Chemistry – handbook – Chemical Thermodynamics

Chemical Thermodynamics

Basic concepts
System, surroundings and Boundary: A specified part of the universe which is under observation is called the system and the remaining portion of the universe which is not a part of the system is called the surroundings.

Types of systems
(i) Isolated system: This type of system has no interaction with its surroundings. The boundary is sealed and insulated. Neither matter nor energy can be exchanged with surrounding. A substance contained in an ideal thermos flask is an example of an isolated system.
(ii) Closed system: This type of system can exchange energy in the form of heat, work or radiations but not matter with its surroundings. The boundary between system and surroundings is sealed but not insulated. For example, liquid in contact with vapor in a sealed tube and pressure cooker.
(iii) Open system: This type of system can exchange matter as well as energy with its surroundings. The boundary is neither sealed nor insulated. Sodium reacting with water in an open beaker is an example of open system. The macroscopic properties can be subdivided into two types,

(i) Intensive properties: The properties which do not depend upon the quantity of matter present in the system or size of the system are called intensive properties. Its examples are pressure, temperature, density, specific heat, surface tension, refractive index, viscosity, melting point, boiling point, volume per mole, concentration etc.

(ii) Mechanical Equilibrium: No chemical work is done between the different parts of the system or between the system and the surrounding. To achieve it keep the pressure constant.

(iii) Thermal Equilibrium : It can be achieved when the temperature remains constant i.e. no flow of heat between the system and the surrounding.

(ii) Extensive properties: The properties whose magnitude depends upon the quantity of matter present in the system are called extensive properties. Its examples are total mass, volume, internal energy, enthalpy, entropy etc. These properties are additive in nature. Any extensive property if expressed as per mole or per gram becomes an intensive property.

State of a system and State Variables
Macroscopic properties that specify the state of a system are referred to as state variables or state functions or parameters of thermodynamics. The change in the state properties is depends only on the system’s initial and final state. It is independent on the path by which the change has been brought out.

Thermodynamic equilibrium: “A system is said to have attained a state of thermodynamic equilibrium when it does not show any further tendency to change its property with time”. The thermodynamic equilibrium criterion specifies that the following three equilibrium forms occur simultaneously in a system;
(i) Chemical Equilibrium: A system where system composition stays constant and definit e.

(ii) Mechanical Equilibrium: No chemical work is done between the different parts of the system or between the system and the surrounding. To achieve it keep the pressure constant.
(iii) Thermal Equilibrium : It can be achieved when the temperature remains constant i.e. no flow of heat between the system and the surrounding.

Thermodynamic process: When the thermodynamic system changes from one state to another, the operation is called a process. The various types of the processes are
(i) Isothermal process: In this process operation is done at constant temperature. dT = 0 thus ΔE = 0 .
(ii) Adiabatic process: In this a process there is no exchange of heat takes place between the system and surroundings. The system is thermally isolated, i.e., dQ = 0 and its boundaries are insulated.
(iii) Isobaric process: In this process the pressure remains constant throughout the change i.e., dP = 0.
(iv) Isochoric process: In this process volume remains constant throughout the change, i.e., dV = 0.
(v) Cyclic process: When a system undergoes a number of different processes and finally return to its initial state, it is termed cyclic process. For a cyclic process dE = 0 and dH = 0.

(vi) Reversible process: A process which occurs
infinitesimally slowly, i.e. opposing force is infinitesimally smaller than driving force and when infinitesimal increase in the opposing force can reverse the process, it is said to be reversible process.
(vii) Irreversible process: When the process occurs from initial to final state in single step in finite time and cannot be reversed, it is termed an irreversible process. Amount of entropy increases in irreversible process. Irreversible processes are spontaneous in nature. All natural processes are irreversible in nature

Internal energy, heat and Work

Internal energy (E): “Every system having some quantity of matter is associated with a definite amount of energy so It is the energy contained within the system. This energy is sum of all types of energies (chemical, mechanical electrical etc.) and is a state function.
First law of thermodynamics
This law is also known as law of conservation of energy. It is proposed by Helmholtz and Robert Mayer. It states that, “Energy can neither be created nor destroyed although it can be converted from one form into another.”

i.e. Change in internal energy = Heat added to the system +Work done on the system

E2 – E1 =ΔE= q +w

If a system does work (w) on the surroundings, its internal energy decreases. In this case, Change in internal energy = Heat added to the system – work done by the system

ΔE = q +(-w) = q – w

Enthalpy and Enthalpy change
The heat content of a system at constant pressure is called enthalpy. It is denoted by ‘H’. From the first law of thermodynamics, q =E+ W ;

q= E+ PV  …………( i)
Heat change in constant pressure can be given,

Δ q=Δ E+ PΔ V ……..( ii)
At constant pressure heat can be replaced at enthalpy.
ΔH=Δ E +P ΔV  ……..( iii)
∴ΔH Is the heat of reaction (in chemical process) or the heat
change at constant pressure.

Specific and Molar heat capacity
Specific heat capacity of a substance is the quantity of heat that is required to raise the temperature of 1g of substance through

The specific heat capacity of a substance is the amount of heat required to raise the temperature of 1g of the substance C .
It can be calculated at constant pressure ( cpand constant volume ( cv )
q= mc ΔT 
The molar heat capacity of a substance is the amount of heat required to raise the temperature of 1 mole of the substance 1° C .

∴ Molar heat capacity = Specific heat capacity  Molecular weight, i.e., Cm = c × M
Since,

Gases on heat show an inclination toward expansion when heated under constant pressure conditions, with an additional energy supplied to raise its temperature 1° C . when required under constant volume conditions.
Thus,
CP  > CV .
CP  = CV + .Work done in expansion, 
Work done in expansion is PΔ V( =R)

Here, Cv and Cp   are molar heat capacities at constant volume and constant pressure respectively.

Some useful relations of Cp and Cv

i) Cp – Cv =R =2 calories =8.314 J

ii)Cv =(3/2) R (for monoatomic gas) and Cv  = (3/2 ) + x(for di and polyatomic gas), where x varies from gas to gas.

iii)cp /cv = γ (Ratio of molar capacities)

iv)For monoatomic gas,  Cv =3 calories whereas,  Cp  = C+ R = 5 calories

v)For monoatomic gas, (γ) = Cp/Cv = (5/2)R / (3/2 ) R = 1.66

(vi) For diatomic gas  (γ) = Cp/Cv = (7/2)R / (5/2 ) R =1.40

(vii) For triatomic gas (γ) =Cp/Cv =  8R /6R =1.33

Second law of thermodynamics
The second law of thermodynamics removes all limitations of the first law of thermodynamics. This law states that for spontaneous processes the total energy change is positive..

The Carnot cycle
The basic condition for a cyclic process is that the net work
done is equal to the heat absorbed. This condition is satisfied in a Carnot cycle.
w/q2  = T2 -T1 / T2   Thermodynamic efficiency
Thus,

The greater the difference in temperature between high and lo w temperature reservoirs, the more heat, and the heat engine converts into work.

Entropy and Entropy change
Definition : Entropy is a measure of randomness or disorder of the system’s molecules. It is a thermodynamic state quantity

S final – S initial =ΔS =q rev /T

If heat is absorbed, then  ΔS =+ve and if heat is evolved, then ΔS =-ve 

Free energy and Free energy change
Gibb’s free energy (G) is a measure of maximum work done which is also known as useful work done from a reversible reaction at constant pressure and temperature.
We can find it by subtracting the product of absolute temperature and entropy from the enthalpy of the system.

i.e., G = H -TS

Third law of thermodynamics
This law was first formulated in 1906 by German chemist Walther Nernst

This law states that, “The entropy of all perfectly crystalline solids approaches zero at the temperature approaches to absolute zero.
Sinceentropy is a measure of disorder it can be interpreted that a perfectly crystalline solid has a perfect order of its constitue nt particles at absolute zero.

Exothermic and Endothermic reactions
Exothermic reactions: Exothermic reactions are those reactions that proceed with the evolution of heat energy.

ΔH = Hp – H r ; Hp <Hr ; ΔH = -ve

Examples : (i) Fermentation is an example of exothermic reaction.

(ii) Formation of CO2 gas is also an example of exothermic reaction.

C(s) +O2(g)  → CO2 (g) + 393.5kJ

It can also be written as; 

C(s) +O2(g)  → CO2 (g) ;    ΔH =-393.5kJ(at constant temperature and pressure)

Endothermic reactions : Endothermic reactions are those reactions that proceed with the absorption of heat energy.

ΔH = Hp – Hr ; Hp > Hr ;  ΔH = +ve

Example : formation of NO is an example of endothermic reaction.

N2(g) + O2(g) → 2NO (g) ;Δ H =+180.5kJ

Heat of reaction or Enthalpy of reaction
The heat of reaction or the reaction enthalpy is the difference between the products and the reactants enthalpies, that we get when the reactants reacted completelly in the che mical reaction.
Enthalpy of reaction (heat of reaction)  = Δ H = ∑Hp – ∑H R

 

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