Thermochemistry

Thermochemistry, a branch of Physical Chemistry, focuses on quantifying the heat changes that occur during chemical reactions and phase transitions. These heat changes are primarily studied in terms of enthalpy, a measure of the total heat content of a system at constant pressure.

Thermochemistry is also commonly referred to as Chemical Energetics, emphasizing its concern with the energy transformations that accompany chemical processes. Through the analysis of enthalpy changes, thermochemistry provides insights into the energy dynamics of chemical reactions, including the amount of heat absorbed or released, as well as the factors influencing reaction spontaneity and equilibrium. This field plays a fundamental role in understanding and predicting the behavior of chemical systems, with applications ranging from industrial processes to environmental phenomena.

1.0Types of Reactions in terms of Heat

In terms of heat, chemical reactions can be further classified based on whether they absorb or release heat energy:

1. Endothermic Reactions: 

Endothermic reactions absorb heat energy from the surroundings. As a result, the surroundings feel colder, and the temperature of the system decreases. These reactions have a positive change in enthalpy (ΔH>0). An example is the thermal decomposition of calcium carbonate:

CaCO₃ (s)  +  heat  →  CaO(s)  +  CO₂(g)

A general representation of endothermic reactions can be in following manner-

A  +  B +  x Cal    ⟶   C  +  D

             A  +  B    ⟶    C  +  D - x Cal

             A  +  B    ⟶    C  +  D  ; ΔH= +x Cal

Examples of Endothermic reactions:

• Dissociation reactions (Mostly)
• Fusion reactions
• Vaporization reactions
• Sublimation reactions
• Photosynthesis

2. Exothermic Reactions: 

Exothermic reactions release heat energy into the surroundings. As a result, the surroundings feel warmer, and the temperature of the system increases. These reactions have a negative change in enthalpy (ΔH<0). 

An example is the combustion of methane:

CH₄(g)  +  2O₂(g)  →  CO₂(g)  +  2H₂O(g)  +  heat

A general representation of exothermic reactions can be in following manner:

              A  +  B - x Cal    ⟶    C  +  D

                        A  +  B     ⟶    C  +  D+  x Cal

                        A  +  B     ⟶    C  +  D  ; ΔH = -x Cal

Examples of Exothermic reactions:

• Combustion reactions
• Neutralisation reactions
• Respiration reaction

3. Thermoneutral reactions:

Thermoneutral reactions, also known as thermally balanced reactions or thermally neutral reactions, are chemical reactions where there is no net release or absorption of heat energy. In other words, the enthalpy change (ΔH) for such reactions is zero.

$\frac{1}{2}$ H₂ (g) +aq.  ⟶ H(aq.) ; ΔH^o = 0

2.0Heat of Reaction (Enthalpy Change Reactions)

These reactions involve the heat change associated with a chemical reaction at constant pressure. The heat of reaction, Δ_rH, is the difference in enthalpy between the products and reactants of a reaction. It can be either endothermic or exothermic depending on whether heat is absorbed or released during the reaction.

3.0Factors affecting Enthalpy of Reaction

1. Physical State of Reactants and Products: 

Changes in the physical state of reactants and products (solid, liquid, gas) can affect the enthalpy change. For example, reactions involving gas-phase species may have different enthalpy changes compared to reactions involving only solid or liquid phases.

2. Stoichiometry of a reaction:

The stoichiometry of a reaction can greatly affect the enthalpy change (ΔH) of the reaction. This is because the enthalpy change is typically given per mole of reaction, so the quantities of reactants and products involved directly impact the magnitude of the enthalpy change.

For Example:

•   A ⟶ B           Δ_rH = x kJ
• 2A ⟶ 2B         Δ_rH = 2x kJ

3. Allotropic Form:

When an element exists in different allotropes, the enthalpy change (ΔH) of reactions involving that element may vary depending on the allotrope present. This is because different allotropes can have different bond energies, lattice energies, or other thermodynamic properties, which in turn affect the overall enthalpy change of a reaction.

C (graphite)    + O₂ (g)  ⟶ CO₂  (g)     ΔH₁ = xJ

C (diamond)   + O₂ (g)  ⟶ CO₂  (g)     ΔH₂ = xJ

                         ΔH₁  > ΔH₁

Graphite has a layered structure consisting of hexagonally arranged carbon atoms that form sheets of graphene. These layers are held together by weak van der Waals forces, allowing them to slide over each other easily. This structure gives graphite its lubricating properties and makes it relatively soft.

Diamond, on the other hand, has a three-dimensional network structure in which each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement. These strong covalent bonds give diamond its hardness and make it one of the hardest known substances.

The difference in structure and bonding between graphite and diamond results in different enthalpies of formation and different Gibbs free energies. Graphite has a lower Gibbs free energy of formation compared to diamond under standard conditions, indicating that it is more thermodynamically stable.

That’s why graphite is more thermodynamically stable than diamond at standard conditions due to its lower Gibbs free energy of formation, which reflects its tendency to exist in a lower-energy state.

4. Reaction carried out at constant pressure or constant volume:

Constant Pressure (Isobaric) Conditions: 

When a reaction is carried out at constant pressure, such as in an open container or a reaction vessel open to the atmosphere, the enthalpy change of the reaction is equal to the heat absorbed or released by the system in thermochemistry. This is because the change in enthalpy, ΔH, is directly related to the heat exchanged at constant pressure. Enthalpy change under constant pressure conditions is commonly measured in laboratory experiments using a calorimeter.

Constant Volume (Isochoric) Conditions: 

When a reaction is carried out at constant volume, such as in a sealed container where the volume cannot change, the enthalpy change of the reaction is not directly equal to the heat exchanged. Instead, the enthalpy change corresponds to the internal energy change (ΔU) of the system in thermochemistry. 

In this case, the enthalpy change can be calculated using the equation:

Δ𝐻=Δ𝑈 + Δ(𝑃𝑉)

Where; Δ(PV) represents the work done by or on the system due to changes in pressure and volume. Under constant volume conditions,

Δ(PV) is typically negligible, so Δ𝐻  is approximately equal to Δ𝑈.

4.0Laws of Thermochemistry

Here are laws of thermochemistry we will discuss in brief.

1. Lavoisier and Laplace's law

Lavoisier and Laplace's law, also known as the law of the conservation of energy in thermochemistry, states that the total heat absorbed or evolved in a chemical process is equal to the total heat absorbed or evolved when the process is reversed.

                A + B ⟶ C + D           Δ_rH = x

                C + D ⟶ A + B           Δ_rH = -x kJ

2. Temperature Effect/Kirchoff’s Relation

Kirchhoff's equation explains how the enthalpy change (ΔH) of a reaction varies with temperature. For a reaction at constant pressure, it is given by:

                           ΔH_T   = ΔH_T +   T0TΔC_pdt

3. Hess's Law of Constant Heat Summation (Brief)

Hess's Law states that the total enthalpy change of a reaction is the same whether it occurs in one step or multiple steps. This is because enthalpy is a state function, depending only on the initial and final states, not the path taken.

• Statement: If a reaction occurs in several steps, the standard reaction enthalpy (ΔH^oreaction) is the sum of the standard enthalpies of the intermediate reactions at the same temperature.
• General representation

For a reaction:

A → B

If it occurs in steps:

A → C  (ΔH₁)

C → D  (ΔH₂)

D → B  (ΔH₃)

Then,

ΔHoverall​      =  ΔH₁  + ΔH₂  + ΔH₃

For example-

C(s)    +      O₂ (g)    →  CO₂(g)   ∆_rH = – 94.0 Kcal/mol

• Mechanism:

C(s)    +  1/2 O₂  (g)  → CO (g)   ∆H₁ = – 26.0 kCal/mol

CO (g) + 1/2  O₂ (g) → CO₂(g)  ∆H₂ = – 68.0 kCal/mol

C(s)    +      O₂(g)    →  CO₂(g)   ∆_rH =  –26 –68= – 94.0 Kcal/mol

5.0Types of Heat of Reactions

1. Heat of Formation (or Standard Enthalpy of Formation) Reactions:

These reactions involve the formation of one mole of a compound from its elements in their standard states with the release or absorption of a specific amount of heat energy. The standard enthalpy of formation, ΔH_f​, is defined as the heat change accompanying the formation of one mole of a compound from its elements in their standard states. 

For example, the formation of water from its elements:

H₂ (g)  +  0.5O₂ ​(g)  →  H₂O (l)  (ΔH_f^o = −285.8kJ/mol)

2. Heat of Combustion

The heat of combustion (also known as enthalpy of combustion) is the amount of heat released when one mole of a substance is completely burned in oxygen.

For example, the standard enthalpy of combustion of methane at 298.15 K is –890.36 kJ mol^–1. This implies the following reaction :

CH₄(g) + 2O₂(g)  → CO₂(g) + 2H₂O(l) ; ΔH° = – 890.36 kJ mol^–1

3. Heat of Neutralization

The heat of neutralization is defined as the heat evolved when one equivalent of an acid reacts completely with one equivalent of a base in a dilute aqueous solution to form water and a salt. This process typically occurs under standard conditions, with all reactants and products in their standard states.

For example:

Strong Acid + Strong Base →   Salt + Water

HCl (aq) + NaOH (aq) → NaCl + H₂O + 13.7 kCal

Note- When one equivalent of SA is neutralized by one equivalent of SB, then heat evolved remains constant and its value is – 13.7 kCal/equivalent or – 57.2 kJ equivalent^–1.

4. Heat of Hydrogenation

The heat of hydrogenation (ΔHHydrogenation) is the heat evolved during the complete hydrogenation of one mole of an unsaturated organic compound to form its saturated counterpart. This process typically involves the addition of hydrogen (H₂) to double or triple bonds in the presence of a catalyst, resulting in the conversion of alkenes or alkynes to alkanes.

Unsaturated organic compounds changes into → Saturated organic compounds

Heat of hydrogenation is an exothermic process.

For example:

C₂H₂ + 2H₂      →    C₂H₆,                 ΔH_hydro= –ve

C₂H₄ + H₂         →     C₂H₆,                ΔH_hydro= –ve

C₂H₂ + H₂         →    C₂H₄,          Heat of Hydrogenation is not possible