Chemical Thermodynamics Class 11 Notes Chemistry Chapter 5 - CBSE
Chapter : 5
What Are Chemical Thermodynamics ?
The System And The Surroundings
A system in thermodynamics refers to that part of universe which is under observation. Everything else in the universe except system is called surroundings.
The universe = The system + The surroundings
The State Of The System
The state of a thermodynamic system is described by its measurable or macroscopic (bulk) properties. Variables like pressure (p), volume (V), temperature (T), amount (n) are called state variables or state functions because their values depend only on the state of the system and not on how it is reached.
Types Of The System
- Open System: There is exchange of energy and matter between system and surroundings. For example, presence of reactants in an open beaker.
- Closed System: There is no exchange of matter, but exchange of energy is possible between system and the surroundings. For example, presence of reactants in a closed vessel made of conducting material such as copper or steel.
- Isolated System: There is no exchange of energy or matter between the system and the surroundings. For
example, presence of reactants in a thermos flask or any other closed insulated vessel.
- Adiabatic Process: It is a process in which there is no transfer of heat between the system and surroundings.
- Reversible Process: It is a process which proceeds infinitely slowly by a series of equilibrium states such that system and the surroundings are always in near equilibrium with each other.
- Irreversible Process: Processes other than reversible processes are known as irreversible processes.
Free Expansion
Expansion of a gas in vacuum (Pex = 0) is called free expansion. No work is done during free expansion of an ideal gas whether the process is reversible or irreversible.
∆U = q - pex∆V
If a process is carried out at constant volume (∆V = 0), then
∆U = qv
the subscript v in qv denotes that heat is supplied at constant volume.
Isothermal and free expansion of an ideal gas
For isothermal (T = constant) expansion of an ideal gas into vacuum; w = 0 since pex = 0. Also, q = 0; therefore, ∆U = 0. Thus,
- For isothermal irreversible change: q = w = pex(Vf - Vi)
- For isothermal reversible change:
$$\text{q = -w = nRT ln}\frac{\text{V}_{r}}{\text{V}_{l}}\\= 2.303\space\text{nRT}\text{log}\frac{V_{f}}{V_{i}}$$
- For adiabatic change, q = 0.
∆U = wad
Enthalpy (H)
It is defined as total heat content of the system. It is equal to the sum of internal energy and pressure-volume work. Mathematically,
H = U + PV
Change in enthalpy is the heat absorbed or evolved by the system at constant pressure.
∆H = qp
For exothermic reaction (System loses energy to Surroundings),
∆H and qp both are negative.
For endothermic reaction (System absorbs energy from the Surroundings).
∆H and qp both are positive.
Relation between ΔH and ΔU
Let us consider a general reaction
$$\text{A}\xrightarrow{}\text{B}$$
Let HA be the enthalpy of reactant A and HB be that of the products
∴ HA = UA + PVA
HB = UB + PVB
∆H = HB- HA
= (UB + PVB) - (UA + PVA)
∆H = ∆U + P∆V (HB - HA)
∆H = ∆U + P∆V
At constant pressure and temperature using ideal gas law,
PVA = nA RT (For reactant A)
PVB = nBRT (For reactant B)
Thus, PVB- PVA = nBRT - nART
= (nB - nA)RT
P∆V = ∆n g RT
∴ ∆H = ∆U + ∆ngRT
Extensive And Intensive Properties
- Extensive property: It is a property whose value depends on the quantity or size of matter present in the system. Mass, volume, enthalpy etc. are known as extensive property.
- Intensive property: The properties which do not depend upon the size of the matter or quantity of the matter present in the system. Temperature, density, pressure etc. are called intensive properties.
Heat Capacity
The increase of temperature is proportional to the heat transferred.
q = coeff X T
The coefficient, C is called the heat capacity and it is directly proportional to amount of substance.
Molar Heat Capacity
The molar heat capacity of a substance,
$$\text{C}_{m}=\bigg(\frac{c}{n}\bigg)$$
is the heat capacity for one mole of the substance and is the quantity of heat needed to raise the temperature of one mole by one degree celsius (or one kelvin).
Specific Heat Capacity
It is the quantity of heat required to raise the temperature of one unit mass of a substance by one degree celsius (or one kelvin).
q = C x m x ∆T
Where, m = mass of the substance
∆T = rise in temperature.
The Relationship Between Cp And Cv For An Ideal Gas
At constant volume heat capacity = Cv
At constant pressure heat capacity = Cp
At constant volume qv = Cv ∆T = ∆U
At constant pressure qp= Cp∆T = ∆H
For one mole of an ideal gas
∆H = ∆U + ∆ (PV) = ∆U + ∆ (RT)
∆H = ∆U + R∆T
On subsitituting the values of ∆H and ∆U, the equation is modified as
Cp ∆T = Cv ∆T + R∆T or Cp -Cv = R
Measurement Of ∆U And ∆H: Calorimetry
- At constant volume, qv
- At constant pressure, qp
∆H Measurements
In an exothermic reaction, heat is evolved, and system loses heat to the surroundings. Therefore, qp will be negative and ∆rH will also be negative. Similarly in an endothermic reaction, heat is absorbed, qp is positive and ∆rH will be positive.
Hess’s Law Of Constant Heat Summation
It states that, “If a reaction takes place in several steps then its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate reactions into which the overall reaction may be divided at the same temperature.”
Standard Enthalpy Of Reactions
The standard enthalpy of reaction is the enthalpy change for a reaction when all the participating substances are in their standard states. The standard state of a substance at a specified temperature is its pure form at 1 bar.
Enthalpy Changes During Phase Transformations
- The enthalpy change that accompanies melting of one mole of a solid substance in standard state is called standard enthalpy of fusion or molar enthalpy of fusion, ∆fusH°.
- The amount of heat required to vaporize one mole of a liquid at constant temperature and under standard pressure (1bar) is called its standard enthalpy of vaporization or molar enthalpy of vaporization, ∆vapH°.
- The standard enthalpy of sublimation, ∆subH° is the change in enthalpy when one mole of a solid substance sublimes at a constant temperature and under standard pressure (1bar).
- The standard enthalpy change for the formation of one mole of a compound from its elements in their most
stable states of aggregation (also known as reference states) is called Standard Molar Enthalpy of Formation, ∆fH°.
Thermochemical Equation
A balanced chemical equation together with the value of its ∆rH is called a thermochemical equation. Conventions regarding thermochemical equations :
- The coefficients in a balanced thermochemical equation refer to the number of moles (never molecules) of
reactants and products involved in the reaction. - The numerical value of ∆rH° refers to the number of moles of substances specified by an equation. Standard enthalpy change ∆rH° will have units as kJ mol–1.
Enthalpy Change, ∆rH Of A Reaction – Reaction Enthalpy
The enthalpy change accompanying a reaction is called the reaction enthalpy.
∆rH = (sum of enthalpies of products) – (sum of enthalpies of reactants)
= ∑ai Hproducts – ∑bi Hreactants
Here symbol ∑ (sigma) is used for summation and ai and bi are the stoichiometric coefficients of the products
and reactants respectively in the balanced chemical equation.
Spontaneity
A spontaneous process is an irreversible process and may only be reversed by some external agency. A few examples of spontaneous process are:
- Common salt dissolves in water of its own.
- Carbon monoxide is oxidised to carbon dioxide of its own.
Entropy
Since entropy is a state property, the change in entropy of a reversible process is given as,
$$\Delta \text{S}_{\text{sys}}=\frac{a_{\text{sys, rev}}}{\text{T}}$$
For reversible and irreversible expansion for an ideal gas, under isothermal conditions, ∆U = 0, but ∆Stotal i.e., ∆Ssys + ∆Ssurr is not zero for irreversible process. Thus, ∆S does discriminate between reversible and irreversible process.
Lattice Enthalpy
The lattice enthalpy of an ionic compound is the enthalpy change which occurs when one mole of an ionic compound dissociates into its ions in gaseous state.
it is impossible to determine lattice enthalpies directly by experiment, therefore, an indirect method is used which involves constructing an enthalpy diagram called a Born-Haber Cycle.
Entropy And Spontaneity
The entropy is a measure of degree of randomness or disorder of a system. Entropy of a substance is minimum in solid state while it is maximum in gaseous state. The change in entropy in a spontaneous process is expressed as ∆S. It is related with q and T for a reversible reaction as :
$$\Delta \text{S}=\frac{\text{q}_{\text{rev}}}{\text{T}}$$
The total entropy change ( ∆Stotal) for the system and surroundings of a spontaneous process is given by,
∆Stotal = ∆Ssystem + ∆Ssurr > 0
When a system is in equilibrium, the entropy is maximum, and the change in entropy, ∆S = 0.
Gibbs Energy And Spontaneity
A new thermodynamic function, the Gibbs energy or Gibbs function G, can be defined as G = H-TS
∆G = ∆H - T∆S
Gibbs energy change = enthalpy change - temperature × entropy change ∆G gives a criteria for spontaneity at constant pressure and temperature,
- If ∆G is negative (<0) the process is spontaneous.
- If ∆G is positive (>0) the process is non-spontaneous
- Free energy change is Reversible Reaction.
Laws Of Thermodynamics
- First Law of Thermodynamics: It states that “The energy of an isolated system is constant”. It is commonly stated as the law of conservation of energy, i.e., energy can neither be created nor be destroyed.
- Entropy and Second Law of Thermodynamics: It states that, “for an isolated system the change in
energy remains constant. Therefore, increase in entropy in such systems is the natural direction of a
spontaneous change”. - Absolute Entropy and Third Law of Thermodynamics: The entropy of any pure crystalline substance
approaches zero as the temperature approaches absolute zero. This is called third law of thermodynamics.
