts of sunlight obliquely illuminating the lunar surface. The LOI burn was only two minutes away, so the crew had little time to appreciate the view. |
Apollo 8 | Lunar orbit | Lunar orbit
The SPS was ignited at 69 hours, 8minutes, and 16 seconds after launch and burned for 4minutes and 7seconds, placing the Apollo8 spacecraft in orbit around the Moon. The crew described the burn as being the longest four minutes of their lives. If the burn had not lasted exactly the correct amount of time, the spacecraft could have ended up in a highly elliptical lunar orbit or even been flung off into space. If it had lasted too long, they could have struck the Moon. After making sure the spacecraft was working, they finally had a chance to look at the Moon, which they would orbit for the next 20 hours.
On Earth, Mission Control continued to wait. If the crew had not burned the engine, or the burn had not lasted the planned length of time, the crew would have appeared early from behind the Moon. Exactly at the calculated moment the signal was received from the spacecraft, indicating it was in a orbit around the Moon.
After reporting on the status of the spacecraft, Lovell gave the first description of what the lunar surface looked like:
thumb|A portion of the lunar far side as seen from Apollo8
Lovell continued to describe the terrain they were passing over. One of the crew's major tasks was reconnaissance of planned future landing sites on the Moon, especially one in Mare Tranquillitatis that was planned as the Apollo11 landing site. The launch time of Apollo8 had been chosen to give the best lighting conditions for examining the site. A film camera had been set up in one of the spacecraft windows to record one frame per second of the Moon below. Bill Anders spent much of the next 20 hours taking as many photographs as possible of targets of interest. By the end of the mission, the crew had taken over eight hundred 70 mm still photographs and of 16 mm movie film.
Throughout the hour that the spacecraft was in contact with Earth, Borman kept asking how the data for the SPS looked. He wanted to make sure that the engine was working and could be used to return early to the Earth if necessary. He also asked that they receive a "go/no go" decision before they passed behind the Moon on each orbit.
As they reappeared for their second pass in front of the Moon, the crew set up equipment to broadcast a view of the lunar surface. Anders described the craters that they were passing over. At the end of this second orbit, they performed an 11-second LOI-2 burn of the SPS to circularize the orbit to .
Throughout the next two orbits, the crew continued to check the spacecraft and to observe and photograph the Moon. During the third pass, Borman read a small prayer for his church. He had been scheduled to participate in a service at St. Christopher's Episcopal Church near Seabrook, Texas, but due to the Apollo8 flight, he was unable to attend. A fellow parishioner and engineer at Mission Control, Rod Rose, suggested that Borman read the prayer, which could be recorded and then replayed during the service. |
Apollo 8 | ''Earthrise'' and Genesis broadcast | Earthrise and Genesis broadcast
thumb|The Earthrise image
thumb|Apollo 8's 1968 Christmas Eve broadcast and reading from the Book of Genesis
When the spacecraft came out from behind the Moon for its fourth pass across the front, the crew witnessed an "Earthrise" in person for the first time in human history. NASA's Lunar Orbiter 1 had taken the first picture of an Earthrise from the vicinity of the Moon, on August 23, 1966. Anders saw the Earth emerging from behind the lunar horizon and called in excitement to the others, taking a black-and-white photograph as he did so. Anders asked Lovell for color film and then took Earthrise, a now famous color photo, later picked by Life magazine as one of its hundred photos of the century.
Due to the synchronous rotation of the Moon about the Earth, Earthrise is not generally visible from the lunar surface. This is because, as seen from any one place on the Moon's surface, Earth remains in approximately the same position in the lunar sky, either above or below the horizon. Earthrise is generally visible only while orbiting the Moon, and at selected surface locations near the Moon's limb, where libration carries the Earth slightly above and below the lunar horizon.
Anders continued to take photographs while Lovell assumed control of the spacecraft so that Borman could rest. Despite the difficulty resting in the cramped and noisy spacecraft, Borman was able to sleep for two orbits, awakening periodically to ask questions about their status. Borman awoke fully when he started to hear his fellow crew members make mistakes. They were beginning to not understand questions and had to ask for the answers to be repeated. Borman realized that everyone was extremely tired from not having a good night's sleep in over three days. He ordered Anders and Lovell to get some sleep and that the rest of the flight plan regarding observing the Moon be scrubbed. Anders initially protested, saying that he was fine, but Borman would not be swayed. Anders finally agreed under the condition that Borman would set up the camera to continue to take automatic pictures of the Moon. Borman also remembered that there was a second television broadcast planned, and with so many people expected to be watching, he wanted the crew to be alert. For the next two orbits, Anders and Lovell slept while Borman sat at the helm.
left|thumb|Apollo 8 Genesis reading
As they rounded the Moon for the ninth time, the astronauts began the second television transmission. Borman introduced the crew, followed by each man giving his impression of the lunar surface and what it was like to be orbiting the Moon. Borman described it as being "a vast, lonely, forbidding expanse of nothing". Then, after talking about what they were flying over, Anders said that the crew had a message for all those on Earth. Each man on board read a section from the Biblical creation story from the Book of Genesis. Borman finished the broadcast by wishing a Merry Christmas to everyone on Earth. His message appeared to sum up the feelings that all three crewmen had from their vantage point in lunar orbit. Borman said, "And from the crew of Apollo8, we close with good night, good luck, a Merry Christmas and God bless all of you—all of you on the good Earth." Ch.20-9.
The only task left for the crew at this point was to perform the trans-Earth injection (TEI), which was scheduled for hours after the end of the television transmission. The TEI was the most critical burn of the flight, as any failure of the SPS to ignite would strand the crew in lunar orbit, with little hope of escape. As with the previous burn, the crew had to perform the maneuver above the far side of the Moon, out of contact with Earth. The burn occurred exactly on time. The spacecraft telemetry was reacquired as it re-emerged from behind the Moon at 89 hours, 28 minutes, and 39 seconds, the exact time calculated. When voice contact was regained, Lovell announced, "Please be informed, there is a Santa Claus", to which Ken Mattingly, the current CAPCOM, replied, "That's affirmative, you are the best ones to know." The spacecraft began its journey back to Earth on December 25, Christmas Day. |
Apollo 8 | Unplanned manual realignment | Unplanned manual realignment
Later, Lovell used some otherwise idle time to do some navigational sightings, maneuvering the module to view various stars by using the computer keyboard. He accidentally erased some of the computer's memory, which caused the inertial measurement unit (IMU) to contain data indicating that the module was in the same relative orientation it had been in before lift-off; the IMU then fired the thrusters to "correct" the module's attitude.
Once the crew realized why the computer had changed the module's attitude, they realized that they would have to reenter data to tell the computer the module's actual orientation. It took Lovell ten minutes to figure out the right numbers, using the thrusters to get the stars Rigel and Sirius aligned, and another 15 minutes to enter the corrected data into the computer. Sixteen months later, during the Apollo13 mission, Lovell would have to perform a similar manual realignment under more critical conditions after the module's IMU had to be turned off to conserve energy. |
Apollo 8 | Cruise back to Earth and reentry | Cruise back to Earth and reentry
thumb|left|Reentry, December 27, 1968, photographed from a KC-135 Stratotanker at 40,000 feet|alt=White streaks of light, with bright spots on the right side of them, fill the bottom of the frame. A larger yellow-tinted sphere with a streak is in the center of the frame. The background is black space.
The cruise back to Earth was mostly a time for the crew to relax and monitor the spacecraft. As long as the trajectory specialists had calculated everything correctly, the spacecraft would reenter Earth's atmosphere two-and-a-half days after TEI and splash down in the Pacific.
On Christmas afternoon, the crew made their fifth television broadcast. This time, they gave a tour of the spacecraft, showing how an astronaut lived in space. When they finished broadcasting, they found a small present from Slayton in the food locker: a real turkey dinner with stuffing, in the same kind of pack given to the troops in Vietnam.
Another Slayton surprise was a gift of three miniature bottles of brandy, which Borman ordered the crew to leave alone until after they landed. They remained unopened, even years after the flight. There were also small presents to the crew from their wives. The next day, at about 124 hours into the mission, the sixth and final TV transmission showed the mission's best video images of the Earth, during a four-minute broadcast. After two uneventful days, the crew prepared for reentry. The computer would control the reentry, and all the crew had to do was put the spacecraft in the correct attitude, with the blunt end forward. In the event of computer failure, Borman was ready to take over.
thumb|Crew of Apollo 8 addressing the crew of USS Yorktown after successful splashdown and recovery
Separation from the service module prepared the command module for reentry by exposing the heat shield and shedding unneeded mass. The service module would burn up in the atmosphere as planned. Six minutes before they hit the top of the atmosphere, the crew saw the Moon rising above the Earth's horizon, just as had been calculated by the trajectory specialists. As the module hit the thin outer atmosphere, the crew noticed that it was becoming hazy outside as glowing plasma formed around the spacecraft. The spacecraft started slowing down, and the deceleration peaked at . With the computer controlling the descent by changing the attitude of the spacecraft, Apollo8 rose briefly like a skipping stone before descending to the ocean. At , the drogue parachute deployed, stabilizing the spacecraft, followed at by the three main parachutes. The spacecraft splashdown position was officially reported as in the North Pacific Ocean, southwest of Hawaii at 15:51:42 UTC on December 27, 1968.
thumb|Command module on the deck of
When the spacecraft hit the water, the parachutes dragged it over and left it upside down, in what was termed Stable2 position. As they were buffeted by a swell, Borman vomited, waiting for the three flotation balloons to right the spacecraft. About six minutes after splashdown, the command module was righted into a normal apex-up (Stable 1) orientation by its inflatable bag uprighting system. The first frogman from aircraft carrier arrived 43 minutes after splashdown. Forty-five minutes later, the crew was safe on the flight deck of the Yorktown. |
Apollo 8 | Legacy | Legacy |
Apollo 8 | Historical importance | Historical importance
Apollo 8 came at the end of 1968, a year that had seen much upheaval in the United States and most of the world. Even though the year saw political assassinations, political unrest in the streets of Europe and America, and the Prague Spring, Time magazine chose the crew of Apollo8 as its Men of the Year for 1968, recognizing them as the people who most influenced events in the preceding year. They had been the first people ever to leave the gravitational influence of the Earth and orbit another celestial body. They had survived a mission that even the crew themselves had rated as having only a fifty-fifty chance of fully succeeding. The effect of Apollo8 was summed up in a telegram from a stranger, received by Borman after the mission, that stated simply, "Thank you Apollo8. You saved 1968."
One of the most famous aspects of the flight was the Earthrise picture that the crew took as they came around for their fourth orbit of the Moon. This was the first time that humans had taken such a picture while actually behind the camera, and it has been credited as one of the inspirations of the first Earth Day in 1970. It was selected as the first of Life magazine's 100 Photographs That Changed the World.
thumb|left|Apollo 8 astronauts return to Houston after their mission
Apollo 11 astronaut Michael Collins said, "Eight's momentous historic significance was foremost"; while space historian Robert K. Poole saw Apollo8 as the most historically significant of all the Apollo missions. The mission was the most widely covered by the media since the first American orbital flight, Mercury-Atlas 6 by John Glenn, in 1962. There were 1,200 journalists covering the mission, with the BBC's coverage broadcast in 54 countries in 15 different languages. The Soviet newspaper Pravda featured a quote from , Chairman of the Soviet Interkosmos program, who described the flight as an "outstanding achievement of American space sciences and technology". It is estimated that a quarter of the people alive at the time saw—either live or delayed—the Christmas Eve transmission during the ninth orbit of the Moon. The Apollo8 broadcasts won an Emmy Award, the highest honor given by the Academy of Television Arts & Sciences.
Madalyn Murray O'Hair, an atheist, later caused controversy by bringing a lawsuit against NASA over the reading from Genesis. O'Hair wanted the courts to ban American astronauts—who were all government employees—from public prayer in space. Though the case was rejected by the Supreme Court of the United States, apparently for lack of jurisdiction in outer space, it caused NASA to be skittish about the issue of religion throughout the rest of the Apollo program. Buzz Aldrin, on Apollo11, self-communicated Presbyterian Communion on the surface of the Moon after landing; he refrained from mentioning this publicly for several years and referred to it only obliquely at the time.
thumb|right|upright|Apollo 8 commemorative stamp
In 1969, the United States Post Office Department issued a postage stamp (Scott catalogue #1371) commemorating the Apollo8 flight around the Moon. The stamp featured a detail of the famous photograph of the Earthrise over the Moon taken by Anders on Christmas Eve, and the words, "In the beginning God...", the first words of the book of Genesis. In January 1969, just 18 days after the crew's return to Earth, they appeared in the Super Bowl III pre-game show, reciting the Pledge of Allegiance, before the national anthem was performed by trumpeter Lloyd Geisler of the Washington National Symphony Orchestra.NFL's website erroneously states that Anita Bryant performed the anthem, but NBC's broadcast of game, available from the Paley Center for Media's collection, shows that Geisler performed it. |
Apollo 8 | Spacecraft location | Spacecraft location
In January 1970, the spacecraft was delivered to Osaka, Japan, for display in the U.S. pavilion at Expo '70. It is now displayed at the Chicago Museum of Science and Industry, along with a collection of personal items from the flight donated by Lovell and the space suit worn by Frank Borman. Jim Lovell's Apollo8 space suit is on public display in the Visitor Center at NASA's Glenn Research Center. Bill Anders's space suit is on display at the Science Museum in London, United Kingdom. |
Apollo 8 | In popular culture | In popular culture
Apollo 8's historic mission has been depicted and referred to in several forms, both documentary and fiction. The various television transmissions and 16 mm footage shot by the crew of Apollo8 were compiled and released by NASA in the 1969 documentary Debrief: Apollo8, hosted by Burgess Meredith. Debrief: Apollo 8 was released as a bonus feature for the Discovery Channel's miniseries DVD release. In addition, Spacecraft Films released, in 2003, a three-disc DVD set containing all of NASA's TV and 16 mm film footage related to the mission, including all TV transmissions from space, training and launch footage, and motion pictures taken in flight. Other documentaries include "Race to the Moon" (2005) as part of season 18 of American Experience and In the Shadow of the Moon (2007). Apollo's Daring Mission aired on PBS' Nova in December 2018, marking the flight's 50th anniversary.
The 1994 album The Songs of Distant Earth by Mike Oldfield uses the Anders' reading for the cut "In The Beginning".
Parts of the mission are dramatized in the 1998 miniseries From the Earth to the Moon episode "1968". The S-IVB stage of Apollo8 was also portrayed as the location of an alien device in the 1970 UFO episode "Conflict". Apollo8's lunar orbit insertion was chronicled with actual recordings in the song "The Other Side", on the 2015 album The Race for Space, by the band Public Service Broadcasting.
A documentary film, First to the Moon: The Journey of Apollo 8 was released in 2018. |
Apollo 8 | See also | See also
Apollo 8 (book)
List of missions to the Moon |
Apollo 8 | Notes | Notes |
Apollo 8 | References | References |
Apollo 8 | Bibliography | Bibliography
|
Apollo 8 | External links | External links
"Apollo 8" at Encyclopedia Astronautica
Article about the 40th anniversary of Apollo8
Multimedia
Apollo 8: Go for TLI 1969 NASA film at the Internet Archive
Debrief: Apollo 8 1969 NASA film at the Internet Archive
"Apollo 07 and 08 16mm Onboard Film (1968)" raw footage taken from Apollos 7and8 at the Internet Archive
Apollo 8 Around the Moon and Back 2018 YouTube video
Apollo 08
Category:Crewed missions to the Moon
Category:Spacecraft launched in 1968
Category:1968 in the United States
Category:Spacecraft which reentered in 1968
Category:December 1968
Category:Spacecraft launched by Saturn rockets
Category:Jim Lovell
Category:William Anders
Category:Frank Borman
Category:Successful space missions |
Apollo 8 | Table of Content | Short description, Background, Framework, Prime crew, Backup crew, Support personnel, Mission insignia and callsign, Preparations, Mission schedule, Saturn V redesign, Mission, Parameter summary, Launch and trans-lunar injection, Lunar trajectory, Lunar sphere of influence, Lunar orbit, ''Earthrise'' and Genesis broadcast, Unplanned manual realignment, Cruise back to Earth and reentry, Legacy, Historical importance, Spacecraft location, In popular culture, See also, Notes, References, Bibliography, External links |
A Modest Proposal | Short description | A Modest Proposal for Preventing the Children of Poor People from Being a Burthen to Their Parents or Country, and for Making Them Beneficial to the Publick, commonly referred to as A Modest Proposal, is a Juvenalian satirical essay written and published by Anglo-Irish writer and clergyman Jonathan Swift in 1729. The essay suggests that poor people in Ireland could ease their economic troubles by selling their children as food to the elite. Swift's use of satirical hyperbole was intended to mock hostile attitudes towards the poor and anti-Catholicism among the Protestant Ascendancy as well as the Dublin Castle administration's policies in general.Swift notes that "the number of Popish infants, is at least three to one in this kingdom, and therefore it will have one another collateral advantage, by lessening the number of Papists among us." In English writing, the phrase "a modest proposal" is now conventionally an allusion to this style of straight-faced satire. |
A Modest Proposal | Synopsis | Synopsis
thumb|A painting of Jonathan Swift
Swift's essay is widely held to be one of the greatest examples of sustained irony in the history of English literature. Much of its shock value derives from the fact that the first portion of the essay describes the plight of starving beggars in Ireland, so that the reader is unprepared for the surprise of Swift's solution when he states: "A young healthy child well nursed, is, at a year old, a most delicious nourishing and wholesome food, whether stewed, roasted, baked, or boiled; and I make no doubt that it will equally serve in a fricassee, or a ragout."
Swift goes to great lengths to support his argument, including a list of possible preparation styles for the children, and calculations showing the financial benefits of his suggestion. He uses methods of argument throughout his essay which lampoon the then-influential William Petty and the social engineering popular among followers of Francis Bacon. These lampoons include appealing to the authority of "a very knowing American of my acquaintance in London" and "the famous Psalmanazar, a native of the island Formosa" (who had already confessed to not being from Formosa in 1706).
In the tradition of Roman satire, Swift introduces the reforms he is actually suggesting by paralipsis: |
A Modest Proposal | Population solutions | Population solutions
George Wittkowsky argued that Swift's main target in A Modest Proposal was not the conditions in Ireland, but rather the can-do spirit of the times that led people to devise a number of illogical schemes that would purportedly solve social and economic ills.Wittkowsky, Swift’s Modest Proposal, p. 76 Swift was especially attacking projects that tried to fix population and labour issues with a simple cure-all solution.Wittkowsky, Swift’s Modest Proposal, p. 85 A memorable example of these sorts of schemes "involved the idea of running the poor through a joint-stock company". In response, Swift's Modest Proposal was "a burlesque of projects concerning the poor"Wittkowsky, Swift's Modest Proposal, p. 88 that were in vogue during the early 18th century.
Ian McBride argues that the point of A Modest Proposal was to "find a suitably decisive means of dehumanizing the settlers who had failed so comprehensively to meet their social responsibilities."McBride, Ian (2019). "The Politics of A Modest Proposal: Swift and the Irish Crisis of the Late 1720s." Past & Present. 244 (1): 89–122.
A Modest Proposal also targets the calculating way people perceived the poor in designing their projects. The pamphlet targets reformers who "regard people as commodities".Wittkowsky, Swift's Modest Proposal, p. 101 In the piece, Swift adopts the "technique of a political arithmetician"Wittkowsky, Swift's Modest Proposal, p. 95 to show the utter ridiculousness of trying to prove any proposal with dispassionate statistics.
Critics differ about Swift's intentions in using this faux-mathematical philosophy. Edmund Wilson argues that statistically "the logic of the 'Modest proposal' can be compared with defence of crime (arrogated to Marx) in which he argues that crime takes care of the superfluous population". Wittkowsky counters that Swift's satiric use of statistical analysis is an effort to enhance his satire that "springs from a spirit of bitter mockery, not from the delight in calculations for their own sake".Wittkowsky, Swift's Modest Proposal, p. 98 |
A Modest Proposal | Rhetoric | Rhetoric
Author Charles K. Smith argues that Swift's rhetorical style persuades the reader to detest the speaker and pity the Irish. Swift's specific strategy is twofold, using a "trap"Smith, Toward a Participatory Rhetoric, p. 135 to create sympathy for the Irish and a dislike of the narrator who, in the span of one sentence, "details vividly and with rhetorical emphasis the grinding poverty" but feels emotion solely for members of his own class.Smith, Toward a Participatory Rhetoric, p. 136 Swift's use of gripping details of poverty and his narrator's cool approach towards them create "two opposing points of view" that "alienate the reader, perhaps unconsciously, from a narrator who can view with 'melancholy' detachment a subject that Swift has directed us, rhetorically, to see in a much less detached way."
Swift has his proposer further degrade the Irish by using language ordinarily reserved for animals. Lewis argues that the speaker uses "the vocabulary of animal husbandry"Smith, Toward a Participatory Rhetoric, p. 138 to describe the Irish. Once the children have been commodified, Swift's rhetoric can easily turn "people into animals, then meat, and from meat, logically, into tonnage worth a price per pound".
Swift uses the proposer's serious tone to highlight the absurdity of his proposal. In making his argument, the speaker uses the conventional, textbook-approved order of argument from Swift's time (which was derived from the Latin rhetorician Quintilian).Smith, Toward a Participatory Rhetoric, p. 139 The contrast between the "careful control against the almost inconceivable perversion of his scheme" and "the ridiculousness of the proposal" create a situation in which the reader has "to consider just what perverted values and assumptions would allow such a diligent, thoughtful, and conventional man to propose so perverse a plan". |
A Modest Proposal | Influences | Influences
Scholars have speculated about which earlier works Swift may have had in mind when he wrote A Modest Proposal. |
A Modest Proposal | Tertullian's ''Apology'' | Tertullian's Apology
James William Johnson argues that A Modest Proposal was largely influenced and inspired by Tertullian's Apology: a satirical attack against early Roman persecution of Christianity. Johnson believes that Swift saw major similarities between the two situations.Johnson, Tertullian and A Modest Proposal, p. 563 Johnson notes Swift's obvious affinity for Tertullian and the bold stylistic and structural similarities between the works A Modest Proposal and Apology.Johnson, Tertullian and A Modest Proposal, p. 562 In structure, Johnson points out the same central theme, that of cannibalism and the eating of babies as well as the same final argument, that "human depravity is such that men will attempt to justify their own cruelty by accusing their victims of being lower than human". Stylistically, Swift and Tertullian share the same command of sarcasm and language. In agreement with Johnson, Donald C. Baker points out the similarity between both authors' tones and use of irony. Baker notes the uncanny way that both authors imply an ironic "justification by ownership" over the subject of sacrificing children—Tertullian while attacking pagan parents, and Swift while attacking the mistreatment of the poor in Ireland.Baker, Tertullian and Swift's A Modest Proposal, p. 219 |
A Modest Proposal | Defoe's ''The Generous Projector'' | Defoe's The Generous Projector
It has also been argued that A Modest Proposal was, at least in part, a response to the 1728 essay The Generous Projector or, A Friendly Proposal to Prevent Murder and Other Enormous Abuses, By Erecting an Hospital for Foundlings and Bastard Children by Swift's rival Daniel Defoe. |
A Modest Proposal | Mandeville's ''Modest Defence of Publick Stews'' | Mandeville's Modest Defence of Publick Stews
Bernard Mandeville's Modest Defence of Publick Stews asked to introduce public and state-controlled bordellos. The 1726 paper acknowledges women's interests and—while not being a completely satirical text—has also been discussed as an inspiration for Jonathan Swift's title.Eine Streitschrift…, Essay von Ursula Pia Jauch. Carl Hanser Verlag, München 2001. Mandeville had by 1705 already become famous for The Fable of the Bees and deliberations on private vices and public benefits. |
A Modest Proposal | John Locke's ''First Treatise of Government'' | John Locke's First Treatise of Government
John Locke commented: "Be it then as Sir Robert says, that Anciently, it was usual for Men to sell and Castrate their Children. Let it be, that they exposed them; Add to it, if you please, for this is still greater Power, that they begat them for their Tables to fat and eat them: If this proves a right to do so, we may, by the same Argument, justifie Adultery, Incest and Sodomy, for there are examples of these too, both Ancient and Modern; Sins, which I suppose, have the Principle Aggravation from this, that they cross the main intention of Nature, which willeth the increase of Mankind, and the continuation of the Species in the highest perfection, and the distinction of Families, with the Security of the Marriage Bed, as necessary thereunto". (First Treatise, sec. 59). |
A Modest Proposal | Economic themes | Economic themes
Robert Phiddian's article "Have you eaten yet? The Reader in A Modest Proposal" focuses on two aspects of A Modest Proposal: the voice of Swift and the voice of the Proposer. Phiddian stresses that a reader of the pamphlet must learn to distinguish between the satirical voice of Jonathan Swift and the apparent economic projections of the Proposer. He reminds readers that "there is a gap between the narrator's meaning and the text's, and that a moral-political argument is being carried out by means of parody".Phiddian, Have You Eaten Yet?, p. 6
While Swift's proposal is obviously not a serious economic proposal, George Wittkowsky, author of "Swift's Modest Proposal: The Biography of an Early Georgian Pamphlet", argues that to understand the piece fully it is important to understand the economics of Swift's time. Wittowsky argued that an insufficient number of critics have taken the time to focus directly on mercantilism and theories of labour in Georgian era Britain. "If one regards the Modest Proposal simply as a criticism of condition, about all one can say is that conditions were bad and that Swift's irony brilliantly underscored this fact".Phiddian, Have You Eaten Yet?, p. 3 |
A Modest Proposal | "People are the riches of a nation" | "People are the riches of a nation"
At the start of a new industrial age in the 18th century, it was believed that "people are the riches of the nation", and there was a general faith in an economy that paid its workers low wages because high wages meant workers would work less.Phiddian, Have You Eaten Yet?, p. 4 Furthermore, "in the mercantilist view no child was too young to go into industry". In those times, the "somewhat more humane attitudes of an earlier day had all but disappeared and the laborer had come to be regarded as a commodity".
Louis A. Landa composed a conducive analysis when he noted that it would have been healthier for the Irish economy to more appropriately utilize their human assets by giving the people an opportunity to "become a source of wealth to the nation" or else they "must turn to begging and thievery". This opportunity may have included giving the farmers more coin to work for, diversifying their professions, or even consider enslaving their people to lower coin usage and build up financial stock in Ireland. Landa wrote that, "Swift is maintaining that the maxim—people are the riches of a nation—applies to Ireland only if Ireland is permitted slavery or cannibalism."
Landa presents Swift's A Modest Proposal as a critique of the popular and unjustified maxim of mercantilism in the 18th century that "people are the riches of a nation".Landa, A Modest Proposal and Populousness, p. 161 Swift presents the dire state of Ireland and shows that mere population itself, in Ireland's case, did not always mean greater wealth and economy.Landa, A Modest Proposal and Populousness, p. 165 The uncontrolled maxim fails to take into account that a person who does not produce in an economic or political way makes a country poorer, not richer. Swift also recognises the implications of this fact in making mercantilist philosophy a paradox: the wealth of a country is based on the poverty of the majority of its citizens. Landa argued that Swift was putting the onus "on England of vitiating the working of natural economic law in Ireland" by denying Irishmen "the same natural rights common to the rest of mankind." |
A Modest Proposal | Public reaction | Public reaction
thumb|upright|Allen Bathurst, 1st Earl Bathurst
Swift's essay created a backlash within Georgian society after its publication. The work was aimed at the elite, and they responded in turn. Several prominent members of society wrote to Swift regarding the work. Lord Bathurst's letter (12 February 1729–30) intimated that he certainly understood the message, and interpreted it as a work of comedy: |
A Modest Proposal | Modern usage | Modern usage
A Modest Video Game Proposal is the title of an open letter sent by activist/former attorney Jack Thompson on 10 October 2005.
The 2012 horror film Butcher Boys, written by the original The Texas Chain Saw Massacre scribe Kim Henkel, is said to be an updating of Jonathan Swift's A Modest Proposal. Henkel imagined the descendants of folks who actually took Swift up on his proposal. The film opens with a quote from J. Swift.
The 2023 song "Eat Your Young" written by Irish musician Hozier might be a reference to "A Modest Proposal". It combines themes regarding the anti-war and anti-income-inequality movement, and uses Swift's essay as a framework to compare those modern problems to those same problems during Swift's time.
The July 2023 Channel 4 mockumentary Gregg Wallace: The British Miracle Meat, written by British comedy writer Matt Edmonds, updates A Modest Proposal and presents it in a similar format to Wallace's Inside the Factory, with human meat given as a potential solution to the UK's cost of living crisis. The words "a modest proposal" are used in Wallace's summing up at the end of the programme, and Swift is credited. |
A Modest Proposal | See also | See also
Cannibalism in literature
Child cannibalism |
A Modest Proposal | Notes | Notes |
A Modest Proposal | References | References
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A Modest Proposal | External links | External links
A Modest Proposal (CELT)
A Modest Proposal (Gutenberg)
A Modest Proposal – Annotated text aligned to Common Core Standards
A Modest Proposal BBC Radio 4 In Our Time with Melvyn Bragg
'A modest proposal For preventing the children of poor people From being a Burthen to their Parents or the Country, And for making them Beneficial to the publick. The Third Edition, Dublin, Printed: And Reprinted at London, for Weaver Bickerton, in Devereux-Court near the Middle-Temple, 1730.
Category:1729 essays
Category:1729 in Great Britain
Category:1729 books
Category:Essays by Jonathan Swift
Category:Satirical essays
Category:Pamphlets
Category:British satire
Category:Fiction about cannibalism
Category:Works about Ireland |
A Modest Proposal | Table of Content | Short description, Synopsis, Population solutions, Rhetoric, Influences, Tertullian's ''Apology'', Defoe's ''The Generous Projector'', Mandeville's ''Modest Defence of Publick Stews'', John Locke's ''First Treatise of Government'', Economic themes, "People are the riches of a nation", Public reaction, Modern usage, See also, Notes, References, External links |
Alkali metal | short description | ↓ Period 2 3 4 5 6 7 Legend
primordial element by radioactive decay
The alkali metals consist of the chemical elements lithium (Li), sodium (Na), potassium (K),The symbols Na and K for sodium and potassium are derived from their Latin names, natrium and kalium; these are still the origins of the names for the elements in some languages, such as German and Russian. rubidium (Rb), caesium (Cs), and francium (Fr). Together with hydrogen they constitute group 1, which lies in the s-block of the periodic table. All alkali metals have their outermost electron in an s-orbital: this shared electron configuration results in their having very similar characteristic properties. Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour. This family of elements is also known as the lithium family after its leading element.
The alkali metals are all shiny, soft, highly reactive metals at standard temperature and pressure and readily lose their outermost electron to form cations with charge +1. They can all be cut easily with a knife due to their softness, exposing a shiny surface that tarnishes rapidly in air due to oxidation by atmospheric moisture and oxygen (and in the case of lithium, nitrogen). Because of their high reactivity, they must be stored under oil to prevent reaction with air, and are found naturally only in salts and never as the free elements. Caesium, the fifth alkali metal, is the most reactive of all the metals. All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.
All of the discovered alkali metals occur in nature as their compounds: in order of abundance, sodium is the most abundant, followed by potassium, lithium, rubidium, caesium, and finally francium, which is very rare due to its extremely high radioactivity; francium occurs only in minute traces in nature as an intermediate step in some obscure side branches of the natural decay chains. Experiments have been conducted to attempt the synthesis of element 119, which is likely to be the next member of the group; none were successful. However, ununennium may not be an alkali metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements; even if it does turn out to be an alkali metal, it is predicted to have some differences in physical and chemical properties from its lighter homologues.
Most alkali metals have many different applications. One of the best-known applications of the pure elements is the use of rubidium and caesium in atomic clocks, of which caesium atomic clocks form the basis of the second. A common application of the compounds of sodium is the sodium-vapour lamp, which emits light very efficiently. Table salt, or sodium chloride, has been used since antiquity. Lithium finds use as a psychiatric medication and as an anode in lithium batteries. Sodium, potassium and possibly lithium are essential elements, having major biological roles as electrolytes, and although the other alkali metals are not essential, they also have various effects on the body, both beneficial and harmful.
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Alkali metal | History | History
thumb|alt=A sample of petalite|Petalite, the lithium mineral from which lithium was first isolated
Sodium compounds have been known since ancient times; salt (sodium chloride) has been an important commodity in human activities. While potash has been used since ancient times, it was not understood for most of its history to be a fundamentally different substance from sodium mineral salts. Georg Ernst Stahl obtained experimental evidence which led him to suggest the fundamental difference of sodium and potassium salts in 1702, and Henri-Louis Duhamel du Monceau was able to prove this difference in 1736. The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus Antoine Lavoisier did not include either alkali in his list of chemical elements in 1789.
Pure potassium was first isolated in 1807 in England by Humphry Davy, who derived it from caustic potash (KOH, potassium hydroxide) by the use of electrolysis of the molten salt with the newly invented voltaic pile. Previous attempts at electrolysis of the aqueous salt were unsuccessful due to potassium's extreme reactivity. Potassium was the first metal that was isolated by electrolysis. Later that same year, Davy reported extraction of sodium from the similar substance caustic soda (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different.
thumb|upright|Johann Wolfgang Döbereiner was among the first to notice similarities between what are now known as the alkali metals.
Petalite () was discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada in a mine on the island of Utö, Sweden. However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jöns Jacob Berzelius, detected the presence of a new element while analysing petalite ore. This new element was noted by him to form compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline than the other alkali metals. Berzelius gave the unknown material the name lithion/lithina, from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium, which was known partly for its high abundance in animal blood. He named the metal inside the material lithium. Lithium, sodium, and potassium were part of the discovery of periodicity, as they are among a series of triads of elements in the same group that were noted by Johann Wolfgang Döbereiner in 1850 as having similar properties.
thumb|upright|alt=A sample of lepidolite|Lepidolite, the rubidium mineral from which rubidium was first isolated
Rubidium and caesium were the first elements to be discovered using the spectroscope, invented in 1859 by Robert Bunsen and Gustav Kirchhoff. The next year, they discovered caesium in the mineral water from Bad Dürkheim, Germany. Their discovery of rubidium came the following year in Heidelberg, Germany, finding it in the mineral lepidolite. The names of rubidium and caesium come from the most prominent lines in their emission spectra: a bright red line for rubidium (from the Latin word rubidus, meaning dark red or bright red), and a sky-blue line for caesium (derived from the Latin word caesius, meaning sky-blue).
Around 1865 John Newlands produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the octaves of music, where notes an octave apart have similar musical functions. His version put all the alkali metals then known (lithium to caesium), as well as copper, silver, and thallium (which show the +1 oxidation state characteristic of the alkali metals), together into a group. His table placed hydrogen with the halogens.
thumb|upright=1.75|Dmitri Mendeleev's periodic system proposed in 1871 showing hydrogen and the alkali metals as part of his group I, along with copper, silver, and gold
After 1869, Dmitri Mendeleev proposed his periodic table placing lithium at the top of a group with sodium, potassium, rubidium, caesium, and thallium. Two years later, Mendeleev revised his table, placing hydrogen in group 1 above lithium, and also moving thallium to the boron group. In this 1871 version, copper, silver, and gold were placed twice, once as part of group IB, and once as part of a "group VIII" encompassing today's groups 8 to 11.In the 1869 version of Mendeleev's periodic table, copper and silver were placed in their own group, aligned with hydrogen and mercury, while gold was tentatively placed under uranium and the undiscovered eka-aluminium in the boron group. After the introduction of the 18-column table, the group IB elements were moved to their current position in the d-block, while alkali metals were left in group IA. Later the group's name was changed to group 1 in 1988. The trivial name "alkali metals" comes from the fact that the hydroxides of the group 1 elements are all strong alkalis when dissolved in water.
There were at least four erroneous and incomplete discoveries before Marguerite Perey of the Curie Institute in Paris, France discovered francium in 1939 by purifying a sample of actinium-227, which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one that was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227.Adloff, Jean-Pierre; Kaufman, George B. (25 September 2005). Francium (Atomic Number 87), the Last Discovered Natural Element . The Chemical Educator 10 (5). Retrieved 26 March 2007. Perey then attempted to determine the proportion of beta decay to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure that she later revised to 1%.
The next element below francium (eka-francium) in the periodic table would be ununennium (Uue), element 119. The synthesis of ununennium was first attempted in 1985 by bombarding a target of einsteinium-254 with calcium-48 ions at the superHILAC accelerator at the Lawrence Berkeley National Laboratory in Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.
+ → * → no atomsThe asterisk denotes an excited state.
It is highly unlikely that this reaction will be able to create any atoms of ununennium in the near future, given the extremely difficult task of making sufficient amounts of einsteinium-254, which is favoured for production of ultraheavy elements because of its large mass, relatively long half-life of 270 days, and availability in significant amounts of several micrograms, to make a large enough target to increase the sensitivity of the experiment to the required level; einsteinium has not been found in nature and has only been produced in laboratories, and in quantities smaller than those needed for effective synthesis of superheavy elements. However, given that ununennium is only the first period 8 element on the extended periodic table, it may well be discovered in the near future through other reactions, and indeed an attempt to synthesise it is currently ongoing in Japan. Currently, none of the period 8 elements has been discovered yet, and it is also possible, due to drip instabilities, that only the lower period 8 elements, up to around element 128, are physically possible. No attempts at synthesis have been made for any heavier alkali metals: due to their extremely high atomic number, they would require new, more powerful methods and technology to make. |
Alkali metal | Occurrence | Occurrence |
Alkali metal | In the Solar System | In the Solar System
thumb|upright=2.5|Estimated abundances of the chemical elements in the Solar System. Hydrogen and helium are most common, from the Big Bang. The next three elements (lithium, beryllium, and boron) are rare because they are poorly synthesised in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers, and (2) a general decrease in abundance, as elements become heavier. Iron is especially common because it represents the minimum-energy nuclide that can be made by fusion of helium in supernovae.
The Oddo–Harkins rule holds that elements with even atomic numbers are more common that those with odd atomic numbers, with the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability. All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the noble gases and the alkaline earth metals) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones as the alkali metals from rubidium onward can only be synthesised in supernovae and not in stellar nucleosynthesis. Lithium is also much less abundant than sodium and potassium as it is poorly synthesised in both Big Bang nucleosynthesis and in stars: the Big Bang could only produce trace quantities of lithium, beryllium and boron due to the absence of a stable nucleus with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this bottleneck by the triple-alpha process, fusing three helium nuclei to form carbon, and skipping over those three elements. |
Alkali metal | On Earth | On Earth
thumb|upright|Spodumene, an important lithium mineral
The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the formation and evolution of the Solar System. In turn, the natural history of the Earth caused parts of this planet to have differing concentrations of the elements. The mass of the Earth is approximately 5.98 kg. It is composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to planetary differentiation, the core region is believed to be primarily composed of iron (88.8%), with smaller amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements.
The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are lithophiles and therefore remain close to the Earth's surface because they combine readily with oxygen and so associate strongly with silica, forming relatively low-density minerals that do not sink down into the Earth's core. Potassium, rubidium and caesium are also incompatible elements due to their large ionic radii.
Sodium and potassium are very abundant on Earth, both being among the ten most common elements in Earth's crust; sodium makes up approximately 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element overall and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element. Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater. Other solid deposits include halite, amphibole, cryolite, nitratine, and zeolite. Many of these solid deposits occur as a result of ancient seas evaporating, which still occurs now in places such as Utah's Great Salt Lake and the Dead Sea. Despite their near-equal abundance in Earth's crust, sodium is far more common than potassium in the ocean, both because potassium's larger size makes its salts less soluble, and because potassium is bound by silicates in soil and what potassium leaches is absorbed far more readily by plant life than sodium.
Despite its chemical similarity, lithium typically does not occur together with sodium or potassium due to its smaller size. Due to its relatively low reactivity, it can be found in seawater in large amounts; it is estimated that lithium concentration in seawater is approximately 0.14 to 0.25 parts per million (ppm) or 25 micromolar. Its diagonal relationship with magnesium often allows it to replace magnesium in ferromagnesium minerals, where its crustal concentration is about 18 ppm, comparable to that of gallium and niobium. Commercially, the most important lithium mineral is spodumene, which occurs in large deposits worldwide.
Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite, although none of these contain only rubidium and no other alkali metals. Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.
Francium-223, the only naturally occurring isotope of francium, is the product of the alpha decay of actinium-227 and can be found in trace amounts in uranium minerals. In a given sample of uranium, there is estimated to be only one francium atom for every 1018 uranium atoms. It has been calculated that there are at most 30 grams of francium in the earth's crust at any time, due to its extremely short half-life of 22 minutes. |
Alkali metal | Properties | Properties |
Alkali metal | Physical and chemical | Physical and chemical
The physical and chemical properties of the alkali metals can be readily explained by their having an ns1 valence electron configuration, which results in weak metallic bonding. Hence, all the alkali metals are soft and have low densities, melting and boiling points, as well as heats of sublimation, vaporisation, and dissociation. They all crystallise in the body-centered cubic crystal structure, and have distinctive flame colours because their outer s electron is very easily excited. Indeed, these flame test colours are the most common way of identifying them since all their salts with common ions are soluble. The ns1 configuration also results in the alkali metals having very large atomic and ionic radii, as well as very high thermal and electrical conductivity. Their chemistry is dominated by the loss of their lone valence electron in the outermost s-orbital to form the +1 oxidation state, due to the ease of ionising this electron and the very high second ionisation energy. Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not well established due to its extreme radioactivity; thus, the presentation of its properties here is limited. What little is known about francium shows that it is very close in behaviour to caesium, as expected. The physical properties of francium are even sketchier because the bulk element has never been observed; hence any data that may be found in the literature are certainly speculative extrapolations.
+ Properties of the alkali metals Name Lithium Sodium Potassium Rubidium Caesium FranciumAtomic number 3 11 19 37 55 87Standard atomic weight 6.94(1) 22.98976928(2) 39.0983(1) 85.4678(3) 132.9054519(2) [223]Electron configuration [He] 2s1 [Ne] 3s1 [Ar] 4s1 [Kr] 5s1 [Xe] 6s1 [Rn] 7s1Melting point (°C) 180.54 97.72 63.38 39.31 28.44 ?Boiling point (°C) 1342 883 759 688 671 ?Density (g·cm−3) 0.534 0.968 0.89 1.532 1.93 ?Heat of fusion (kJ·mol−1) 3.00 2.60 2.321 2.19 2.09 ?Heat of vaporisation (kJ·mol−1) 136 97.42 79.1 69 66.1 ?Heat of formation of monatomic gas (kJ·mol−1) 162 108 89.6 82.0 78.2 ?Electrical resistivity at 25 °C (nΩ·cm) 94.7 48.8 73.9 131 208 ?Atomic radius (pm) 152 186 227 248 265 ?Ionic radius of hexacoordinate M+ ion (pm) 76 102 138 152 167 ?First ionisation energy (kJ·mol−1) 520.2 495.8 418.8 403.0 375.7 392.8Electron affinity (kJ·mol−1) 59.62 52.87 48.38 46.89 45.51 ?Enthalpy of dissociation of M2 (kJ·mol−1) 106.5 73.6 57.3 45.6 44.77 ?Pauling electronegativity 0.98 0.93 0.82 0.82 0.79 ?Allen electronegativity0.910.870.730.710.660.67Standard electrode potential (E°(M+→M0); V)Vanýsek, Petr (2011). “Electrochemical Series”, in Handbook of Chemistry and Physics: 92nd Edition (Chemical Rubber Company). −3.04 −2.71 −2.93 −2.98 −3.03 ?Flame test colourPrincipal emission/absorption wavelength (nm) Crimson670.8 Yellow589.2 Violet766.5 Red-violet780.0 Blue455.5 ?
The alkali metals are more similar to each other than the elements in any other group are to each other. Indeed, the similarity is so great that it is quite difficult to separate potassium, rubidium, and caesium, due to their similar ionic radii; lithium and sodium are more distinct. For instance, when moving down the table, all known alkali metals show increasing atomic radius, decreasing electronegativity, increasing reactivity, and decreasing melting and boiling points as well as heats of fusion and vaporisation. In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium. One of the very few properties of the alkali metals that does not display a very smooth trend is their reduction potentials: lithium's value is anomalous, being more negative than the others. This is because the Li+ ion has a very high hydration energy in the gas phase: though the lithium ion disrupts the structure of water significantly, causing a higher change in entropy, this high hydration energy is enough to make the reduction potentials indicate it as being the most electropositive alkali metal, despite the difficulty of ionising it in the gas phase.
The stable alkali metals are all silver-coloured metals except for caesium, which has a pale golden tint: it is one of only three metals that are clearly coloured (the other two being copper and gold). Additionally, the heavy alkaline earth metals calcium, strontium, and barium, as well as the divalent lanthanides europium and ytterbium, are pale yellow, though the colour is much less prominent than it is for caesium. Their lustre tarnishes rapidly in air due to oxidation.
thumb|right|Potassium reacts violently with water at room temperature
thumb|right|Caesium reacts explosively with water even at low temperatures
All the alkali metals are highly reactive and are never found in elemental forms in nature. Because of this, they are usually stored in mineral oil or kerosene (paraffin oil). They react aggressively with the halogens to form the alkali metal halides, which are white ionic crystalline compounds that are all soluble in water except lithium fluoride (LiF). The alkali metals also react with water to form strongly alkaline hydroxides and thus should be handled with great care. The heavier alkali metals react more vigorously than the lighter ones; for example, when dropped into water, caesium produces a larger explosion than potassium if the same number of moles of each metal is used. The alkali metals have the lowest first ionisation energies in their respective periods of the periodic table because of their low effective nuclear charge and the ability to attain a noble gas configuration by losing just one electron. Not only do the alkali metals react with water, but also with proton donors like alcohols and phenols, gaseous ammonia, and alkynes, the last demonstrating the phenomenal degree of their reactivity. Their great power as reducing agents makes them very useful in liberating other metals from their oxides or halides.
The second ionisation energy of all of the alkali metals is very high as it is in a full shell that is also closer to the nucleus; thus, they almost always lose a single electron, forming cations. The alkalides are an exception: they are unstable compounds which contain alkali metals in a −1 oxidation state, which is very unusual as before the discovery of the alkalides, the alkali metals were not expected to be able to form anions and were thought to be able to appear in salts only as cations. The alkalide anions have filled s-subshells, which gives them enough stability to exist. All the stable alkali metals except lithium are known to be able to form alkalides, and the alkalides have much theoretical interest due to their unusual stoichiometry and low ionisation potentials. Alkalides are chemically similar to the electrides, which are salts with trapped electrons acting as anions. A particularly striking example of an alkalide is "inverse sodium hydride", H+Na− (both ions being complexed), as opposed to the usual sodium hydride, Na+H−: it is unstable in isolation, due to its high energy resulting from the displacement of two electrons from hydrogen to sodium, although several derivatives are predicted to be metastable or stable.
In aqueous solution, the alkali metal ions form aqua ions of the formula [M(H2O)n]+, where n is the solvation number. Their coordination numbers and shapes agree well with those expected from their ionic radii. In aqueous solution the water molecules directly attached to the metal ion are said to belong to the first coordination sphere, also known as the first, or primary, solvation shell. The bond between a water molecule and the metal ion is a dative covalent bond, with the oxygen atom donating both electrons to the bond. Each coordinated water molecule may be attached by hydrogen bonds to other water molecules. The latter are said to reside in the second coordination sphere. However, for the alkali metal cations, the second coordination sphere is not well-defined as the +1 charge on the cation is not high enough to polarise the water molecules in the primary solvation shell enough for them to form strong hydrogen bonds with those in the second coordination sphere, producing a more stable entity. The solvation number for Li+ has been experimentally determined to be 4, forming the tetrahedral [Li(H2O)4]+: while solvation numbers of 3 to 6 have been found for lithium aqua ions, solvation numbers less than 4 may be the result of the formation of contact ion pairs, and the higher solvation numbers may be interpreted in terms of water molecules that approach [Li(H2O)4]+ through a face of the tetrahedron, though molecular dynamic simulations may indicate the existence of an octahedral hexaaqua ion. There are also probably six water molecules in the primary solvation sphere of the sodium ion, forming the octahedral [Na(H2O)6]+ ion. While it was previously thought that the heavier alkali metals also formed octahedral hexaaqua ions, it has since been found that potassium and rubidium probably form the [K(H2O)8]+ and [Rb(H2O)8]+ ions, which have the square antiprismatic structure, and that caesium forms the 12-coordinate [Cs(H2O)12]+ ion. |
Alkali metal | Lithium | Lithium
The chemistry of lithium shows several differences from that of the rest of the group as the small Li+ cation polarises anions and gives its compounds a more covalent character. Lithium and magnesium have a diagonal relationship due to their similar atomic radii, so that they show some similarities. For example, lithium forms a stable nitride, a property common among all the alkaline earth metals (magnesium's group) but unique among the alkali metals. In addition, among their respective groups, only lithium and magnesium form organometallic compounds with significant covalent character (e.g. LiMe and MgMe2).
Lithium fluoride is the only alkali metal halide that is poorly soluble in water, and lithium hydroxide is the only alkali metal hydroxide that is not deliquescent. Conversely, lithium perchlorate and other lithium salts with large anions that cannot be polarised are much more stable than the analogous compounds of the other alkali metals, probably because Li+ has a high solvation energy. This effect also means that most simple lithium salts are commonly encountered in hydrated form, because the anhydrous forms are extremely hygroscopic: this allows salts like lithium chloride and lithium bromide to be used in dehumidifiers and air-conditioners. |
Alkali metal | Francium | Francium
Francium is also predicted to show some differences due to its high atomic weight, causing its electrons to travel at considerable fractions of the speed of light and thus making relativistic effects more prominent. In contrast to the trend of decreasing electronegativities and ionisation energies of the alkali metals, francium's electronegativity and ionisation energy are predicted to be higher than caesium's due to the relativistic stabilisation of the 7s electrons; also, its atomic radius is expected to be abnormally low. Thus, contrary to expectation, caesium is the most reactive of the alkali metals, not francium. All known physical properties of francium also deviate from the clear trends going from lithium to caesium, such as the first ionisation energy, electron affinity, and anion polarisability, though due to the paucity of known data about francium many sources give extrapolated values, ignoring that relativistic effects make the trend from lithium to caesium become inapplicable at francium. Some of the few properties of francium that have been predicted taking relativity into account are the electron affinity (47.2 kJ/mol) and the enthalpy of dissociation of the Fr2 molecule (42.1 kJ/mol). The CsFr molecule is polarised as Cs+Fr−, showing that the 7s subshell of francium is much more strongly affected by relativistic effects than the 6s subshell of caesium. Additionally, francium superoxide (FrO2) is expected to have significant covalent character, unlike the other alkali metal superoxides, because of bonding contributions from the 6p electrons of francium. |
Alkali metal | Nuclear | Nuclear
+Primordial isotopes of the alkali metals Z Alkali metal Stable Decaysunstable: italicsodd–odd isotopes coloured pink 3 lithium 2 — 11 sodium 1 — 19 potassium 2 1 37 rubidium 1 1 55 caesium 1 — 87 francium — — No primordial isotopes( is a radiogenic nuclide)Radioactive:
All the alkali metals have odd atomic numbers; hence, their isotopes must be either odd–odd (both proton and neutron number are odd) or odd–even (proton number is odd, but neutron number is even). Odd–odd nuclei have even mass numbers, whereas odd–even nuclei have odd mass numbers. Odd–odd primordial nuclides are rare because most odd–odd nuclei are highly unstable with respect to beta decay, because the decay products are even–even, and are therefore more strongly bound, due to nuclear pairing effects.
Due to the great rarity of odd–odd nuclei, almost all the primordial isotopes of the alkali metals are odd–even (the exceptions being the light stable isotope lithium-6 and the long-lived radioisotope potassium-40). For a given odd mass number, there can be only a single beta-stable nuclide, since there is not a difference in binding energy between even–odd and odd–even comparable to that between even–even and odd–odd, leaving other nuclides of the same mass number (isobars) free to beta decay toward the lowest-mass nuclide. An effect of the instability of an odd number of either type of nucleons is that odd-numbered elements, such as the alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 monoisotopic elements that have only a single stable isotope, all but one have an odd atomic number and all but one also have an even number of neutrons. Beryllium is the single exception to both rules, due to its low atomic number.
All of the alkali metals except lithium and caesium have at least one naturally occurring radioisotope: sodium-22 and sodium-24 are trace radioisotopes produced cosmogenically, potassium-40 and rubidium-87 have very long half-lives and thus occur naturally, and all isotopes of francium are radioactive. Caesium was also thought to be radioactive in the early 20th century, although it has no naturally occurring radioisotopes. (Francium had not been discovered yet at that time.) The natural long-lived radioisotope of potassium, potassium-40, makes up about 0.012% of natural potassium, and thus natural potassium is weakly radioactive. This natural radioactivity became a basis for a mistaken claim of the discovery for element 87 (the next alkali metal after caesium) in 1925. Natural rubidium is similarly slightly radioactive, with 27.83% being the long-lived radioisotope rubidium-87.
Caesium-137, with a half-life of 30.17 years, is one of the two principal medium-lived fission products, along with strontium-90, which are responsible for most of the radioactivity of spent nuclear fuel after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left from the Chernobyl accident. Caesium-137 undergoes high-energy beta decay and eventually becomes stable barium-137. It is a strong emitter of gamma radiation. Caesium-137 has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay. Caesium-137 has been used as a tracer in hydrologic studies, analogous to the use of tritium.Radioisotope Brief: Cesium-137 (Cs-137). U.S. National Center for Environmental Health Small amounts of caesium-134 and caesium-137 were released into the environment during nearly all nuclear weapon tests and some nuclear accidents, most notably the Goiânia accident and the Chernobyl disaster. As of 2005, caesium-137 is the principal source of radiation in the zone of alienation around the Chernobyl nuclear power plant. Its chemical properties as one of the alkali metals make it one of the most problematic of the short-to-medium-lifetime fission products because it easily moves and spreads in nature due to the high water solubility of its salts, and is taken up by the body, which mistakes it for its essential congeners sodium and potassium. |
Alkali metal | Periodic trends | Periodic trends
The alkali metals are more similar to each other than the elements in any other group are to each other. For instance, when moving down the table, all known alkali metals show increasing atomic radius, decreasing electronegativity, increasing reactivity, and decreasing melting and boiling points as well as heats of fusion and vaporisation. In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium. |
Alkali metal | Atomic and ionic radii | Atomic and ionic radii
thumb|250px|Effective nuclear charge on an atomic electron
The atomic radii of the alkali metals increase going down the group. Because of the shielding effect, when an atom has more than one electron shell, each electron feels electric repulsion from the other electrons as well as electric attraction from the nucleus. In the alkali metals, the outermost electron only feels a net charge of +1, as some of the nuclear charge (which is equal to the atomic number) is cancelled by the inner electrons; the number of inner electrons of an alkali metal is always one less than the nuclear charge. Therefore, the only factor which affects the atomic radius of the alkali metals is the number of electron shells. Since this number increases down the group, the atomic radius must also increase down the group.
The ionic radii of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different electron shell than the inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the effective nuclear charge has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases. |
Alkali metal | First ionisation energy | First ionisation energy
thumb|upright=2.7|Periodic trend for ionisation energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases. Predicted values are used for elements beyond 104.
The first ionisation energy of an element or molecule is the energy required to move the most loosely held electron from one mole of gaseous atoms of the element or molecules to form one mole of gaseous ions with electric charge +1. The factors affecting the first ionisation energy are the nuclear charge, the amount of shielding by the inner electrons and the distance from the most loosely held electron from the nucleus, which is always an outer electron in main group elements. The first two factors change the effective nuclear charge the most loosely held electron feels. Since the outermost electron of alkali metals always feels the same effective nuclear charge (+1), the only factor which affects the first ionisation energy is the distance from the outermost electron to the nucleus. Since this distance increases down the group, the outermost electron feels less attraction from the nucleus and thus the first ionisation energy decreases. This trend is broken in francium due to the relativistic stabilisation and contraction of the 7s orbital, bringing francium's valence electron closer to the nucleus than would be expected from non-relativistic calculations. This makes francium's outermost electron feel more attraction from the nucleus, increasing its first ionisation energy slightly beyond that of caesium.
The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled electron shell and is thus difficult to remove. |
Alkali metal | Reactivity | Reactivity
The reactivities of the alkali metals increase going down the group. This is the result of a combination of two factors: the first ionisation energies and atomisation energies of the alkali metals. Because the first ionisation energy of the alkali metals decreases down the group, it is easier for the outermost electron to be removed from the atom and participate in chemical reactions, thus increasing reactivity down the group. The atomisation energy measures the strength of the metallic bond of an element, which falls down the group as the atoms increase in radius and thus the metallic bond must increase in length, making the delocalised electrons further away from the attraction of the nuclei of the heavier alkali metals. Adding the atomisation and first ionisation energies gives a quantity closely related to (but not equal to) the activation energy of the reaction of an alkali metal with another substance. This quantity decreases going down the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity increases down the group. |
Alkali metal | Electronegativity | Electronegativity
thumb|upright=1.25|Periodic variation of Pauling electronegativities as one descends the main groups of the periodic table from the second to the sixth period.
Electronegativity is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself. If the bond between sodium and chlorine in sodium chloride were covalent, the pair of shared electrons would be attracted to the chlorine because the effective nuclear charge on the outer electrons is +7 in chlorine but is only +1 in sodium. The electron pair is attracted so close to the chlorine atom that they are practically transferred to the chlorine atom (an ionic bond). However, if the sodium atom was replaced by a lithium atom, the electrons will not be attracted as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more strongly attracted to the closer effective nuclear charge from lithium. Hence, the larger alkali metal atoms (further down the group) will be less electronegative as the bonding pair is less strongly attracted towards them. As mentioned previously, francium is expected to be an exception.
Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example, lithium iodide (LiI) will dissolve in organic solvents, a property of most covalent compounds. Lithium fluoride (LiF) is the only alkali halide that is not soluble in water, and lithium hydroxide (LiOH) is the only alkali metal hydroxide that is not deliquescent. |
Alkali metal | Melting and boiling points | Melting and boiling points
The melting point of a substance is the point where it changes state from solid to liquid while the boiling point of a substance (in liquid state) is the point where the vapour pressure of the liquid equals the environmental pressure surrounding the liquid Section 17.43, page 321 and all the liquid changes state to gas. As a metal is heated to its melting point, the metallic bonds keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal's boiling point. Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali metals decreases down the group. This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons. As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points. The increased nuclear charge is not a relevant factor due to the shielding effect. |
Alkali metal | Density | Density
The alkali metals all have the same crystal structure (body-centred cubic) and thus the only relevant factors are the number of atoms that can fit into a certain volume and the mass of one of the atoms, since density is defined as mass per unit volume. The first factor depends on the volume of the atom and thus the atomic radius, which increases going down the group; thus, the volume of an alkali metal atom increases going down the group. The mass of an alkali metal atom also increases going down the group. Thus, the trend for the densities of the alkali metals depends on their atomic weights and atomic radii; if figures for these two factors are known, the ratios between the densities of the alkali metals can then be calculated. The resultant trend is that the densities of the alkali metals increase down the table, with an exception at potassium. Due to having the lowest atomic weight and the largest atomic radius of all the elements in their periods, the alkali metals are the least dense metals in the periodic table. Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water: in fact, lithium is the least dense known solid at room temperature. |
Alkali metal | Compounds | Compounds
The alkali metals form complete series of compounds with all usually encountered anions, which well illustrate group trends. These compounds can be described as involving the alkali metals losing electrons to acceptor species and forming monopositive ions. This description is most accurate for alkali halides and becomes less and less accurate as cationic and anionic charge increase, and as the anion becomes larger and more polarisable. For instance, ionic bonding gives way to metallic bonding along the series NaCl, Na2O, Na2S, Na3P, Na3As, Na3Sb, Na3Bi, Na. |
Alkali metal | [[Hydroxides]] | Hydroxides
thumb|right|alt=A large orange-yellow explosion|A reaction of 3 pounds (≈ 1.4 kg) of sodium with water
All the alkali metals react vigorously or explosively with cold water, producing an aqueous solution of a strongly basic alkali metal hydroxide and releasing hydrogen gas. This reaction becomes more vigorous going down the group: lithium reacts steadily with effervescence, but sodium and potassium can ignite, and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers. When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing hydrogen gas; this takes place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water). The alkali metal hydroxides are the most basic known hydroxides.
Recent research has suggested that the explosive behavior of alkali metals in water is driven by a Coulomb explosion rather than solely by rapid generation of hydrogen itself. All alkali metals melt as a part of the reaction with water. Water molecules ionise the bare metallic surface of the liquid metal, leaving a positively charged metal surface and negatively charged water ions. The attraction between the charged metal and water ions will rapidly increase the surface area, causing an exponential increase of ionisation. When the repulsive forces within the liquid metal surface exceeds the forces of the surface tension, it vigorously explodes.
The hydroxides themselves are the most basic hydroxides known, reacting with acids to give salts and with alcohols to give oligomeric alkoxides. They easily react with carbon dioxide to form carbonates or bicarbonates, or with hydrogen sulfide to form sulfides or bisulfides, and may be used to separate thiols from petroleum. They react with amphoteric oxides: for example, the oxides of aluminium, zinc, tin, and lead react with the alkali metal hydroxides to give aluminates, zincates, stannates, and plumbates. Silicon dioxide is acidic, and thus the alkali metal hydroxides can also attack silicate glass. |
Alkali metal | Intermetallic compounds | Intermetallic compounds
thumb|right|Liquid NaK alloy at room temperature
The alkali metals form many intermetallic compounds with each other and the elements from groups 2 to 13 in the periodic table of varying stoichiometries, such as the sodium amalgams with mercury, including Na5Hg8 and Na3Hg. Some of these have ionic characteristics: taking the alloys with gold, the most electronegative of metals, as an example, NaAu and KAu are metallic, but RbAu and CsAu are semiconductors. NaK is an alloy of sodium and potassium that is very useful because it is liquid at room temperature, although precautions must be taken due to its extreme reactivity towards water and air. The eutectic mixture melts at −12.6 °C. An alloy of 41% caesium, 47% sodium, and 12% potassium has the lowest known melting point of any metal or alloy, −78 °C. |
Alkali metal | Compounds with the group 13 elements | Compounds with the group 13 elements
The intermetallic compounds of the alkali metals with the heavier group 13 elements (aluminium, gallium, indium, and thallium), such as NaTl, are poor conductors or semiconductors, unlike the normal alloys with the preceding elements, implying that the alkali metal involved has lost an electron to the Zintl anions involved. Nevertheless, while the elements in group 14 and beyond tend to form discrete anionic clusters, group 13 elements tend to form polymeric ions with the alkali metal cations located between the giant ionic lattice. For example, NaTl consists of a polymeric anion (—Tl−—)n with a covalent diamond cubic structure with Na+ ions located between the anionic lattice. The larger alkali metals cannot fit similarly into an anionic lattice and tend to force the heavier group 13 elements to form anionic clusters.S.M. Kauzlarich, Encyclopedia of Inorganic chemistry, 1994, John Wiley & Sons,
Boron is a special case, being the only nonmetal in group 13. The alkali metal borides tend to be boron-rich, involving appreciable boron–boron bonding involving deltahedral structures, and are thermally unstable due to the alkali metals having a very high vapour pressure at elevated temperatures. This makes direct synthesis problematic because the alkali metals do not react with boron below 700 °C, and thus this must be accomplished in sealed containers with the alkali metal in excess. Furthermore, exceptionally in this group, reactivity with boron decreases down the group: lithium reacts completely at 700 °C, but sodium at 900 °C and potassium not until 1200 °C, and the reaction is instantaneous for lithium but takes hours for potassium. Rubidium and caesium borides have not even been characterised. Various phases are known, such as LiB10, NaB6, NaB15, and KB6. Under high pressure the boron–boron bonding in the lithium borides changes from following Wade's rules to forming Zintl anions like the rest of group 13. |
Alkali metal | Compounds with the group 14 elements | Compounds with the group 14 elements
Lithium and sodium react with carbon to form acetylides, Li2C2 and Na2C2, which can also be obtained by reaction of the metal with acetylene. Potassium, rubidium, and caesium react with graphite; their atoms are intercalated between the hexagonal graphite layers, forming graphite intercalation compounds of formulae MC60 (dark grey, almost black), MC48 (dark grey, almost black), MC36 (blue), MC24 (steel blue), and MC8 (bronze) (M = K, Rb, or Cs). These compounds are over 200 times more electrically conductive than pure graphite, suggesting that the valence electron of the alkali metal is transferred to the graphite layers (e.g. ). Upon heating of KC8, the elimination of potassium atoms results in the conversion in sequence to KC24, KC36, KC48 and finally KC60. KC8 is a very strong reducing agent and is pyrophoric and explodes on contact with water.NIST Ionizing Radiation Division 2001 – Technical Highlights. physics.nist.gov While the larger alkali metals (K, Rb, and Cs) initially form MC8, the smaller ones initially form MC6, and indeed they require reaction of the metals with graphite at high temperatures around 500 °C to form. Apart from this, the alkali metals are such strong reducing agents that they can even reduce buckminsterfullerene to produce solid fullerides MnC60; sodium, potassium, rubidium, and caesium can form fullerides where n = 2, 3, 4, or 6, and rubidium and caesium additionally can achieve n = 1.
When the alkali metals react with the heavier elements in the carbon group (silicon, germanium, tin, and lead), ionic substances with cage-like structures are formed, such as the silicides M4Si4 (M = K, Rb, or Cs), which contains M+ and tetrahedral ions. The chemistry of alkali metal germanides, involving the germanide ion Ge4− and other cluster (Zintl) ions such as , , , and [(Ge9)2]6−, is largely analogous to that of the corresponding silicides. Alkali metal stannides are mostly ionic, sometimes with the stannide ion (Sn4−), and sometimes with more complex Zintl ions such as , which appears in tetrapotassium nonastannide (K4Sn9). The monatomic plumbide ion (Pb4−) is unknown, and indeed its formation is predicted to be energetically unfavourable; alkali metal plumbides have complex Zintl ions, such as . These alkali metal germanides, stannides, and plumbides may be produced by reducing germanium, tin, and lead with sodium metal in liquid ammonia. |
Alkali metal | Nitrides and pnictides | Nitrides and pnictides
thumb|Unit cell ball-and-stick model of lithium nitride. On the basis of size a tetrahedral structure would be expected, but that would be geometrically impossible: thus lithium nitride takes on this unique crystal structure.
Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with nitrogen at standard conditions, and its nitride is the only stable alkali metal nitride. Nitrogen is an unreactive gas because breaking the strong triple bond in the dinitrogen molecule (N2) requires a lot of energy. The formation of an alkali metal nitride would consume the ionisation energy of the alkali metal (forming M+ ions), the energy required to break the triple bond in N2 and the formation of N3− ions, and all the energy released from the formation of an alkali metal nitride is from the lattice energy of the alkali metal nitride. The lattice energy is maximised with small, highly charged ions; the alkali metals do not form highly charged ions, only forming ions with a charge of +1, so only lithium, the smallest alkali metal, can release enough lattice energy to make the reaction with nitrogen exothermic, forming lithium nitride. The reactions of the other alkali metals with nitrogen would not release enough lattice energy and would thus be endothermic, so they do not form nitrides at standard conditions. Sodium nitride (Na3N) and potassium nitride (K3N), while existing, are extremely unstable, being prone to decomposing back into their constituent elements, and cannot be produced by reacting the elements with each other at standard conditions.. 'Elusive Binary Compound Prepared' Chemical & Engineering News 80 No. 20 (20 May 2002) Steric hindrance forbids the existence of rubidium or caesium nitride. However, sodium and potassium form colourless azide salts involving the linear anion; due to the large size of the alkali metal cations, they are thermally stable enough to be able to melt before decomposing.
All the alkali metals react readily with phosphorus and arsenic to form phosphides and arsenides with the formula M3Pn (where M represents an alkali metal and Pn represents a pnictogen – phosphorus, arsenic, antimony, or bismuth). This is due to the greater size of the P3− and As3− ions, so that less lattice energy needs to be released for the salts to form. These are not the only phosphides and arsenides of the alkali metals: for example, potassium has nine different known phosphides, with formulae K3P, K4P3, K5P4, KP, K4P6, K3P7, K3P11, KP10.3, and KP15.H.G. Von Schnering, W. Hönle Phosphides – Solid-state Chemistry Encyclopedia of Inorganic Chemistry Ed. R. Bruce King (1994) John Wiley & Sons While most metals form arsenides, only the alkali and alkaline earth metals form mostly ionic arsenides. The structure of Na3As is complex with unusually short Na–Na distances of 328–330 pm which are shorter than in sodium metal, and this indicates that even with these electropositive metals the bonding cannot be straightforwardly ionic. Other alkali metal arsenides not conforming to the formula M3As are known, such as LiAs, which has a metallic lustre and electrical conductivity indicating the presence of some metallic bonding. The antimonides are unstable and reactive as the Sb3− ion is a strong reducing agent; reaction of them with acids form the toxic and unstable gas stibine (SbH3). Indeed, they have some metallic properties, and the alkali metal antimonides of stoichiometry MSb involve antimony atoms bonded in a spiral Zintl structure. Bismuthides are not even wholly ionic; they are intermetallic compounds containing partially metallic and partially ionic bonds. |
Alkali metal | Oxides and chalcogenides | Oxides and chalcogenides
All the alkali metals react vigorously with oxygen at standard conditions. They form various types of oxides, such as simple oxides (containing the O2− ion), peroxides (containing the ion, where there is a single bond between the two oxygen atoms), superoxides (containing the ion), and many others. Lithium burns in air to form lithium oxide, but sodium reacts with oxygen to form a mixture of sodium oxide and sodium peroxide. Potassium forms a mixture of potassium peroxide and potassium superoxide, while rubidium and caesium form the superoxide exclusively. Their reactivity increases going down the group: while lithium, sodium and potassium merely burn in air, rubidium and caesium are pyrophoric (spontaneously catch fire in air).
The smaller alkali metals tend to polarise the larger anions (the peroxide and superoxide) due to their small size. This attracts the electrons in the more complex anions towards one of its constituent oxygen atoms, forming an oxide ion and an oxygen atom. This causes lithium to form the oxide exclusively on reaction with oxygen at room temperature. This effect becomes drastically weaker for the larger sodium and potassium, allowing them to form the less stable peroxides. Rubidium and caesium, at the bottom of the group, are so large that even the least stable superoxides can form. Because the superoxide releases the most energy when formed, the superoxide is preferentially formed for the larger alkali metals where the more complex anions are not polarised. The oxides and peroxides for these alkali metals do exist, but do not form upon direct reaction of the metal with oxygen at standard conditions. In addition, the small size of the Li+ and O2− ions contributes to their forming a stable ionic lattice structure. Under controlled conditions, however, all the alkali metals, with the exception of francium, are known to form their oxides, peroxides, and superoxides. The alkali metal peroxides and superoxides are powerful oxidising agents. Sodium peroxide and potassium superoxide react with carbon dioxide to form the alkali metal carbonate and oxygen gas, which allows them to be used in submarine air purifiers; the presence of water vapour, naturally present in breath, makes the removal of carbon dioxide by potassium superoxide even more efficient. All the stable alkali metals except lithium can form red ozonides (MO3) through low-temperature reaction of the powdered anhydrous hydroxide with ozone: the ozonides may be then extracted using liquid ammonia. They slowly decompose at standard conditions to the superoxides and oxygen, and hydrolyse immediately to the hydroxides when in contact with water. Potassium, rubidium, and caesium also form sesquioxides M2O3, which may be better considered peroxide disuperoxides, .
Rubidium and caesium can form a great variety of suboxides with the metals in formal oxidation states below +1. Rubidium can form Rb6O and Rb9O2 (copper-coloured) upon oxidation in air, while caesium forms an immense variety of oxides, such as the ozonide CsO3 and several brightly coloured suboxides, such as Cs7O (bronze), Cs4O (red-violet), Cs11O3 (violet), Cs3O (dark green), CsO, Cs3O2, as well as Cs7O2. The last of these may be heated under vacuum to generate Cs2O.
The alkali metals can also react analogously with the heavier chalcogens (sulfur, selenium, tellurium, and polonium), and all the alkali metal chalcogenides are known (with the exception of francium's). Reaction with an excess of the chalcogen can similarly result in lower chalcogenides, with chalcogen ions containing chains of the chalcogen atoms in question. For example, sodium can react with sulfur to form the sulfide (Na2S) and various polysulfides with the formula Na2Sx (x from 2 to 6), containing the ions. Due to the basicity of the Se2− and Te2− ions, the alkali metal selenides and tellurides are alkaline in solution; when reacted directly with selenium and tellurium, alkali metal polyselenides and polytellurides are formed along with the selenides and tellurides with the and ions. They may be obtained directly from the elements in liquid ammonia or when air is not present, and are colourless, water-soluble compounds that air oxidises quickly back to selenium or tellurium. The alkali metal polonides are all ionic compounds containing the Po2− ion; they are very chemically stable and can be produced by direct reaction of the elements at around 300–400 °C. |
Alkali metal | Halides, hydrides, and pseudohalides | Halides, hydrides, and pseudohalides
The alkali metals are among the most electropositive elements on the periodic table and thus tend to bond ionically to the most electronegative elements on the periodic table, the halogens (fluorine, chlorine, bromine, iodine, and astatine), forming salts known as the alkali metal halides. The reaction is very vigorous and can sometimes result in explosions. All twenty stable alkali metal halides are known; the unstable ones are not known, with the exception of sodium astatide, because of the great instability and rarity of astatine and francium. The most well-known of the twenty is certainly sodium chloride, otherwise known as common salt. All of the stable alkali metal halides have the formula MX where M is an alkali metal and X is a halogen. They are all white ionic crystalline solids that have high melting points. All the alkali metal halides are soluble in water except for lithium fluoride (LiF), which is insoluble in water due to its very high lattice enthalpy. The high lattice enthalpy of lithium fluoride is due to the small sizes of the Li+ and F− ions, causing the electrostatic interactions between them to be strong: a similar effect occurs for magnesium fluoride, consistent with the diagonal relationship between lithium and magnesium.
The alkali metals also react similarly with hydrogen to form ionic alkali metal hydrides, where the hydride anion acts as a pseudohalide: these are often used as reducing agents, producing hydrides, complex metal hydrides, or hydrogen gas. Other pseudohalides are also known, notably the cyanides. These are isostructural to the respective halides except for lithium cyanide, indicating that the cyanide ions may rotate freely. Ternary alkali metal halide oxides, such as Na3ClO, K3BrO (yellow), Na4Br2O, Na4I2O, and K4Br2O, are also known. The polyhalides are rather unstable, although those of rubidium and caesium are greatly stabilised by the feeble polarising power of these extremely large cations. |
Alkali metal | Coordination complexes | Coordination complexes
Alkali metal cations do not usually form coordination complexes with simple Lewis bases due to their low charge of just +1 and their relatively large size; thus the Li+ ion forms most complexes and the heavier alkali metal ions form less and less (though exceptions occur for weak complexes). Lithium in particular has a very rich coordination chemistry in which it exhibits coordination numbers from 1 to 12, although octahedral hexacoordination is its preferred mode. In aqueous solution, the alkali metal ions exist as octahedral hexahydrate complexes [M(H2O)6]+, with the exception of the lithium ion, which due to its small size forms tetrahedral tetrahydrate complexes [Li(H2O)4]+; the alkali metals form these complexes because their ions are attracted by electrostatic forces of attraction to the polar water molecules. Because of this, anhydrous salts containing alkali metal cations are often used as desiccants. Alkali metals also readily form complexes with crown ethers (e.g. 12-crown-4 for Li+, 15-crown-5 for Na+, 18-crown-6 for K+, and 21-crown-7 for Rb+) and cryptands due to electrostatic attraction. |
Alkali metal | Ammonia solutions | Ammonia solutions
The alkali metals dissolve slowly in liquid ammonia, forming ammoniacal solutions of solvated metal cation M+ and solvated electron e−, which react to form hydrogen gas and the alkali metal amide (MNH2, where M represents an alkali metal): this was first noted by Humphry Davy in 1809 and rediscovered by W. Weyl in 1864. The process may be speeded up by a catalyst. Similar solutions are formed by the heavy divalent alkaline earth metals calcium, strontium, barium, as well as the divalent lanthanides, europium and ytterbium. The amide salt is quite insoluble and readily precipitates out of solution, leaving intensely coloured ammonia solutions of the alkali metals. In 1907, Charles A. Kraus identified the colour as being due to the presence of solvated electrons, which contribute to the high electrical conductivity of these solutions. At low concentrations (below 3 M), the solution is dark blue and has ten times the conductivity of aqueous sodium chloride; at higher concentrations (above 3 M), the solution is copper-coloured and has approximately the conductivity of liquid metals like mercury. In addition to the alkali metal amide salt and solvated electrons, such ammonia solutions also contain the alkali metal cation (M+), the neutral alkali metal atom (M), diatomic alkali metal molecules (M2) and alkali metal anions (M−). These are unstable and eventually become the more thermodynamically stable alkali metal amide and hydrogen gas. Solvated electrons are powerful reducing agents and are often used in chemical synthesis. |
Alkali metal | Organometallic | Organometallic |
Alkali metal | Organolithium | Organolithium
thumb|upright=1.15|Structure of the octahedral n-butyllithium hexamer, (C4H9Li)6. The aggregates are held together by delocalised covalent bonds between lithium and the terminal carbon of the butyl chain.Elschenbroich, C. "Organometallics" (2006) Wiley-VCH: Weinheim. . There is no direct lithium–lithium bonding in any organolithium compound.
thumb|upright=1.15|Solid phenyllithium forms monoclinic crystals that can be described as consisting of dimeric Li2(C6H5)2 subunits. The lithium atoms and the ipso carbons of the phenyl rings form a planar four-membered ring. The plane of the phenyl groups is perpendicular to the plane of this Li2C2 ring. Additional strong intermolecular bonding occurs between these phenyllithium dimers and the π electrons of the phenyl groups in the adjacent dimers, resulting in an infinite polymeric ladder structure.
Being the smallest alkali metal, lithium forms the widest variety of and most stable organometallic compounds, which are bonded covalently. Organolithium compounds are electrically non-conducting volatile solids or liquids that melt at low temperatures, and tend to form oligomers with the structure (RLi)x where R is the organic group. As the electropositive nature of lithium puts most of the charge density of the bond on the carbon atom, effectively creating a carbanion, organolithium compounds are extremely powerful bases and nucleophiles. For use as bases, butyllithiums are often used and are commercially available. An example of an organolithium compound is methyllithium ((CH3Li)x), which exists in tetrameric (x = 4, tetrahedral) and hexameric (x = 6, octahedral) forms. Organolithium compounds, especially n-butyllithium, are useful reagents in organic synthesis, as might be expected given lithium's diagonal relationship with magnesium, which plays an important role in the Grignard reaction. For example, alkyllithiums and aryllithiums may be used to synthesise aldehydes and ketones by reaction with metal carbonyls. The reaction with nickel tetracarbonyl, for example, proceeds through an unstable acyl nickel carbonyl complex which then undergoes electrophilic substitution to give the desired aldehyde (using H+ as the electrophile) or ketone (using an alkyl halide) product.
LiR \ + \ Ni(CO)4 \ \longrightarrow Li^{+}[RCONi(CO)3]^{-}
Li^{+}[RCONi(CO)3]^{-}->[\ce{H^{+}}][\ce{solvent}] \ Li^{+} \ + \ RCHO \ + \ [(solvent)Ni(CO)3]
Li^{+}[RCONi(CO)3]^{-}->[\ce{R^{'}Br}][\ce{solvent}] \ Li^{+} \ + \ RR^{'}CO \ + \ [(solvent)Ni(CO)3]
Alkyllithiums and aryllithiums may also react with N,N-disubstituted amides to give aldehydes and ketones, and symmetrical ketones by reacting with carbon monoxide. They thermally decompose to eliminate a β-hydrogen, producing alkenes and lithium hydride: another route is the reaction of ethers with alkyl- and aryllithiums that act as strong bases. In non-polar solvents, aryllithiums react as the carbanions they effectively are, turning carbon dioxide to aromatic carboxylic acids (ArCO2H) and aryl ketones to tertiary carbinols (Ar'2C(Ar)OH). Finally, they may be used to synthesise other organometallic compounds through metal-halogen exchange. |
Alkali metal | Heavier alkali metals | Heavier alkali metals
Unlike the organolithium compounds, the organometallic compounds of the heavier alkali metals are predominantly ionic. The application of organosodium compounds in chemistry is limited in part due to competition from organolithium compounds, which are commercially available and exhibit more convenient reactivity. The principal organosodium compound of commercial importance is sodium cyclopentadienide. Sodium tetraphenylborate can also be classified as an organosodium compound since in the solid state sodium is bound to the aryl groups. Organometallic compounds of the higher alkali metals are even more reactive than organosodium compounds and of limited utility. A notable reagent is Schlosser's base, a mixture of n-butyllithium and potassium tert-butoxide. This reagent reacts with propene to form the compound allylpotassium (KCH2CHCH2). cis-2-Butene and trans-2-butene equilibrate when in contact with alkali metals. Whereas isomerisation is fast with lithium and sodium, it is slow with the heavier alkali metals. The heavier alkali metals also favour the sterically congested conformation. Several crystal structures of organopotassium compounds have been reported, establishing that they, like the sodium compounds, are polymeric. Organosodium, organopotassium, organorubidium and organocaesium compounds are all mostly ionic and are insoluble (or nearly so) in nonpolar solvents.
Alkyl and aryl derivatives of sodium and potassium tend to react with air. They cause the cleavage of ethers, generating alkoxides. Unlike alkyllithium compounds, alkylsodiums and alkylpotassiums cannot be made by reacting the metals with alkyl halides because Wurtz coupling occurs:
RM + R'X → R–R' + MX
As such, they have to be made by reacting alkylmercury compounds with sodium or potassium metal in inert hydrocarbon solvents. While methylsodium forms tetramers like methyllithium, methylpotassium is more ionic and has the nickel arsenide structure with discrete methyl anions and potassium cations.
The alkali metals and their hydrides react with acidic hydrocarbons, for example cyclopentadienes and terminal alkynes, to give salts. Liquid ammonia, ether, or hydrocarbon solvents are used, the most common of which being tetrahydrofuran. The most important of these compounds is sodium cyclopentadienide, NaC5H5, an important precursor to many transition metal cyclopentadienyl derivatives. Similarly, the alkali metals react with cyclooctatetraene in tetrahydrofuran to give alkali metal cyclooctatetraenides; for example, dipotassium cyclooctatetraenide (K2C8H8) is an important precursor to many metal cyclooctatetraenyl derivatives, such as uranocene. The large and very weakly polarising alkali metal cations can stabilise large, aromatic, polarisable radical anions, such as the dark-green sodium naphthalenide, Na+[C10H8•]−, a strong reducing agent. |
Alkali metal | Representative reactions of alkali metals | Representative reactions of alkali metals |
Alkali metal | Reaction with oxygen | Reaction with oxygen
Upon reacting with oxygen, alkali metals form oxides, peroxides, superoxides and suboxides. However, the first three are more common. The table below"Inorganic Chemistry" by Gary L. Miessler and Donald A. Tar, 6th edition, Pearson shows the types of compounds formed in reaction with oxygen. The compound in brackets represents the minor product of combustion.
Alkali metalOxidePeroxideSuperoxideLiLi2O(Li2O2)Na(Na2O)Na2O2K KO2Rb RbO2Cs CsO2
The alkali metal peroxides are ionic compounds that are unstable in water. The peroxide anion is weakly bound to the cation, and it is hydrolysed, forming stronger covalent bonds.
Na2O2 + 2H2O → 2NaOH + H2O2
The other oxygen compounds are also unstable in water.
2KO2 + 2H2O → 2KOH + H2O2 + O2Kumar De, Anil (2007). A Text Book of Inorganic Chemistry. New Age International. p. 247. .
Li2O + H2O → 2LiOH |
Alkali metal | Reaction with sulfur | Reaction with sulfur
With sulfur, they form sulfides and polysulfides."The chemistry of the Elements" by Greenwood and Earnshaw, 2nd edition, Elsevier
2Na + 1/8S8 → Na2S + 1/8S8 → Na2S2...Na2S7
Because alkali metal sulfides are essentially salts of a weak acid and a strong base, they form basic solutions.
S2- + H2O → HS− + HO−
HS− + H2O → H2S + HO− |
Alkali metal | Reaction with nitrogen | Reaction with nitrogen
Lithium is the only metal that combines directly with nitrogen at room temperature.
3Li + 1/2N2 → Li3N
Li3N can react with water to liberate ammonia.
Li3N + 3H2O → 3LiOH + NH3 |
Alkali metal | Reaction with hydrogen | Reaction with hydrogen
With hydrogen, alkali metals form saline hydrides that hydrolyse in water.
2 Na \ + H2 \ ->[\ce{\Delta}] \ 2 NaH
2 NaH \ + \ 2 H2O \ \longrightarrow \ 2 NaOH \ + \ H2 \uparrow |
Alkali metal | Reaction with carbon | Reaction with carbon
Lithium is the only metal that reacts directly with carbon to give dilithium acetylide. Na and K can react with acetylene to give acetylides."Inorganic Chemistry" by Cotton and Wilkinson
2 Li \ + \ 2 C \ \longrightarrow \ Li2C2
2 Na \ + \ 2 C2H2 \ ->[\ce{150 \ ^{o}C}] \ 2 NaC2H \ + \ H2
2 Na \ + \ 2 NaC2H \ ->[\ce{220 \ ^{o}C}] \ 2 Na2C2 \ + \ H2 |
Alkali metal | Reaction with water | Reaction with water
On reaction with water, they generate hydroxide ions and hydrogen gas. This reaction is vigorous and highly exothermic and the hydrogen resulted may ignite in air or even explode in the case of Rb and Cs.
Na + H2O → NaOH + 1/2H2 |
Alkali metal | Reaction with other salts | Reaction with other salts
The alkali metals are very good reducing agents. They can reduce metal cations that are less electropositive. Titanium is produced industrially by the reduction of titanium tetrachloride with Na at 400 °C (van Arkel–de Boer process).
TiCl4 + 4Na → 4NaCl + Ti |
Alkali metal | Reaction with organohalide compounds | Reaction with organohalide compounds
Alkali metals react with halogen derivatives to generate hydrocarbon via the Wurtz reaction.
2CH3-Cl + 2Na → H3C-CH3 + 2NaCl |
Alkali metal | Alkali metals in liquid ammonia | Alkali metals in liquid ammonia
Alkali metals dissolve in liquid ammonia or other donor solvents like aliphatic amines or hexamethylphosphoramide to give blue solutions. These solutions are believed to contain free electrons.
Na + xNH3 → Na+ + e(NH3)x−
Due to the presence of solvated electrons, these solutions are very powerful reducing agents used in organic synthesis.
thumb|upright=1.25|centre|Reduction reactions using sodium in liquid ammonia
Reaction 1) is known as Birch reduction.
Other reductions that can be carried by these solutions are:
S8 + 2e− → S82-
Fe(CO)5 + 2e− → Fe(CO)42- + CO |
Alkali metal | Extensions | Extensions
thumb|upright=1.12|Empirical (Na–Cs, Mg–Ra) and predicted (Fr–Uhp, Ubn–Uhh) atomic radius of the alkali and alkaline earth metals from the third to the ninth period, measured in angstroms
Although francium is the heaviest alkali metal that has been discovered, there has been some theoretical work predicting the physical and chemical characteristics of hypothetical heavier alkali metals. Being the first period 8 element, the undiscovered element ununennium (element 119) is predicted to be the next alkali metal after francium and behave much like their lighter congeners; however, it is also predicted to differ from the lighter alkali metals in some properties. Its chemistry is predicted to be closer to that of potassium or rubidium instead of caesium or francium. This is unusual as periodic trends, ignoring relativistic effects would predict ununennium to be even more reactive than caesium and francium. This lowered reactivity is due to the relativistic stabilisation of ununennium's valence electron, increasing ununennium's first ionisation energy and decreasing the metallic and ionic radii; this effect is already seen for francium. This assumes that ununennium will behave chemically as an alkali metal, which, although likely, may not be true due to relativistic effects. The relativistic stabilisation of the 8s orbital also increases ununennium's electron affinity far beyond that of caesium and francium; indeed, ununennium is expected to have an electron affinity higher than all the alkali metals lighter than it. Relativistic effects also cause a very large drop in the polarisability of ununennium. On the other hand, ununennium is predicted to continue the trend of melting points decreasing going down the group, being expected to have a melting point between 0 °C and 30 °C.
thumb|left|Empirical (Na–Fr) and predicted (Uue) electron affinity of the alkali metals from the third to the eighth period, measured in electron volts
The stabilisation of ununennium's valence electron and thus the contraction of the 8s orbital cause its atomic radius to be lowered to 240 pm, very close to that of rubidium (247 pm), so that the chemistry of ununennium in the +1 oxidation state should be more similar to the chemistry of rubidium than to that of francium. On the other hand, the ionic radius of the Uue+ ion is predicted to be larger than that of Rb+, because the 7p orbitals are destabilised and are thus larger than the p-orbitals of the lower shells. Ununennium may also show the +3 and +5 oxidation states, which are not seen in any other alkali metal, in addition to the +1 oxidation state that is characteristic of the other alkali metals and is also the main oxidation state of all the known alkali metals: this is because of the destabilisation and expansion of the 7p3/2 spinor, causing its outermost electrons to have a lower ionisation energy than what would otherwise be expected. Indeed, many ununennium compounds are expected to have a large covalent character, due to the involvement of the 7p3/2 electrons in the bonding.
thumb|Empirical (Na–Fr, Mg–Ra) and predicted (Uue–Uhp, Ubn–Uhh) ionisation energy of the alkali and alkaline earth metals from the third to the ninth period, measured in electron volts
Not as much work has been done predicting the properties of the alkali metals beyond ununennium. Although a simple extrapolation of the periodic table (by the Aufbau principle) would put element 169, unhexennium, under ununennium, Dirac-Fock calculations predict that the next element after ununennium with alkali-metal-like properties may be element 165, unhexpentium, which is predicted to have the electron configuration [Og] 5g18 6f14 7d10 8s2 8p1/22 9s1. This element would be intermediate in properties between an alkali metal and a group 11 element, and while its physical and atomic properties would be closer to the former, its chemistry may be closer to that of the latter. Further calculations show that unhexpentium would follow the trend of increasing ionisation energy beyond caesium, having an ionisation energy comparable to that of sodium, and that it should also continue the trend of decreasing atomic radii beyond caesium, having an atomic radius comparable to that of potassium. However, the 7d electrons of unhexpentium may also be able to participate in chemical reactions along with the 9s electron, possibly allowing oxidation states beyond +1, whence the likely transition metal behaviour of unhexpentium. Due to the alkali and alkaline earth metals both being s-block elements, these predictions for the trends and properties of ununennium and unhexpentium also mostly hold quite similarly for the corresponding alkaline earth metals unbinilium (Ubn) and unhexhexium (Uhh). Unsepttrium, element 173, may be an even better heavier homologue of ununennium; with a predicted electron configuration of [Usb] 6g1, it returns to the alkali-metal-like situation of having one easily removed electron far above a closed p-shell in energy, and is expected to be even more reactive than caesium.
The probable properties of further alkali metals beyond unsepttrium have not been explored yet as of 2019, and they may or may not be able to exist. In periods 8 and above of the periodic table, relativistic and shell-structure effects become so strong that extrapolations from lighter congeners become completely inaccurate. In addition, the relativistic and shell-structure effects (which stabilise the s-orbitals and destabilise and expand the d-, f-, and g-orbitals of higher shells) have opposite effects, causing even larger difference between relativistic and non-relativistic calculations of the properties of elements with such high atomic numbers. Interest in the chemical properties of ununennium, unhexpentium, and unsepttrium stems from the fact that they are located close to the expected locations of islands of stability, centered at elements 122 (306Ubb) and 164 (482Uhq).Nuclear scientists eye future landfall on a second 'island of stability' . EurekAlert! (2008-04-06). Retrieved on 2016-11-25. |
Alkali metal | Pseudo-alkali metals | Pseudo-alkali metals
Many other substances are similar to the alkali metals in their tendency to form monopositive cations. Analogously to the pseudohalogens, they have sometimes been called "pseudo-alkali metals". These substances include some elements and many more polyatomic ions; the polyatomic ions are especially similar to the alkali metals in their large size and weak polarising power. |
Alkali metal | Hydrogen | Hydrogen
The element hydrogen, with one electron per neutral atom, is usually placed at the top of Group 1 of the periodic table because of its electron configuration. But hydrogen is not normally considered to be an alkali metal. Metallic hydrogen, which only exists at very high pressures, is known for its electrical and magnetic properties, not its chemical properties. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule (H2); however, the alkali metals form diatomic molecules (such as dilithium, Li2) only at high temperatures, when they are in the gaseous state.Winter, Mark J. (1994) Chemical Bonding, Oxford University Press,
Hydrogen, like the alkali metals, has one valence electron and reacts easily with the halogens, but the similarities mostly end there because of the small size of a bare proton H+ compared to the alkali metal cations. Its placement above lithium is primarily due to its electron configuration. It is sometimes placed above fluorine due to their similar chemical properties, though the resemblance is likewise not absolute.
The first ionisation energy of hydrogen (1312.0 kJ/mol) is much higher than that of the alkali metals.Huheey, J.E.; Keiter, E.A. and Keiter, R.L. (1993) Inorganic Chemistry: Principles of Structure and Reactivity, 4th edition, HarperCollins, New York, USA.James, A.M. and Lord, M.P. (1992) Macmillan's Chemical and Physical Data, Macmillan, London, UK. As only one additional electron is required to fill in the outermost shell of the hydrogen atom, hydrogen often behaves like a halogen, forming the negative hydride ion, and is very occasionally considered to be a halogen on that basis. (The alkali metals can also form negative ions, known as alkalides, but these are little more than laboratory curiosities, being unstable.) An argument against this placement is that formation of hydride from hydrogen is endothermic, unlike the exothermic formation of halides from halogens. The radius of the H− anion also does not fit the trend of increasing size going down the halogens: indeed, H− is very diffuse because its single proton cannot easily control both electrons. It was expected for some time that liquid hydrogen would show metallic properties; while this has been shown to not be the case, under extremely high pressures, such as those found at the cores of Jupiter and Saturn, hydrogen does become metallic and behaves like an alkali metal; in this phase, it is known as metallic hydrogen. The electrical resistivity of liquid metallic hydrogen at 3000 K is approximately equal to that of liquid rubidium and caesium at 2000 K at the respective pressures when they undergo a nonmetal-to-metal transition.
The 1s1 electron configuration of hydrogen, while analogous to that of the alkali metals (ns1), is unique because there is no 1p subshell. Hence it can lose an electron to form the hydron H+, or gain one to form the hydride ion H−. In the former case it resembles superficially the alkali metals; in the latter case, the halogens, but the differences due to the lack of a 1p subshell are important enough that neither group fits the properties of hydrogen well. Group 14 is also a good fit in terms of thermodynamic properties such as ionisation energy and electron affinity, but hydrogen cannot be tetravalent. Thus none of the three placements are entirely satisfactory, although group 1 is the most common placement (if one is chosen) because of the electron configuration and the fact that the hydron is by far the most important of all monatomic hydrogen species, being the foundation of acid-base chemistry. As an example of hydrogen's unorthodox properties stemming from its unusual electron configuration and small size, the hydrogen ion is very small (radius around 150 fm compared to the 50–220 pm size of most other atoms and ions) and so is nonexistent in condensed systems other than in association with other atoms or molecules. Indeed, transferring of protons between chemicals is the basis of acid-base chemistry. Also unique is hydrogen's ability to form hydrogen bonds, which are an effect of charge-transfer, electrostatic, and electron correlative contributing phenomena. While analogous lithium bonds are also known, they are mostly electrostatic. Nevertheless, hydrogen can take on the same structural role as the alkali metals in some molecular crystals, and has a close relationship with the lightest alkali metals (especially lithium). |
Alkali metal | Ammonium and derivatives | Ammonium and derivatives
thumb|right|Similarly to the alkali metals, ammonia reacts with hydrochloric acid to form the salt ammonium chloride.
The ammonium ion () has very similar properties to the heavier alkali metals, acting as an alkali metal intermediate between potassium and rubidium, and is often considered a close relative. For example, most alkali metal salts are soluble in water, a property which ammonium salts share. Ammonium is expected to behave stably as a metal ( ions in a sea of delocalised electrons) at very high pressures (though less than the typical pressure where transitions from insulating to metallic behaviour occur around, 100 GPa), and could possibly occur inside the ice giants Uranus and Neptune, which may have significant impacts on their interior magnetic fields. It has been estimated that the transition from a mixture of ammonia and dihydrogen molecules to metallic ammonium may occur at pressures just below 25 GPa. Under standard conditions, ammonium can form a metallic amalgam with mercury.
Other "pseudo-alkali metals" include the alkylammonium cations, in which some of the hydrogen atoms in the ammonium cation are replaced by alkyl or aryl groups. In particular, the quaternary ammonium cations () are very useful since they are permanently charged, and they are often used as an alternative to the expensive Cs+ to stabilise very large and very easily polarisable anions such as . Tetraalkylammonium hydroxides, like alkali metal hydroxides, are very strong bases that react with atmospheric carbon dioxide to form carbonates. Furthermore, the nitrogen atom may be replaced by a phosphorus, arsenic, or antimony atom (the heavier nonmetallic pnictogens), creating a phosphonium () or arsonium () cation that can itself be substituted similarly; while stibonium () itself is not known, some of its organic derivatives are characterised. |
Alkali metal | Cobaltocene and derivatives | Cobaltocene and derivatives
Cobaltocene, Co(C5H5)2, is a metallocene, the cobalt analogue of ferrocene. It is a dark purple solid. Cobaltocene has 19 valence electrons, one more than usually found in organotransition metal complexes, such as its very stable relative, ferrocene, in accordance with the 18-electron rule. This additional electron occupies an orbital that is antibonding with respect to the Co–C bonds. Consequently, many chemical reactions of Co(C5H5)2 are characterized by its tendency to lose this "extra" electron, yielding a very stable 18-electron cation known as cobaltocenium. Many cobaltocenium salts coprecipitate with caesium salts, and cobaltocenium hydroxide is a strong base that absorbs atmospheric carbon dioxide to form cobaltocenium carbonate. Like the alkali metals, cobaltocene is a strong reducing agent, and decamethylcobaltocene is stronger still due to the combined inductive effect of the ten methyl groups. Cobalt may be substituted by its heavier congener rhodium to give rhodocene, an even stronger reducing agent. Iridocene (involving iridium) would presumably be still more potent, but is not very well-studied due to its instability. |
Alkali metal | Thallium | Thallium
thumb|right|Very pure thallium pieces in a glass ampoule, stored under argon gas
Thallium is the heaviest stable element in group 13 of the periodic table. At the bottom of the periodic table, the inert-pair effect is quite strong, because of the relativistic stabilisation of the 6s orbital and the decreasing bond energy as the atoms increase in size so that the amount of energy released in forming two more bonds is not worth the high ionisation energies of the 6s electrons. It displays the +1 oxidation state that all the known alkali metals display, and thallium compounds with thallium in its +1 oxidation state closely resemble the corresponding potassium or silver compounds stoichiometrically due to the similar ionic radii of the Tl+ (164 pm), K+ (152 pm) and Ag+ (129 pm) ions. It was sometimes considered an alkali metal in continental Europe (but not in England) in the years immediately following its discovery, and was placed just after caesium as the sixth alkali metal in Dmitri Mendeleev's 1869 periodic table and Julius Lothar Meyer's 1868 periodic table. Mendeleev's 1871 periodic table and Meyer's 1870 periodic table put thallium in its current position in the boron group and left the space below caesium blank. However, thallium also displays the oxidation state +3, which no known alkali metal displays (although ununennium, the undiscovered seventh alkali metal, is predicted to possibly display the +3 oxidation state). The sixth alkali metal is now considered to be francium.. While Tl+ is stabilised by the inert-pair effect, this inert pair of 6s electrons is still able to participate chemically, so that these electrons are stereochemically active in aqueous solution. Additionally, the thallium halides (except TlF) are quite insoluble in water, and TlI has an unusual structure because of the presence of the stereochemically active inert pair in thallium. |
Alkali metal | Copper, silver, and gold | Copper, silver, and gold
The group 11 metals (or coinage metals), copper, silver, and gold, are typically categorised as transition metals given they can form ions with incomplete d-shells. Physically, they have the relatively low melting points and high electronegativity values associated with post-transition metals. "The filled d subshell and free s electron of Cu, Ag, and Au contribute to their high electrical and thermal conductivity. Transition metals to the left of group 11 experience interactions between s electrons and the partially filled d subshell that lower electron mobility."Russell AM & Lee KL (2005) Structure-property relations in nonferrous metals. Wiley-Interscience, New York. p. 302. Chemically, the group 11 metals behave like main-group metals in their +1 valence states, and are hence somewhat related to the alkali metals: this is one reason for their previously being labelled as "group IB", paralleling the alkali metals' "group IA". They are occasionally classified as post-transition metals.Deming HG (1940) Fundamental Chemistry, John Wiley & Sons, New York, pp. 705–7 Their spectra are analogous to those of the alkali metals. Their monopositive ions are paramagnetic and contribute no colour to their salts, like those of the alkali metals.Bailar, J. C. (1973) Comprehensive inorganic chemistry, vol. 3, p. 16.
In Mendeleev's 1871 periodic table, copper, silver, and gold are listed twice, once under group VIII (with the iron triad and platinum group metals), and once under group IB. Group IB was nonetheless parenthesised to note that it was tentative. Mendeleev's main criterion for group assignment was the maximum oxidation state of an element: on that basis, the group 11 elements could not be classified in group IB, due to the existence of copper(II) and gold(III) compounds being known at that time. However, eliminating group IB would make group I the only main group (group VIII was labelled a transition group) to lack an A–B bifurcation. Soon afterward, a majority of chemists chose to classify these elements in group IB and remove them from group VIII for the resulting symmetry: this was the predominant classification until the rise of the modern medium-long 18-column periodic table, which separated the alkali metals and group 11 metals.
The coinage metals were traditionally regarded as a subdivision of the alkali metal group, due to them sharing the characteristic s1 electron configuration of the alkali metals (group 1: p6s1; group 11: d10s1). However, the similarities are largely confined to the stoichiometries of the +1 compounds of both groups, and not their chemical properties. This stems from the filled d subshell providing a much weaker shielding effect on the outermost s electron than the filled p subshell, so that the coinage metals have much higher first ionisation energies and smaller ionic radii than do the corresponding alkali metals. Furthermore, they have higher melting points, hardnesses, and densities, and lower reactivities and solubilities in liquid ammonia, as well as having more covalent character in their compounds. Finally, the alkali metals are at the top of the electrochemical series, whereas the coinage metals are almost at the very bottom. The coinage metals' filled d shell is much more easily disrupted than the alkali metals' filled p shell, so that the second and third ionisation energies are lower, enabling higher oxidation states than +1 and a richer coordination chemistry, thus giving the group 11 metals clear transition metal character. Particularly noteworthy is gold forming ionic compounds with rubidium and caesium, in which it forms the auride ion (Au−) which also occurs in solvated form in liquid ammonia solution: here gold behaves as a pseudohalogen because its 5d106s1 configuration has one electron less than the quasi-closed shell 5d106s2 configuration of mercury. |
Alkali metal | Production and isolation | Production and isolation
The production of pure alkali metals is somewhat complicated due to their extreme reactivity with commonly used substances, such as water. From their silicate ores, all the stable alkali metals may be obtained the same way: sulfuric acid is first used to dissolve the desired alkali metal ion and aluminium(III) ions from the ore (leaching), whereupon basic precipitation removes aluminium ions from the mixture by precipitating it as the hydroxide. The remaining insoluble alkali metal carbonate is then precipitated selectively; the salt is then dissolved in hydrochloric acid to produce the chloride. The result is then left to evaporate and the alkali metal can then be isolated. Lithium and sodium are typically isolated through electrolysis from their liquid chlorides, with calcium chloride typically added to lower the melting point of the mixture. The heavier alkali metals, however, are more typically isolated in a different way, where a reducing agent (typically sodium for potassium and magnesium or calcium for the heaviest alkali metals) is used to reduce the alkali metal chloride. The liquid or gaseous product (the alkali metal) then undergoes fractional distillation for purification. Most routes to the pure alkali metals require the use of electrolysis due to their high reactivity; one of the few which does not is the pyrolysis of the corresponding alkali metal azide, which yields the metal for sodium, potassium, rubidium, and caesium and the nitride for lithium.
Lithium salts have to be extracted from the water of mineral springs, brine pools, and brine deposits. The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride.
Sodium occurs mostly in seawater and dried seabed, but is now produced through electrolysis of sodium chloride by lowering the melting point of the substance to below 700 °C through the use of a Downs cell. Extremely pure sodium can be produced through the thermal decomposition of sodium azide.Merck Index, 9th ed., monograph 8325 Potassium occurs in many minerals, such as sylvite (potassium chloride). Previously, potassium was generally made from the electrolysis of potassium chloride or potassium hydroxide, found extensively in places such as Canada, Russia, Belarus, Germany, Israel, United States, and Jordan, in a method similar to how sodium was produced in the late 1800s and early 1900s. It can also be produced from seawater. However, these methods are problematic because the potassium metal tends to dissolve in its molten chloride and vaporises significantly at the operating temperatures, potentially forming the explosive superoxide. As a result, pure potassium metal is now produced by reducing molten potassium chloride with sodium metal at 850 °C.
Na (g) + KCl (l) NaCl (l) + K (g)
Although sodium is less reactive than potassium, this process works because at such high temperatures potassium is more volatile than sodium and can easily be distilled off, so that the equilibrium shifts towards the right to produce more potassium gas and proceeds almost to completion.
Metals like sodium are obtained by electrolysis of molten salts. Rb & Cs obtained mainly as by products of Li processing. To make pure caesium, ores of caesium and rubidium are crushed and heated to 650 °C with sodium metal, generating an alloy that can then be separated via a fractional distillation technique. Because metallic caesium is too reactive to handle, it is normally offered as caesium azide (CsN3). Caesium hydroxide is formed when caesium interacts aggressively with water and ice (CsOH).
Rubidium is the 16th most abundant element in the earth's crust; however, it is quite rare. Some minerals found in North America, South Africa, Russia, and Canada contain rubidium. Some potassium minerals (lepidolites, biotites, feldspar, carnallite) contain it, together with caesium. Pollucite, carnallite, leucite, and lepidolite are all minerals that contain rubidium. As a by-product of lithium extraction, it is commercially obtained from lepidolite. Rubidium is also found in potassium rocks and brines, which is a commercial supply. The majority of rubidium is now obtained as a byproduct of refining lithium. Rubidium is used in vacuum tubes as a getter, a material that combines with and removes trace gases from vacuum tubes.Liu, Jinlian & Yin, Zhoulan & Li, Xinhai & Hu, Qiyang & Liu, Wei. (2019). A novel process for the selective precipitation of valuable metals from lepidolite. Minerals Engineering. 135. 29–36. 10.1016/j.mineng.2018.11.046.thumb|This sample of uraninite contains about 100,000 atoms (3.3 g) of francium-223 at any given time.|alt=A shiny gray 5-centimeter piece of matter with a rough surface.
For several years in the 1950s and 1960s, a by-product of the potassium production called Alkarb was a main source for rubidium. Alkarb contained 21% rubidium while the rest was potassium and a small fraction of caesium. Today the largest producers of caesium, for example the Tanco Mine in Manitoba, Canada, produce rubidium as by-product from pollucite. Today, a common method for separating rubidium from potassium and caesium is the fractional crystallisation of a rubidium and caesium alum (Cs, Rb)Al(SO4)2·12H2O, which yields pure rubidium alum after approximately 30 recrystallisations. The limited applications and the lack of a mineral rich in rubidium limit the production of rubidium compounds to 2 to 4 tonnes per year. Caesium, however, is not produced from the above reaction. Instead, the mining of pollucite ore is the main method of obtaining pure caesium, extracted from the ore mainly by three methods: acid digestion, alkaline decomposition, and direct reduction. Both metals are produced as by-products of lithium production: after 1958, when interest in lithium's thermonuclear properties increased sharply, the production of rubidium and caesium also increased correspondingly. Pure rubidium and caesium metals are produced by reducing their chlorides with calcium metal at 750 °C and low pressure.
As a result of its extreme rarity in nature, most francium is synthesised in the nuclear reaction 197Au + 18O → 210Fr + 5 n, yielding francium-209, francium-210, and francium-211. The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms, which were synthesised using the nuclear reaction given above. When the only natural isotope francium-223 is specifically required, it is produced as the alpha daughter of actinium-227, itself produced synthetically from the neutron irradiation of natural radium-226, one of the daughters of natural uranium-238. |
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