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States of Matter - An Overview of the Gaseous State of Matter

Updated on 15 December 2021
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Updated on 15 December 2021

Definition of Matter


The matter is any substance that has mass and occupies space. It has various physical and chemical properties. Solid, liquid and gas comprise the 3 states of matter. These three types of matter have distinct physical properties as a result of the presence of forces of interactions between the atoms and molecules of which the matter is made up.



Intermolecular forces


Intermolecular forces are the attractive and repulsive forces present between atoms and molecules of the different states of matter. Dipole-dipole, dipole-induced dipole and dispersion forces are all part of attractive intermolecular forces called the van der Waal forces. However, van der Waals forces do not comprise the ion-dipole and ion-induced dipole interactions. Hydrogen bonding is a very strong type of intermolecular force and is one of the dipole-dipole forces.

 

Types of intermolecular forces


  1. Dispersion forces or London forces
  2. Dipole-dipole force
  3. Dipole-induced dipole forces
  4. Hydrogen bond

 

Thermal Energy


There is a production of thermal energy in a substance as a result of the motion of its atoms and molecules. Thus, thermal energy is also dependent upon the average kinetic energy arising due to the movement of atoms and molecules present in it. Changes in temperature mark the changes in thermal energy of the substance.


Thermal energy versus Intermolecular forces


The three states of matter are only possible because of the balance between the intermolecular forces (that are responsible for keeping the atoms and molecules bound together), and thermal energy (that is responsible for separating the molecules due to the motion of the atoms and molecules).


Gas as a state of matter


  1. Gases have no distinct shape and take on the shape of the container.
  2. Because of the weakest intermolecular forces, they lack definite volume.
  3. Gases are easily compressible due to the presence of excess space between their particles, which undergo compression when pressure is applied.
  4. Gases are not rigid.
  5. Gases can easily diffuse because the molecules in a gas move at a very fast rate, resulting in a very high rate of diffusion.
  6. Gases can flow in any direction.


The gas laws


Boyle’s law (Pressure- volume relationship):

Provided the temperature remains constant, the pressure of n 

number of moles of a gas is inversely proportional to the volume of the gas.


    V ∝ 1/ p  (T and n are constant)


Charle’s law (temperature volume relationship)

Provided the pressure remains constant, the volume of a particular mass of a gas varies directly with its absolute temperature.


    V ∝ T ( p and n are constant)


Gay lussac’s law (Pressure temperature relationship)

Provided the volume remains constant, the pressure of a particular amount of gas is directly proportional to its temperature.


     P∝T (V and N remain constant)


Avogadro’s law (relationship between volume and amount of gas)


According to this law, provided the pressure and temperature remain the same, the volume of the gas varies directly with the number of moles of that gas.


    V ∝ n at constant p and T


Ideal Gas Equation


An ideal gas equation is a type of equation formed by the combination of three gas laws, which are T


Boyle’s Law V ∝ 1/ p at constant T and n,


Charles' Law; V ∝ T at constant p and n


Avogadro Law ; V ∝ n at constant p and T


Thus, V n ∝ T p

 

V n = R T p where R is proportionality constant.


On rearranging the equations we get;


pV = nRT


R = pV Tn

 

R is called a gas constant. It applies to all the gases. and is thus also known as Universal Gas Constant.


Density and Molar Mass of a Gaseous Substance

According to the ideal gas equation,

pV = nRT

Or n/V = p/RT

 n = Mass of Gas (m) / Molar Mass of Gas (M) = m/M

So, m/MV = p/RT

d/M= p/RT

 

Dalton’s Law of Partial Pressures


According to Dalton's law of partial pressure, the total pressure exerted by the mixture of non-reactive gases is equivalent to the sum of the partial pressures of individual gases. Mathematically,


pTotal = p1+p2+p3+......(at constant T, V).


Where p1, p2, p3……….. = Partial pressures of gases 1,2, and 3.


Partial Pressure in terms of Mole Fraction


Let the three gases be at

T = temperature three gases,

V = volume,

p1, p2 and p3 = partial pressure exerted on gases 1, 2 and 3 respectively.

ptotal = p1+p2+p3

ptotal = n1RT/V + n2RT/V + n3RT/V

ptotal = (n1+n2+n3)RT/V

Dividing p1 by pTotal = n1/ (n1+n2+n3)

P1 = x1 ptotal

Similarly, p2= x2 ptotal

         P3 = x3 ptotal


Kinetic Molecular Theory of Gases


  1. Molecules are volumeless point masses.
  2. Unless they collide, gas atoms impose no constraints on other atoms.
  3. Collisions of particles with each other or with the container's boundaries do not lower the energy of the system.
  4. A gas's atoms move in both regular and irregular patterns.
  5. A gas's temperature is determined by its average kinetic energy. 3/2 KT = 1/2 mv2
  6. Ideal gases do contain kinetic energy.

 

Gases and their Liquefaction


By combining the effects of compression and temperature, gases can be easily liquefied. Compression brings the particles closer together, whereas cooling slows the motion of the molecules. Intermolecular interactions may hold the rapidly and gradually moving particles together, causing the gas to liquefy. Underneath the normal temperature of the liquid, the fluid and gas phase can be easily recognized. At normal temperature, the fluid gradually and continuously transitions into a vaporous state, and the surface isolating the two stages fades away. Vapour pressure refers to the ability of gas below its basic normal temperature to be condensed by applying weight.


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