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1 | 1905-1908 | 3) This implies
that the pressure of hydrogen gas is one bar
and the concentration of hydrogen ion in the
solution is one molar 2 2 |
1 | 1906-1909 | This implies
that the pressure of hydrogen gas is one bar
and the concentration of hydrogen ion in the
solution is one molar 2 2 1
Measurement
of Electrode
Potential
Fig |
1 | 1907-1910 | 2 2 1
Measurement
of Electrode
Potential
Fig 2 |
1 | 1908-1911 | 2 1
Measurement
of Electrode
Potential
Fig 2 3: Standard Hydrogen Electrode (SHE) |
1 | 1909-1912 | 1
Measurement
of Electrode
Potential
Fig 2 3: Standard Hydrogen Electrode (SHE) Rationalised 2023-24
35
Electrochemistry
At 298 K the emf of the cell, standard hydrogen electrode ççsecond
half-cell constructed by taking standard hydrogen electrode as anode
(reference half-cell) and the other half-cell as cathode, gives the reduction
potential of the other half-cell |
1 | 1910-1913 | 2 3: Standard Hydrogen Electrode (SHE) Rationalised 2023-24
35
Electrochemistry
At 298 K the emf of the cell, standard hydrogen electrode ççsecond
half-cell constructed by taking standard hydrogen electrode as anode
(reference half-cell) and the other half-cell as cathode, gives the reduction
potential of the other half-cell If the concentrations of the oxidised and
the reduced forms of the species in the right hand half-cell are unity,
then the cell potential is equal to standard electrode potential, Eo
R of
the given half-cell |
1 | 1911-1914 | 3: Standard Hydrogen Electrode (SHE) Rationalised 2023-24
35
Electrochemistry
At 298 K the emf of the cell, standard hydrogen electrode ççsecond
half-cell constructed by taking standard hydrogen electrode as anode
(reference half-cell) and the other half-cell as cathode, gives the reduction
potential of the other half-cell If the concentrations of the oxidised and
the reduced forms of the species in the right hand half-cell are unity,
then the cell potential is equal to standard electrode potential, Eo
R of
the given half-cell Eo = Eo
R – Eo
L
As Eo
L for standard hydrogen electrode is zero |
1 | 1912-1915 | Rationalised 2023-24
35
Electrochemistry
At 298 K the emf of the cell, standard hydrogen electrode ççsecond
half-cell constructed by taking standard hydrogen electrode as anode
(reference half-cell) and the other half-cell as cathode, gives the reduction
potential of the other half-cell If the concentrations of the oxidised and
the reduced forms of the species in the right hand half-cell are unity,
then the cell potential is equal to standard electrode potential, Eo
R of
the given half-cell Eo = Eo
R – Eo
L
As Eo
L for standard hydrogen electrode is zero Eo = Eo
R – 0 = Eo
R
The measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Cu
2+ (aq, 1 M) ú Cu
is 0 |
1 | 1913-1916 | If the concentrations of the oxidised and
the reduced forms of the species in the right hand half-cell are unity,
then the cell potential is equal to standard electrode potential, Eo
R of
the given half-cell Eo = Eo
R – Eo
L
As Eo
L for standard hydrogen electrode is zero Eo = Eo
R – 0 = Eo
R
The measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Cu
2+ (aq, 1 M) ú Cu
is 0 34 V and it is also the value for the standard electrode potential
of the half-cell corresponding to the reaction:
Cu
2+ (aq, 1M) + 2 e
– ® Cu(s)
Similarly, the measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Zn
2+ (aq, 1M) ç Zn
is -0 |
1 | 1914-1917 | Eo = Eo
R – Eo
L
As Eo
L for standard hydrogen electrode is zero Eo = Eo
R – 0 = Eo
R
The measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Cu
2+ (aq, 1 M) ú Cu
is 0 34 V and it is also the value for the standard electrode potential
of the half-cell corresponding to the reaction:
Cu
2+ (aq, 1M) + 2 e
– ® Cu(s)
Similarly, the measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Zn
2+ (aq, 1M) ç Zn
is -0 76 V corresponding to the standard electrode potential of the
half-cell reaction:
Zn
2+ (aq, 1 M) + 2e
– ® Zn(s)
The positive value of the standard electrode potential in the first
case indicates that Cu
2+ ions get reduced more easily than H
+ ions |
1 | 1915-1918 | Eo = Eo
R – 0 = Eo
R
The measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Cu
2+ (aq, 1 M) ú Cu
is 0 34 V and it is also the value for the standard electrode potential
of the half-cell corresponding to the reaction:
Cu
2+ (aq, 1M) + 2 e
– ® Cu(s)
Similarly, the measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Zn
2+ (aq, 1M) ç Zn
is -0 76 V corresponding to the standard electrode potential of the
half-cell reaction:
Zn
2+ (aq, 1 M) + 2e
– ® Zn(s)
The positive value of the standard electrode potential in the first
case indicates that Cu
2+ ions get reduced more easily than H
+ ions The
reverse process cannot occur, that is, hydrogen ions cannot oxidise Cu
(or alternatively we can say that hydrogen gas can reduce copper ion)
under the standard conditions described above |
1 | 1916-1919 | 34 V and it is also the value for the standard electrode potential
of the half-cell corresponding to the reaction:
Cu
2+ (aq, 1M) + 2 e
– ® Cu(s)
Similarly, the measured emf of the cell:
Pt(s) ç H2(g, 1 bar) ç H
+ (aq, 1 M) çç Zn
2+ (aq, 1M) ç Zn
is -0 76 V corresponding to the standard electrode potential of the
half-cell reaction:
Zn
2+ (aq, 1 M) + 2e
– ® Zn(s)
The positive value of the standard electrode potential in the first
case indicates that Cu
2+ ions get reduced more easily than H
+ ions The
reverse process cannot occur, that is, hydrogen ions cannot oxidise Cu
(or alternatively we can say that hydrogen gas can reduce copper ion)
under the standard conditions described above Thus, Cu does not
dissolve in HCl |
1 | 1917-1920 | 76 V corresponding to the standard electrode potential of the
half-cell reaction:
Zn
2+ (aq, 1 M) + 2e
– ® Zn(s)
The positive value of the standard electrode potential in the first
case indicates that Cu
2+ ions get reduced more easily than H
+ ions The
reverse process cannot occur, that is, hydrogen ions cannot oxidise Cu
(or alternatively we can say that hydrogen gas can reduce copper ion)
under the standard conditions described above Thus, Cu does not
dissolve in HCl In nitric acid it is oxidised by nitrate ion and not by
hydrogen ion |
1 | 1918-1921 | The
reverse process cannot occur, that is, hydrogen ions cannot oxidise Cu
(or alternatively we can say that hydrogen gas can reduce copper ion)
under the standard conditions described above Thus, Cu does not
dissolve in HCl In nitric acid it is oxidised by nitrate ion and not by
hydrogen ion The negative value of the standard electrode potential
in the second case indicates that hydrogen ions can oxidise zinc (or
zinc can reduce hydrogen ions) |
1 | 1919-1922 | Thus, Cu does not
dissolve in HCl In nitric acid it is oxidised by nitrate ion and not by
hydrogen ion The negative value of the standard electrode potential
in the second case indicates that hydrogen ions can oxidise zinc (or
zinc can reduce hydrogen ions) In view of this convention, the half reaction for the Daniell cell in
Fig |
1 | 1920-1923 | In nitric acid it is oxidised by nitrate ion and not by
hydrogen ion The negative value of the standard electrode potential
in the second case indicates that hydrogen ions can oxidise zinc (or
zinc can reduce hydrogen ions) In view of this convention, the half reaction for the Daniell cell in
Fig 2 |
1 | 1921-1924 | The negative value of the standard electrode potential
in the second case indicates that hydrogen ions can oxidise zinc (or
zinc can reduce hydrogen ions) In view of this convention, the half reaction for the Daniell cell in
Fig 2 1 can be written as:
Left electrode: Zn(s) ® Zn
2+ (aq, 1 M) + 2 e
–
Right electrode: Cu
2+ (aq, 1 M) + 2 e
– ® Cu(s)
The overall reaction of the cell is the sum of above two reactions
and we obtain the equation:
Zn(s) + Cu
2+ (aq) ® Zn
2+ (aq) + Cu(s)
emf of the cell = Eo
cell = Eo
R – Eo
L
= 0 |
1 | 1922-1925 | In view of this convention, the half reaction for the Daniell cell in
Fig 2 1 can be written as:
Left electrode: Zn(s) ® Zn
2+ (aq, 1 M) + 2 e
–
Right electrode: Cu
2+ (aq, 1 M) + 2 e
– ® Cu(s)
The overall reaction of the cell is the sum of above two reactions
and we obtain the equation:
Zn(s) + Cu
2+ (aq) ® Zn
2+ (aq) + Cu(s)
emf of the cell = Eo
cell = Eo
R – Eo
L
= 0 34V – (– 0 |
1 | 1923-1926 | 2 1 can be written as:
Left electrode: Zn(s) ® Zn
2+ (aq, 1 M) + 2 e
–
Right electrode: Cu
2+ (aq, 1 M) + 2 e
– ® Cu(s)
The overall reaction of the cell is the sum of above two reactions
and we obtain the equation:
Zn(s) + Cu
2+ (aq) ® Zn
2+ (aq) + Cu(s)
emf of the cell = Eo
cell = Eo
R – Eo
L
= 0 34V – (– 0 76)V = 1 |
1 | 1924-1927 | 1 can be written as:
Left electrode: Zn(s) ® Zn
2+ (aq, 1 M) + 2 e
–
Right electrode: Cu
2+ (aq, 1 M) + 2 e
– ® Cu(s)
The overall reaction of the cell is the sum of above two reactions
and we obtain the equation:
Zn(s) + Cu
2+ (aq) ® Zn
2+ (aq) + Cu(s)
emf of the cell = Eo
cell = Eo
R – Eo
L
= 0 34V – (– 0 76)V = 1 10 V
Sometimes metals like platinum or gold are used as inert electrodes |
1 | 1925-1928 | 34V – (– 0 76)V = 1 10 V
Sometimes metals like platinum or gold are used as inert electrodes They do not participate in the reaction but provide their surface for
oxidation or reduction reactions and for the conduction of electrons |
1 | 1926-1929 | 76)V = 1 10 V
Sometimes metals like platinum or gold are used as inert electrodes They do not participate in the reaction but provide their surface for
oxidation or reduction reactions and for the conduction of electrons For example, Pt is used in the following half-cells:
Hydrogen electrode: Pt(s)|H2(g)| H+(aq)
With half-cell reaction: H+ (aq)+ e– ® ½ H2(g)
Bromine electrode: Pt(s)|Br2(aq)| Br–(aq)
Rationalised 2023-24
36
Chemistry
With half-cell reaction: ½ Br2(aq) + e– ® Br–(aq)
The standard electrode potentials are very important and we can
extract a lot of useful information from them |
1 | 1927-1930 | 10 V
Sometimes metals like platinum or gold are used as inert electrodes They do not participate in the reaction but provide their surface for
oxidation or reduction reactions and for the conduction of electrons For example, Pt is used in the following half-cells:
Hydrogen electrode: Pt(s)|H2(g)| H+(aq)
With half-cell reaction: H+ (aq)+ e– ® ½ H2(g)
Bromine electrode: Pt(s)|Br2(aq)| Br–(aq)
Rationalised 2023-24
36
Chemistry
With half-cell reaction: ½ Br2(aq) + e– ® Br–(aq)
The standard electrode potentials are very important and we can
extract a lot of useful information from them The values of standard
electrode potentials for some selected half-cell reduction reactions are
given in Table 2 |
1 | 1928-1931 | They do not participate in the reaction but provide their surface for
oxidation or reduction reactions and for the conduction of electrons For example, Pt is used in the following half-cells:
Hydrogen electrode: Pt(s)|H2(g)| H+(aq)
With half-cell reaction: H+ (aq)+ e– ® ½ H2(g)
Bromine electrode: Pt(s)|Br2(aq)| Br–(aq)
Rationalised 2023-24
36
Chemistry
With half-cell reaction: ½ Br2(aq) + e– ® Br–(aq)
The standard electrode potentials are very important and we can
extract a lot of useful information from them The values of standard
electrode potentials for some selected half-cell reduction reactions are
given in Table 2 1 |
1 | 1929-1932 | For example, Pt is used in the following half-cells:
Hydrogen electrode: Pt(s)|H2(g)| H+(aq)
With half-cell reaction: H+ (aq)+ e– ® ½ H2(g)
Bromine electrode: Pt(s)|Br2(aq)| Br–(aq)
Rationalised 2023-24
36
Chemistry
With half-cell reaction: ½ Br2(aq) + e– ® Br–(aq)
The standard electrode potentials are very important and we can
extract a lot of useful information from them The values of standard
electrode potentials for some selected half-cell reduction reactions are
given in Table 2 1 If the standard electrode potential of an electrode
is greater than zero then its reduced form is more stable compared to
hydrogen gas |
1 | 1930-1933 | The values of standard
electrode potentials for some selected half-cell reduction reactions are
given in Table 2 1 If the standard electrode potential of an electrode
is greater than zero then its reduced form is more stable compared to
hydrogen gas Similarly, if the standard electrode potential is negative
then hydrogen gas is more stable than the reduced form of the species |
1 | 1931-1934 | 1 If the standard electrode potential of an electrode
is greater than zero then its reduced form is more stable compared to
hydrogen gas Similarly, if the standard electrode potential is negative
then hydrogen gas is more stable than the reduced form of the species It can be seen that the standard electrode potential for fluorine is the
highest in the Table indicating that fluorine gas (F2) has the maximum
tendency to get reduced to fluoride ions (F–) and therefore fluorine
gas is the strongest oxidising agent and fluoride ion is the weakest
reducing agent |
1 | 1932-1935 | If the standard electrode potential of an electrode
is greater than zero then its reduced form is more stable compared to
hydrogen gas Similarly, if the standard electrode potential is negative
then hydrogen gas is more stable than the reduced form of the species It can be seen that the standard electrode potential for fluorine is the
highest in the Table indicating that fluorine gas (F2) has the maximum
tendency to get reduced to fluoride ions (F–) and therefore fluorine
gas is the strongest oxidising agent and fluoride ion is the weakest
reducing agent Lithium has the lowest electrode potential indicating
that lithium ion is the weakest oxidising agent while lithium metal is
the most powerful reducing agent in an aqueous solution |
1 | 1933-1936 | Similarly, if the standard electrode potential is negative
then hydrogen gas is more stable than the reduced form of the species It can be seen that the standard electrode potential for fluorine is the
highest in the Table indicating that fluorine gas (F2) has the maximum
tendency to get reduced to fluoride ions (F–) and therefore fluorine
gas is the strongest oxidising agent and fluoride ion is the weakest
reducing agent Lithium has the lowest electrode potential indicating
that lithium ion is the weakest oxidising agent while lithium metal is
the most powerful reducing agent in an aqueous solution It may be
seen that as we go from top to bottom in Table 2 |
1 | 1934-1937 | It can be seen that the standard electrode potential for fluorine is the
highest in the Table indicating that fluorine gas (F2) has the maximum
tendency to get reduced to fluoride ions (F–) and therefore fluorine
gas is the strongest oxidising agent and fluoride ion is the weakest
reducing agent Lithium has the lowest electrode potential indicating
that lithium ion is the weakest oxidising agent while lithium metal is
the most powerful reducing agent in an aqueous solution It may be
seen that as we go from top to bottom in Table 2 1 the standard
electrode potential decreases and with this, decreases the oxidising
power of the species on the left and increases the reducing power of
the species on the right hand side of the reaction |
1 | 1935-1938 | Lithium has the lowest electrode potential indicating
that lithium ion is the weakest oxidising agent while lithium metal is
the most powerful reducing agent in an aqueous solution It may be
seen that as we go from top to bottom in Table 2 1 the standard
electrode potential decreases and with this, decreases the oxidising
power of the species on the left and increases the reducing power of
the species on the right hand side of the reaction Electrochemical
cells are extensively used for determining the pH of solutions, solubility
product, equilibrium constant and other thermodynamic properties
and for potentiometric titrations |
1 | 1936-1939 | It may be
seen that as we go from top to bottom in Table 2 1 the standard
electrode potential decreases and with this, decreases the oxidising
power of the species on the left and increases the reducing power of
the species on the right hand side of the reaction Electrochemical
cells are extensively used for determining the pH of solutions, solubility
product, equilibrium constant and other thermodynamic properties
and for potentiometric titrations Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
2 |
1 | 1937-1940 | 1 the standard
electrode potential decreases and with this, decreases the oxidising
power of the species on the left and increases the reducing power of
the species on the right hand side of the reaction Electrochemical
cells are extensively used for determining the pH of solutions, solubility
product, equilibrium constant and other thermodynamic properties
and for potentiometric titrations Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
2 1 How would you determine the standard electrode potential of the system
Mg2+|Mg |
1 | 1938-1941 | Electrochemical
cells are extensively used for determining the pH of solutions, solubility
product, equilibrium constant and other thermodynamic properties
and for potentiometric titrations Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
2 1 How would you determine the standard electrode potential of the system
Mg2+|Mg 2 |
1 | 1939-1942 | Intext Questions
Intext Questions
Intext Questions
Intext Questions
Intext Questions
2 1 How would you determine the standard electrode potential of the system
Mg2+|Mg 2 2 Can you store copper sulphate solutions in a zinc pot |
1 | 1940-1943 | 1 How would you determine the standard electrode potential of the system
Mg2+|Mg 2 2 Can you store copper sulphate solutions in a zinc pot 2 |
1 | 1941-1944 | 2 2 Can you store copper sulphate solutions in a zinc pot 2 3 Consult the table of standard electrode potentials and suggest three
substances that can oxidise ferrous ions under suitable conditions |
1 | 1942-1945 | 2 Can you store copper sulphate solutions in a zinc pot 2 3 Consult the table of standard electrode potentials and suggest three
substances that can oxidise ferrous ions under suitable conditions 2 |
1 | 1943-1946 | 2 3 Consult the table of standard electrode potentials and suggest three
substances that can oxidise ferrous ions under suitable conditions 2 3
2 |
1 | 1944-1947 | 3 Consult the table of standard electrode potentials and suggest three
substances that can oxidise ferrous ions under suitable conditions 2 3
2 3
2 |
1 | 1945-1948 | 2 3
2 3
2 3
2 |
1 | 1946-1949 | 3
2 3
2 3
2 3
2 |
1 | 1947-1950 | 3
2 3
2 3
2 3 Nernst
Nernst
Nernst
Nernst
Nernst
Equation
Equation
Equation
Equation
Equation
We have assumed in the previous section that the concentration of all
the species involved in the electrode reaction is unity |
1 | 1948-1951 | 3
2 3
2 3 Nernst
Nernst
Nernst
Nernst
Nernst
Equation
Equation
Equation
Equation
Equation
We have assumed in the previous section that the concentration of all
the species involved in the electrode reaction is unity This need not be
always true |
1 | 1949-1952 | 3
2 3 Nernst
Nernst
Nernst
Nernst
Nernst
Equation
Equation
Equation
Equation
Equation
We have assumed in the previous section that the concentration of all
the species involved in the electrode reaction is unity This need not be
always true Nernst showed that for the electrode reaction:
Mn+(aq) + ne–® M(s)
the electrode potential at any concentration measured with respect to
standard hydrogen electrode can be represented by:
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/ M
E
E
– RT
nF ln [M]
[M
]
n+
but concentration of solid M is taken as unity and we have
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/M
E
E
–
RT
nF ln
n+
1
[M
]
(2 |
1 | 1950-1953 | 3 Nernst
Nernst
Nernst
Nernst
Nernst
Equation
Equation
Equation
Equation
Equation
We have assumed in the previous section that the concentration of all
the species involved in the electrode reaction is unity This need not be
always true Nernst showed that for the electrode reaction:
Mn+(aq) + ne–® M(s)
the electrode potential at any concentration measured with respect to
standard hydrogen electrode can be represented by:
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/ M
E
E
– RT
nF ln [M]
[M
]
n+
but concentration of solid M is taken as unity and we have
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/M
E
E
–
RT
nF ln
n+
1
[M
]
(2 8)
(
)
n+
o
M
/ M
E
has already been defined, R is gas constant (8 |
1 | 1951-1954 | This need not be
always true Nernst showed that for the electrode reaction:
Mn+(aq) + ne–® M(s)
the electrode potential at any concentration measured with respect to
standard hydrogen electrode can be represented by:
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/ M
E
E
– RT
nF ln [M]
[M
]
n+
but concentration of solid M is taken as unity and we have
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/M
E
E
–
RT
nF ln
n+
1
[M
]
(2 8)
(
)
n+
o
M
/ M
E
has already been defined, R is gas constant (8 314 JK–1 mol–1),
F is Faraday constant (96487 C mol–1), T is temperature in kelvin and
[Mn+] is the concentration of the species, Mn+ |
1 | 1952-1955 | Nernst showed that for the electrode reaction:
Mn+(aq) + ne–® M(s)
the electrode potential at any concentration measured with respect to
standard hydrogen electrode can be represented by:
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/ M
E
E
– RT
nF ln [M]
[M
]
n+
but concentration of solid M is taken as unity and we have
(
)
(
)
+
+
=
n
n
o
M
/ M
M
/M
E
E
–
RT
nF ln
n+
1
[M
]
(2 8)
(
)
n+
o
M
/ M
E
has already been defined, R is gas constant (8 314 JK–1 mol–1),
F is Faraday constant (96487 C mol–1), T is temperature in kelvin and
[Mn+] is the concentration of the species, Mn+ Rationalised 2023-24
37
Electrochemistry
F2(g) + 2e–
® 2F–
2 |
1 | 1953-1956 | 8)
(
)
n+
o
M
/ M
E
has already been defined, R is gas constant (8 314 JK–1 mol–1),
F is Faraday constant (96487 C mol–1), T is temperature in kelvin and
[Mn+] is the concentration of the species, Mn+ Rationalised 2023-24
37
Electrochemistry
F2(g) + 2e–
® 2F–
2 87
Co3+ + e–
® Co2+
1 |
1 | 1954-1957 | 314 JK–1 mol–1),
F is Faraday constant (96487 C mol–1), T is temperature in kelvin and
[Mn+] is the concentration of the species, Mn+ Rationalised 2023-24
37
Electrochemistry
F2(g) + 2e–
® 2F–
2 87
Co3+ + e–
® Co2+
1 81
H2O2 + 2H+ + 2e–
® 2H2O
1 |
1 | 1955-1958 | Rationalised 2023-24
37
Electrochemistry
F2(g) + 2e–
® 2F–
2 87
Co3+ + e–
® Co2+
1 81
H2O2 + 2H+ + 2e–
® 2H2O
1 78
MnO4
– + 8H+ + 5e–
® Mn2+ + 4H2O
1 |
1 | 1956-1959 | 87
Co3+ + e–
® Co2+
1 81
H2O2 + 2H+ + 2e–
® 2H2O
1 78
MnO4
– + 8H+ + 5e–
® Mn2+ + 4H2O
1 51
Au3+ + 3e–
® Au(s)
1 |
1 | 1957-1960 | 81
H2O2 + 2H+ + 2e–
® 2H2O
1 78
MnO4
– + 8H+ + 5e–
® Mn2+ + 4H2O
1 51
Au3+ + 3e–
® Au(s)
1 40
Cl2(g) + 2e–
® 2Cl–
1 |
1 | 1958-1961 | 78
MnO4
– + 8H+ + 5e–
® Mn2+ + 4H2O
1 51
Au3+ + 3e–
® Au(s)
1 40
Cl2(g) + 2e–
® 2Cl–
1 36
Cr2O7
2– + 14H+ + 6e–
® 2Cr3+ + 7H2O
1 |
1 | 1959-1962 | 51
Au3+ + 3e–
® Au(s)
1 40
Cl2(g) + 2e–
® 2Cl–
1 36
Cr2O7
2– + 14H+ + 6e–
® 2Cr3+ + 7H2O
1 33
O2(g) + 4H+ + 4e–
® 2H2O
1 |
1 | 1960-1963 | 40
Cl2(g) + 2e–
® 2Cl–
1 36
Cr2O7
2– + 14H+ + 6e–
® 2Cr3+ + 7H2O
1 33
O2(g) + 4H+ + 4e–
® 2H2O
1 23
MnO2(s) + 4H+ + 2e–
® Mn2+ + 2H2O
1 |
1 | 1961-1964 | 36
Cr2O7
2– + 14H+ + 6e–
® 2Cr3+ + 7H2O
1 33
O2(g) + 4H+ + 4e–
® 2H2O
1 23
MnO2(s) + 4H+ + 2e–
® Mn2+ + 2H2O
1 23
Br2 + 2e–
® 2Br–
1 |
1 | 1962-1965 | 33
O2(g) + 4H+ + 4e–
® 2H2O
1 23
MnO2(s) + 4H+ + 2e–
® Mn2+ + 2H2O
1 23
Br2 + 2e–
® 2Br–
1 09
NO3
– + 4H+ + 3e–
® NO(g) + 2H2O
0 |
1 | 1963-1966 | 23
MnO2(s) + 4H+ + 2e–
® Mn2+ + 2H2O
1 23
Br2 + 2e–
® 2Br–
1 09
NO3
– + 4H+ + 3e–
® NO(g) + 2H2O
0 97
2Hg2+ + 2e–
® Hg2
2+
0 |
1 | 1964-1967 | 23
Br2 + 2e–
® 2Br–
1 09
NO3
– + 4H+ + 3e–
® NO(g) + 2H2O
0 97
2Hg2+ + 2e–
® Hg2
2+
0 92
Ag+ + e–
® Ag(s)
0 |
1 | 1965-1968 | 09
NO3
– + 4H+ + 3e–
® NO(g) + 2H2O
0 97
2Hg2+ + 2e–
® Hg2
2+
0 92
Ag+ + e–
® Ag(s)
0 80
Fe3+ + e–
® Fe2+
0 |
1 | 1966-1969 | 97
2Hg2+ + 2e–
® Hg2
2+
0 92
Ag+ + e–
® Ag(s)
0 80
Fe3+ + e–
® Fe2+
0 77
O2(g) + 2H+ + 2e–
® H2O2
0 |
1 | 1967-1970 | 92
Ag+ + e–
® Ag(s)
0 80
Fe3+ + e–
® Fe2+
0 77
O2(g) + 2H+ + 2e–
® H2O2
0 68
I2 + 2e–
® 2I–
0 |
1 | 1968-1971 | 80
Fe3+ + e–
® Fe2+
0 77
O2(g) + 2H+ + 2e–
® H2O2
0 68
I2 + 2e–
® 2I–
0 54
Cu+ + e–
® Cu(s)
0 |
1 | 1969-1972 | 77
O2(g) + 2H+ + 2e–
® H2O2
0 68
I2 + 2e–
® 2I–
0 54
Cu+ + e–
® Cu(s)
0 52
Cu2+ + 2e–
® Cu(s)
0 |
1 | 1970-1973 | 68
I2 + 2e–
® 2I–
0 54
Cu+ + e–
® Cu(s)
0 52
Cu2+ + 2e–
® Cu(s)
0 34
AgCl(s) + e–
® Ag(s) + Cl–
0 |
1 | 1971-1974 | 54
Cu+ + e–
® Cu(s)
0 52
Cu2+ + 2e–
® Cu(s)
0 34
AgCl(s) + e–
® Ag(s) + Cl–
0 22
AgBr(s) + e–
® Ag(s) + Br–
0 |
1 | 1972-1975 | 52
Cu2+ + 2e–
® Cu(s)
0 34
AgCl(s) + e–
® Ag(s) + Cl–
0 22
AgBr(s) + e–
® Ag(s) + Br–
0 10
2H+ + 2e–
® H2(g)
0 |
1 | 1973-1976 | 34
AgCl(s) + e–
® Ag(s) + Cl–
0 22
AgBr(s) + e–
® Ag(s) + Br–
0 10
2H+ + 2e–
® H2(g)
0 00
Pb2+ + 2e–
® Pb(s)
–0 |
1 | 1974-1977 | 22
AgBr(s) + e–
® Ag(s) + Br–
0 10
2H+ + 2e–
® H2(g)
0 00
Pb2+ + 2e–
® Pb(s)
–0 13
Sn2+ + 2e–
® Sn(s)
–0 |
1 | 1975-1978 | 10
2H+ + 2e–
® H2(g)
0 00
Pb2+ + 2e–
® Pb(s)
–0 13
Sn2+ + 2e–
® Sn(s)
–0 14
Ni2+ + 2e–
® Ni(s)
–0 |
1 | 1976-1979 | 00
Pb2+ + 2e–
® Pb(s)
–0 13
Sn2+ + 2e–
® Sn(s)
–0 14
Ni2+ + 2e–
® Ni(s)
–0 25
Fe2+ + 2e–
® Fe(s)
–0 |
1 | 1977-1980 | 13
Sn2+ + 2e–
® Sn(s)
–0 14
Ni2+ + 2e–
® Ni(s)
–0 25
Fe2+ + 2e–
® Fe(s)
–0 44
Cr3+ + 3e–
® Cr(s)
–0 |
1 | 1978-1981 | 14
Ni2+ + 2e–
® Ni(s)
–0 25
Fe2+ + 2e–
® Fe(s)
–0 44
Cr3+ + 3e–
® Cr(s)
–0 74
Zn2+ + 2e–
® Zn(s)
–0 |
1 | 1979-1982 | 25
Fe2+ + 2e–
® Fe(s)
–0 44
Cr3+ + 3e–
® Cr(s)
–0 74
Zn2+ + 2e–
® Zn(s)
–0 76
2H2O + 2e–
® H2(g) + 2OH–(aq)
–0 |
1 | 1980-1983 | 44
Cr3+ + 3e–
® Cr(s)
–0 74
Zn2+ + 2e–
® Zn(s)
–0 76
2H2O + 2e–
® H2(g) + 2OH–(aq)
–0 83
Al3+ + 3e–
® Al(s)
–1 |
1 | 1981-1984 | 74
Zn2+ + 2e–
® Zn(s)
–0 76
2H2O + 2e–
® H2(g) + 2OH–(aq)
–0 83
Al3+ + 3e–
® Al(s)
–1 66
Mg2+ + 2e–
® Mg(s)
–2 |
1 | 1982-1985 | 76
2H2O + 2e–
® H2(g) + 2OH–(aq)
–0 83
Al3+ + 3e–
® Al(s)
–1 66
Mg2+ + 2e–
® Mg(s)
–2 36
Na+ + e–
® Na(s)
–2 |
1 | 1983-1986 | 83
Al3+ + 3e–
® Al(s)
–1 66
Mg2+ + 2e–
® Mg(s)
–2 36
Na+ + e–
® Na(s)
–2 71
Ca2+ + 2e–
® Ca(s)
–2 |
1 | 1984-1987 | 66
Mg2+ + 2e–
® Mg(s)
–2 36
Na+ + e–
® Na(s)
–2 71
Ca2+ + 2e–
® Ca(s)
–2 87
K+ + e–
® K(s)
–2 |
1 | 1985-1988 | 36
Na+ + e–
® Na(s)
–2 71
Ca2+ + 2e–
® Ca(s)
–2 87
K+ + e–
® K(s)
–2 93
Li+ + e–
® Li(s)
–3 |
1 | 1986-1989 | 71
Ca2+ + 2e–
® Ca(s)
–2 87
K+ + e–
® K(s)
–2 93
Li+ + e–
® Li(s)
–3 05
Table 2 |
1 | 1987-1990 | 87
K+ + e–
® K(s)
–2 93
Li+ + e–
® Li(s)
–3 05
Table 2 1: Standard Electrode Potentials at 298 K
Ions are present as aqueous species and H2O as liquid; gases and solids are shown by g and s |
1 | 1988-1991 | 93
Li+ + e–
® Li(s)
–3 05
Table 2 1: Standard Electrode Potentials at 298 K
Ions are present as aqueous species and H2O as liquid; gases and solids are shown by g and s Reaction (Oxidised form + ne–
®
®
®
®
® Reduced form)
E o/V
Increasing strength of oxidising agent
Increasing strength of reducing agent
1 |
1 | 1989-1992 | 05
Table 2 1: Standard Electrode Potentials at 298 K
Ions are present as aqueous species and H2O as liquid; gases and solids are shown by g and s Reaction (Oxidised form + ne–
®
®
®
®
® Reduced form)
E o/V
Increasing strength of oxidising agent
Increasing strength of reducing agent
1 A negative Eo means that the redox couple is a stronger reducing agent than the H+/H2 couple |
1 | 1990-1993 | 1: Standard Electrode Potentials at 298 K
Ions are present as aqueous species and H2O as liquid; gases and solids are shown by g and s Reaction (Oxidised form + ne–
®
®
®
®
® Reduced form)
E o/V
Increasing strength of oxidising agent
Increasing strength of reducing agent
1 A negative Eo means that the redox couple is a stronger reducing agent than the H+/H2 couple 2 |
1 | 1991-1994 | Reaction (Oxidised form + ne–
®
®
®
®
® Reduced form)
E o/V
Increasing strength of oxidising agent
Increasing strength of reducing agent
1 A negative Eo means that the redox couple is a stronger reducing agent than the H+/H2 couple 2 A positive Eo means that the redox couple is a weaker reducing agent than the H+/H2 couple |
1 | 1992-1995 | A negative Eo means that the redox couple is a stronger reducing agent than the H+/H2 couple 2 A positive Eo means that the redox couple is a weaker reducing agent than the H+/H2 couple Rationalised 2023-24
38
Chemistry
In Daniell cell, the electrode potential for any given concentration of
Cu2+ and Zn2+ ions, we write
For Cathode:
Cu2
/Cu
E
=
(
)
2+
o
Cu
/Cu
E
– RT
F
2 ln
2
1
Cu
aq
(2 |
1 | 1993-1996 | 2 A positive Eo means that the redox couple is a weaker reducing agent than the H+/H2 couple Rationalised 2023-24
38
Chemistry
In Daniell cell, the electrode potential for any given concentration of
Cu2+ and Zn2+ ions, we write
For Cathode:
Cu2
/Cu
E
=
(
)
2+
o
Cu
/Cu
E
– RT
F
2 ln
2
1
Cu
aq
(2 9)
For Anode:
Zn2
/Zn
E
=
(
)
2+
o
Zn
/ Zn
E
– RT
F
2 ln
2
1
Zn
aq
(2 |
1 | 1994-1997 | A positive Eo means that the redox couple is a weaker reducing agent than the H+/H2 couple Rationalised 2023-24
38
Chemistry
In Daniell cell, the electrode potential for any given concentration of
Cu2+ and Zn2+ ions, we write
For Cathode:
Cu2
/Cu
E
=
(
)
2+
o
Cu
/Cu
E
– RT
F
2 ln
2
1
Cu
aq
(2 9)
For Anode:
Zn2
/Zn
E
=
(
)
2+
o
Zn
/ Zn
E
– RT
F
2 ln
2
1
Zn
aq
(2 10)
The cell potential, E(cell) =
Cu2
/Cu
E
–
Zn2
/Zn
E
=
(
)
2+
o
Cu
/ Cu
E
– RT
F
2 ln
2+
1
Cu
(aq)
–
(
)
2+
o
Zn
/ Zn
E
+ RT
2F
ln
2+
1
Zn
(aq)
=
(
)
2+
o
Cu
/ Cu
E
–
(
)
2+
o
Zn
/ Zn
E
– RT
F
2
2+
2+
1
1
ln
– ln
Cu
aq
Zn
aq
E(cell) =
(
)
o
Ecell
– RT
2F
ln [
]
+
[
]
Zn2
2
Cu
(2 |
1 | 1995-1998 | Rationalised 2023-24
38
Chemistry
In Daniell cell, the electrode potential for any given concentration of
Cu2+ and Zn2+ ions, we write
For Cathode:
Cu2
/Cu
E
=
(
)
2+
o
Cu
/Cu
E
– RT
F
2 ln
2
1
Cu
aq
(2 9)
For Anode:
Zn2
/Zn
E
=
(
)
2+
o
Zn
/ Zn
E
– RT
F
2 ln
2
1
Zn
aq
(2 10)
The cell potential, E(cell) =
Cu2
/Cu
E
–
Zn2
/Zn
E
=
(
)
2+
o
Cu
/ Cu
E
– RT
F
2 ln
2+
1
Cu
(aq)
–
(
)
2+
o
Zn
/ Zn
E
+ RT
2F
ln
2+
1
Zn
(aq)
=
(
)
2+
o
Cu
/ Cu
E
–
(
)
2+
o
Zn
/ Zn
E
– RT
F
2
2+
2+
1
1
ln
– ln
Cu
aq
Zn
aq
E(cell) =
(
)
o
Ecell
– RT
2F
ln [
]
+
[
]
Zn2
2
Cu
(2 11)
It can be seen that E(cell) depends on the concentration of both Cu2+
and Zn2+ ions |
1 | 1996-1999 | 9)
For Anode:
Zn2
/Zn
E
=
(
)
2+
o
Zn
/ Zn
E
– RT
F
2 ln
2
1
Zn
aq
(2 10)
The cell potential, E(cell) =
Cu2
/Cu
E
–
Zn2
/Zn
E
=
(
)
2+
o
Cu
/ Cu
E
– RT
F
2 ln
2+
1
Cu
(aq)
–
(
)
2+
o
Zn
/ Zn
E
+ RT
2F
ln
2+
1
Zn
(aq)
=
(
)
2+
o
Cu
/ Cu
E
–
(
)
2+
o
Zn
/ Zn
E
– RT
F
2
2+
2+
1
1
ln
– ln
Cu
aq
Zn
aq
E(cell) =
(
)
o
Ecell
– RT
2F
ln [
]
+
[
]
Zn2
2
Cu
(2 11)
It can be seen that E(cell) depends on the concentration of both Cu2+
and Zn2+ ions It increases with increase in the concentration of Cu2+
ions and decrease in the concentration of Zn2+ ions |
1 | 1997-2000 | 10)
The cell potential, E(cell) =
Cu2
/Cu
E
–
Zn2
/Zn
E
=
(
)
2+
o
Cu
/ Cu
E
– RT
F
2 ln
2+
1
Cu
(aq)
–
(
)
2+
o
Zn
/ Zn
E
+ RT
2F
ln
2+
1
Zn
(aq)
=
(
)
2+
o
Cu
/ Cu
E
–
(
)
2+
o
Zn
/ Zn
E
– RT
F
2
2+
2+
1
1
ln
– ln
Cu
aq
Zn
aq
E(cell) =
(
)
o
Ecell
– RT
2F
ln [
]
+
[
]
Zn2
2
Cu
(2 11)
It can be seen that E(cell) depends on the concentration of both Cu2+
and Zn2+ ions It increases with increase in the concentration of Cu2+
ions and decrease in the concentration of Zn2+ ions By converting the natural logarithm in Eq |
1 | 1998-2001 | 11)
It can be seen that E(cell) depends on the concentration of both Cu2+
and Zn2+ ions It increases with increase in the concentration of Cu2+
ions and decrease in the concentration of Zn2+ ions By converting the natural logarithm in Eq (2 |
1 | 1999-2002 | It increases with increase in the concentration of Cu2+
ions and decrease in the concentration of Zn2+ ions By converting the natural logarithm in Eq (2 11) to the base 10 and
substituting the values of R, F and T = 298 K, it reduces to
E(cell) =
(
)
o
Ecell
–
0 059
2
2
2 |
1 | 2000-2003 | By converting the natural logarithm in Eq (2 11) to the base 10 and
substituting the values of R, F and T = 298 K, it reduces to
E(cell) =
(
)
o
Ecell
–
0 059
2
2
2 [
]
[
]
log Zn
Cu
+
+
(2 |
1 | 2001-2004 | (2 11) to the base 10 and
substituting the values of R, F and T = 298 K, it reduces to
E(cell) =
(
)
o
Ecell
–
0 059
2
2
2 [
]
[
]
log Zn
Cu
+
+
(2 12)
We should use the same number of electrons (n) for both the
electrodes and thus for the following cell
Ni(s)ú Ni2+(aq) úú Ag+(aq)ú Ag
The cell reaction is Ni(s) + 2Ag+(aq) ® Ni2+(aq) + 2Ag(s)
The Nernst equation can be written as
E(cell) =
(
)
o
Ecell
– RT
F
2 ln [Ni
]
[Ag ]
2+
2
+
and for a general electrochemical reaction of the type:
a A + bB
ne–
cC + dD
Nernst equation can be written as:
E(cell) =
(
)
o
Ecell
–
RT
nF 1nQ
=
(
)
o
Ecell
–
RT
nF ln
[C] [D]
[A] [B]
c
d
a
b
(2 |
1 | 2002-2005 | 11) to the base 10 and
substituting the values of R, F and T = 298 K, it reduces to
E(cell) =
(
)
o
Ecell
–
0 059
2
2
2 [
]
[
]
log Zn
Cu
+
+
(2 12)
We should use the same number of electrons (n) for both the
electrodes and thus for the following cell
Ni(s)ú Ni2+(aq) úú Ag+(aq)ú Ag
The cell reaction is Ni(s) + 2Ag+(aq) ® Ni2+(aq) + 2Ag(s)
The Nernst equation can be written as
E(cell) =
(
)
o
Ecell
– RT
F
2 ln [Ni
]
[Ag ]
2+
2
+
and for a general electrochemical reaction of the type:
a A + bB
ne–
cC + dD
Nernst equation can be written as:
E(cell) =
(
)
o
Ecell
–
RT
nF 1nQ
=
(
)
o
Ecell
–
RT
nF ln
[C] [D]
[A] [B]
c
d
a
b
(2 13)
Rationalised 2023-24
39
Electrochemistry
If the circuit in Daniell cell (Fig |
1 | 2003-2006 | [
]
[
]
log Zn
Cu
+
+
(2 12)
We should use the same number of electrons (n) for both the
electrodes and thus for the following cell
Ni(s)ú Ni2+(aq) úú Ag+(aq)ú Ag
The cell reaction is Ni(s) + 2Ag+(aq) ® Ni2+(aq) + 2Ag(s)
The Nernst equation can be written as
E(cell) =
(
)
o
Ecell
– RT
F
2 ln [Ni
]
[Ag ]
2+
2
+
and for a general electrochemical reaction of the type:
a A + bB
ne–
cC + dD
Nernst equation can be written as:
E(cell) =
(
)
o
Ecell
–
RT
nF 1nQ
=
(
)
o
Ecell
–
RT
nF ln
[C] [D]
[A] [B]
c
d
a
b
(2 13)
Rationalised 2023-24
39
Electrochemistry
If the circuit in Daniell cell (Fig 2 |
1 | 2004-2007 | 12)
We should use the same number of electrons (n) for both the
electrodes and thus for the following cell
Ni(s)ú Ni2+(aq) úú Ag+(aq)ú Ag
The cell reaction is Ni(s) + 2Ag+(aq) ® Ni2+(aq) + 2Ag(s)
The Nernst equation can be written as
E(cell) =
(
)
o
Ecell
– RT
F
2 ln [Ni
]
[Ag ]
2+
2
+
and for a general electrochemical reaction of the type:
a A + bB
ne–
cC + dD
Nernst equation can be written as:
E(cell) =
(
)
o
Ecell
–
RT
nF 1nQ
=
(
)
o
Ecell
–
RT
nF ln
[C] [D]
[A] [B]
c
d
a
b
(2 13)
Rationalised 2023-24
39
Electrochemistry
If the circuit in Daniell cell (Fig 2 1) is closed then we note that the reaction
Zn(s) + Cu2+(aq) ® Zn2+(aq) + Cu(s)
(2 |
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