Gases – General Characteristics of gases

States of the matter

– All matter exists in three states: gases, liquids and solids.

– A gas consists of molecules separated wide apart in empty space. The molecules are free to move about throughout the container.

– A liquid has molecules touching each other. However, the intermolecular space, permits the movement of molecules throughout the liquid.

– A solid has molecules, atoms or ions arranged in a certain order in fixed positions in the crystal lattice.

– The particles in a solid are not free to move about but vibrate in their fixed positions. 

– Of the three states of matter, the gaseous state is the one most studied and best understood. We shall consider it first.

General Characteristics of gases

(1) Expansibility

– Gases have limitless expansibility.

– They expand to fill the entire vessel they are placed in.

(2) Compressibility

– Gases are easily compressed by application of pressure to a movable piston fitted in the container

(3) Diffusibility

– Gases can diffuse rapidly through each other to form a homogeneous mixture.

(4) Pressure

– Gases exert pressure on the walls of the container in all directions.

(5) Effect of Heat

– When a gas, confined in a vessel is heated, its pressure increases.

– Upon heating in a vessel fitted with a piston, the volume of the gas increases.

– The above properties of gases can be easily explained by the Kinetic Molecular Theory.

Parameters of gases

– A gas sample can be described in terms of four parameters (measurable properties):

(1) the volume, V of the gas

(2) its pressure, P

(3) its temperature, T

(4) the number of moles, n, of gas in the container

(1) Volume of the gas, V

– The volume of the container is the volume of the gas sample.

– It is usually given in litre (l or L) or millilitres (ml or mL).

1 1itre(l) = 1000 ml and 1 ml = 10^–3 

– The pressure of a gas is defined as the force exerted by the impacts of its molecules per unit surface area in contact.

– The pressure of a gas sample can be measured with the help of a mercury manometer (Fig. 1) Similarly, the atmospheric pressure can be determined with a mercury barometer (Fig. 2).

 

– The pressure of air that can support a 760 mm Hg column at sea level, is called one atmosphere (1 atm).

– The unit of pressure, a millimetre of mercury, is also called torr. Thus,

1 atm = 760 mm Hg = 760 torr 

– The SI unit of pressure is the Pascal (Pa). The relation between the atmosphere, torr and Pascal is :

1 atm = 760 torr = 1.013 × 105 Pa 

– The unit of pressure (Pascal) is not in common use.

(3) Temperature, T

– The temperature of a gas may be measured in Centigrade degrees (°C) or Celsius degrees.

– The SI unit of temperature is Kelvin (K) or Absolute degree.

– The centigrade degrees can be converted to kelvins by using the equation.

– The Kelvin temperature (or absolute temperature) is always used in calculations of other parameters of gases.

– Remember that the degree sign (°) is not used with K.

(4) The Moles of a Gas Sample, n

The number of moles, n, of a sample of a gas in a container can be found by dividing the mass, m, of the sample by the molar mass, M (molecular mass).

The gas Laws

– The volume of a given sample of gas depends on the temperature and pressure applied to it.

– Any change in temperature or pressure will affect the volume of the gas.

– As a result of experimental studies from the 17th to 19th century, scientists derived the relationships among the pressure, temperature and volume of a given mass of gas.

– These relationships, which describe the general behaviour of gases, are called gas laws.

– The Ideal Gas Equation:

 – This is called the Universal Gas Law. It is also called the Ideal Gas Law

– Ideal Gas Law as it applies to all gases which exhibit ideal behaviour i.e., obey the gas laws perfectly.

– The ideal gas law may be stated as the volume of a given amount of gas is directly proportional to the number of moles of gas, directly proportional to the temperature, and inversely proportional to the pressure.