States of Matter
Gas, Solids and liquids are considered as the three states of matter. The chemical constitution of a substance remains the same in all three states. All these states are interconvertible. State change is also known as a phase change. Every phase change is accompanied by free energy change ∆G.
∆G = ∆H - T∆S
∆H is the enthalpy change associated with making/breaking of bonds and ∆S is the entropy change. Entropy is the degree of disorderliness. As compared to liquids, gases are more disordered and have more entropy.
By the process of Sublimation, solid can directly change into a gas. By condensation, the gaseous state changes into a liquid state. By solidification, liquid changes into a solid-state.
Characteristics of The Gaseous State of Matter
Gases have no definite shape. They fill the containers.
Gases are highly compressible.
Gases have low density.
Large empty spaces between the particles of gases are present.
Kinetic- Molecular Theory of Gases
This theory is based on certain assumptions. This theory deals with the microscopic model of a gas. The properties studied above is a macroscopic model.
Gases consist of tiny particles moving randomly.
The volume of gas molecules is negligible compared to the total volume of the container.
The gas particles act independently of one another. The attractive forces are absent between the particles.
Collisions of a gas particle with other particles or with the wall are elastic collisions. This leads to the total kinetic energy to be constant.
The average kinetic energy of the particles of the gas is directly proportional to the temperature of the gas.
The pressure exerted by a gas is a result of collisions of gas particles with walls of the container.
The speed of particles in gas at a particular time is different. Thus, the kinetic energy will also be different. On collision, the speed changes but the distribution of speed always remains constant. This distribution is Maxwell-Boltzmann distribution.
This is true for fixed amounts of gas.
According to this law, pressure increases when the volume decreases.
P is inversely proportional to V
PV is constant
=> P1V1 = P2V2
Gay - Lussac's Law
The Ideal Gas Equation
All the variables can be related according to the gas laws in a single equation known as the Ideal Gas Equation.
PV = nRT
This is the ideal gas equation and R is the ideal gas constant.
R is expressed in the unit of work or energy per mole per Kelvin.
Ideal Gas Real Gas Difference
An ideal gas is that gas that follows all the gas laws.
For an ideal gas,
PV/RT= Z = 1
Z is the compressibility factor. The graph of PV/RT against p at 0°C is a horizontal straight line.
However, at high and low temperatures, gases are not found to behave ideally. There are various deviations seen from the ideal state.
The Ideal Gas and Real Gas
Real gases behave differently from ideal gas behaviour because of the assumptions taken in the kinetic theory of the gases.
The volume of molecules of a gas is not negligible. This volume correction is required.
It was assumed that there were no intermolecular forces between the gas molecules. Thus a pressure correction is also needed.
At low temperatures, the above effects were negligible but at high temperatures, these assumptions do not prove right.
[Image will be Uploaded Soon]
The figure shows the deviation from ideal gas behaviour.
Difference Between Ideal and Non- Ideal gas
The deviation of the compressibility factor from 1 is the measure of the non-ideality of a gas. For an ideal gas, Z is 1 but for real gas, Z is greater than or less than 1 but not equal to 1.
Van Der Waals Equation represents the behaviour of real gas-
[ p + n2a/V2 ] [ V - nb ] = nRT
a and b are called van der Waals constants. a is the correction of intermolecular attraction and b is the volume correction.