Why do some gases liquefy easily under pressure and normal room temperature?
The gas laws you describe, Boyle's and Charles' Laws, can be combined into one relationship known as the Ideal Gas Law. According to this law,

$$PV=nRT$$

Where 

$P$ is the pressure of the gas ($\ce{Pa}$), 

$V$ is the volume of the gas ($\ce{m^3}$),

$n$ is the number of moles or amount of gas ($\ce{mol}$),

$R$ is a constant ($R\sim 8.31 \ce{\frac{J}{mol.K}}$), and

$T$ is the temperature of the gas ($\ce{K}$).

To see how changing one property affects another property of the sample of gas, hold the other values constant and see what must happen.

For example, according to the law, an increase in Pressure will result in a decrease in Volume, if Temperature is held constant. 
But if the Volume is instead held constant, then the Temperature will increase instead.

In effect, the ideal gas law is the combination of the other laws affecting a gas and its properties.

But this does not explain why some substances are liquids at room temperature while others are not and require cooling and/or a change of pressure.

To describe this, an understanding of what are known as Intermolecular forces must be expressed.

Between molecules of a given substance, there exist attractions. These attractions vary in strength and are what determine the pressures and temperatures needed to liquify specific chemicals.

The (usually) weakest kind of attractions are known as London Dispersion Forces. These forces are generated randomly due to the random movement of electrons about a chemical. Since the movement is random, there come times where the electron density is more negative on one side of a molecule than another. This results in a very slight charge, which can induce charges onto neighboring molecules and make them attract each other.

The next class of attractions are known as Dipole-Dipole Interactions. This is a result of a permanent charge density difference across a molecule. In these molecules, one or more atoms are much greater in their affinities for electrons and will draw the density away from the atoms with lesser affinity. This makes for a permeant slight charge on one or more sides of a molecule that are positive, while the other sides become negative. The negative ends of one molecule will attract the positive ends of neighboring molecules, resulting in a usually pretty strong attractive force.

The strongest by far class of interactions are known as Hydrogen Bonding. This is when hydrogen is bonded to en extremely electronegative atom, like oxygen or fluorine. Take water for example. The Oxygen atom greatly skews the electron density, making the oxygen end much more negative in charge than the positive hydrogen ends. This makes for very strong intermolecular attractions.

With these forces understood, you can now see that chemicals vary in the strength of their intermolecular attractions. 

All that said, the weaker the intermolecular attractions, the easier it is for the molecules to break their attractions and become gaseous. Such chemicals with weak attractions require so little energy to become gaseous that at room temperature and standard pressure, the molecules act in the gas phase. To make these chemicals turn into liquids, one must either increase the pressure to make the formation of attractions more favorable or lower the temperature to lessen the energy of the molecules so that they don't break their attractive forces.

Some chemicals have extremely weak attractive forces, meaning that the temperature will have to be very low or the pressure very high to force the chemical into its liquid state. 

Other chemicals, like water, exhibit such strong forces that the temperature has to be raised or the pressure lowered to vaporize them.

So yes, the gas properties of pressure, temperature, and volume do affect changes of state, but its the strength of the intermolecular attractions that are the most important and are the reasons that chemicals behave so differently under the same conditions.