Revision: Thermodynamics Chemistry HSC Science Class 11 Tamil Nadu Board of Secondary Education

A part of the universe chosen for analysis, separated by boundaries, where energy and matter exchanges are observed, is called a thermodynamic system.

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An assembly of an extremely large number of particles (atoms or molecules) in solid, liquid, gaseous, or a combination of two or more phases is called a thermodynamic system.

Everything outside the system which has direct effect on the system is called its surrounding.

A system that exchanges neither energy nor matter with its surroundings is called an isolated system.

A system that exchanges only energy (not matter) with its surroundings is called a closed system.

The branch of science which deals with exchange of heat energy between bodies and conversion of heat energy into mechanical energy and vice versa is called thermodynamics.

A system that exchanges both energy and matter with its surroundings is called an open system.

A thermodynamic process during which there is no transfer of heat (energy) from the system to the surroundings or from the surroundings to the system is called an adiabatic process.

It is defined as one in which the volume of the system remains constant during its change from the initial to the final stage of the process.

It is defined as one in which the pressure of the system remains constant during its change from the initial to the final state.

When two bodies at different temperatures are brought into contact through a diathermic wall, heat flows from the hotter body to the cooler one. This continues until both reach the same temperature, at which point heat flow stops. This state is called thermal equilibrium.

The temperature at which pure water freezes at 1 atm pressure is called the ice point.

The temperature at which pure water boils and vapourises into steam at 1 atm pressure is called the steam point or boiling point.

Thermometry is the branch of physics dealing with temperature measurement. It relies on the principle that certain physical properties of materials change continuously and predictably with temperature.

An adiabatic wall is an ideal partition that completely prevents heat transfer between two systems. In diagrams, it is shown as a thick, cross-hatched (slanting lines) region.

A diathermic wall is a partition that freely allows heat to flow between two systems. It is shown as a thin dark line in diagrams. A thin copper sheet is a good example.

The thermodynamic state variables that do not depend on the size of the system (e.g., pressure, temperature) are called intensive variables.

The specific values of macroscopic variables that completely describe every equilibrium state of a thermodynamic system are called thermodynamic state variables.

The thermodynamic state variables that depend on the size of the system (e.g., internal energy, volume) are called extensive variables.

It is the enthalpy change during the hydration of 1 mole of anhydrous salt by the addition of a specific number of moles of water.

The total heat content of a system at constant pressure is known as enthalpу.

At constant pressure: ΔH = q_p​ (heat exchanged at constant pressure).

It is defined as the change in heat enthalpy when one mole of a substance is completely burnt in oxygen.

ΔΗ = Σ (Heat of Combustion of reactant)- Σ (Heat of Combustion of product)

It is the enthalpy change when one mole of it dissolves in a specified amount of solvent

It is the enthalpy change associated with diluting a component in a solution at constant pressure and temperature.

Enthalpy of a system is sum of internal energy of a system and the energy equivalent to PV work.

H = U + PV

The heat of combustion of a substance is defined as “The change in enthalpy of a system when one mole of the substance is completely burnt in excess of air or oxygen”. It is denoted by ∆H_C.

It is defined as the heat evolved or decrease in enthalpy when 1 gm equivalent of an acid is neutralised by 1 gm equivalent of a base in solution.

The enthalpy of neutralization is defined as the change in enthalpy of the system when one gram equivalent of an acid is neutralized by one gram equivalent of a base or vice versa in dilute solution.

\[\ce{H^+_{(aq)} + OH^-_{(aq)} -> H2O_{(l)}}\] = 57.32 kJ

The heat capacity for 1 mole of a substance, is called molar heat capacity (cm). It is defined as “The amount of heat absorbed by one mole of the substance to raise its temperature by 1 kelvin”.

A useful state variable that measures the change in heat divided by the temperature of the system, where the combined entropy of the system and its environment remains constant if the process approaches reversibility, is called entropy.

Gibbs free energy is defined as the part of the total energy of a system that can be converted (or) available for conversion into work.

G = H − TS

where G = Gibb’s free energy

H = enthalpy

T = temperature

S = entropy

Master Conversion Formula:

\[\frac{T_F-32}{180}=\frac{T_C}{100}\] = \[\frac {T_K−273.15}{100}\]

By substituting equation W = −p_ex . ΔV in the equation ΔU = q + W, we get

ΔU = q − p_ex . ΔV  ...(1)

If the reaction is carried out in a closed container so that the volume of the system is constant, then Δ = 0. In such a case, no work is involved.

The equation (1) becomes ΔU = q_v

Equation (1) suggests that the change in internal energy of the system is due to heat transfer. The subscript v indicates a constant volume process. As U is a state function, qv is also a state function. We see that an increase in the internal energy of a system is numerically equal to the heat absorbed by the system in a constant volume (isochoric) process.

ΔS = \[\frac {ΔQ}{T}\]

If system A is in thermal equilibrium with system C, and system B is also in thermal equilibrium with system C, then systems A and B are in thermal equilibrium with each other.

By substituting equation W = −p_ex . ΔV in the equation ΔU = q + W, we get

ΔU = q − p_ex . ΔV  ...(1)

If the reaction is carried out in a closed container so that the volume of the system is constant, then Δ = 0. In such a case, no work is involved.

The equation (1) becomes ΔU = q_v

Equation (1) suggests that the change in internal energy of the system is due to heat transfer. The subscript v indicates a constant volume process. As U is a state function, qv is also a state function. We see that an increase in the internal energy of a system is numerically equal to the heat absorbed by the system in a constant volume (isochoric) process.

Statement:
The net heat energy supplied to a system is equal to the sum of the change in internal energy of the system and the work done by the system. It is based on the law of conservation of energy.

Formula:

where Q = heat added, ΔU = change in internal energy, W = work done by the system.

The second law of thermodynamics states that the total entropy of a system and its surroundings increases in a spontaneous process.

Mathematically,

ΔS_total = `Delta S_"system" + Delta S_"surroundings" gt 0`

For an equilibrium:

ΔS_total = 0

The third law of thermodynamics states that the entropy of a pure crystalline substance at absolute zero is zero. Otherwise, it can be stated that it is impossible to lower the temperature of an object to absolute zero in a finite number of steps. Mathematically,

`lim_(T->0)` S = 0 for a perfectly ordered crystalline state.

First Law: Energy of system + surroundings remains constant → ΔU = q + W

ΔU: change in internal energy, q: heat, W: work done on system

Sign convention:

• Work by system (−)
• on system (+)
• Heat absorbed (+)
• released (−)

ΔU > 0: energy enters system; ΔU < 0: energy leaves system

• Isothermal: ΔU = 0 → q = −W
• Adiabatic: q = 0 → ΔU = W
• Isochoric: W = 0 → ΔU = q
• Isobaric: ΔU = q + W

Entropy (S): A thermodynamic property that measures the degree of randomness or disorder of a system.

\[\Delta S=\frac{q_{rev}}{T}\]

Second Law: The entropy of the universe always tends to increase during any spontaneous process.

Total entropy change:

Entropy of mixing:

(ΔS for mixing is always positive, since ΔS is always fractional/positive)

Third Law: The entropy of a perfectly ordered crystalline substance at absolute zero (0 K) is zero.

Absolute Entropy (Standard Molar Entropy, S°): The entropy value of a substance at 298 K and 1 bar pressure, relative to 0 K.

• Unlike the enthalpy of formation, the absolute entropy of elements is not zero at standard conditions.
• Standard entropy of reaction: \[\Delta_{r}S^{\circ}=\sum S_{products}^{\circ}-\sum S_{reactants}^{\circ}\]