Course Topics

Introduction to Thermochemistry

Thermochemistry: The study of heat changes during a chemical reaction.

Energy: “Ability of a body to do work is called energy.”

Units Relation
Joule (SI Unit) J = kg·m²·s⁻²
Calorie 1 cal = 4.184 J
CRITICAL CONCEPT: Energy is never created or destroyed, only converted from one form to another (First Law of Thermodynamics).

Types of energies:

  • Kinetic energy (K.E): Energy due to motion
  • Potential energy (P.E): Energy due to position, shape and orientation
Memory Tip

Thermochemistry = Thermo (heat) + Chemistry = Study of heat in chemical reactions!

Exothermic vs Endothermic Reactions

Exothermic Reactions

  • Heat is evolved from the system
  • HR > HP (Reactants have higher heat content)
  • ∆H = -ive (Negative)
  • More bonds formed than broken
  • Heat flows from system to surroundings
  • Most spontaneous reactions are exothermic

Endothermic Reactions

  • Heat is absorbed by the system
  • HP > HR (Products have higher heat content)
  • ∆H = +ive (Positive)
  • More bonds broken than formed
  • Heat flows from surroundings to system
  • Most non-spontaneous reactions are endothermic
∆H = HProducts – HReactants
CRITICAL CONCEPT: The sign of ∆H determines if a reaction is exothermic (negative) or endothermic (positive). Bond breaking requires energy (endothermic), while bond formation releases energy (exothermic).
Memory Tip

EXOthermic = EXits heat (gives off heat), ENDOthermic = ENters heat (absorbs heat)!

System, Surrounding & State Functions

System: “The portion of universe which is set aside for consideration, observation, discussion, argumentation or experimentation.”

Types of Systems:

  • Open System: Can exchange both matter and energy with surroundings
  • Closed System: Can exchange only energy (not matter) with surroundings
  • Isolated System: Cannot exchange either matter or energy with surroundings

Surrounding: The portion of universe except system.

State Function: A macroscopic property of a system which has definite values for initial and final states, independent of the path.

State Functions (Path Independent) Path Functions (Path Dependent)
  • Pressure (P)
  • Temperature (T)
  • Volume (V)
  • Internal Energy (E)
  • Enthalpy (H)
  • Entropy (S)
  • Gibbs Free Energy (G)
  • Heat (q)
  • Work (w)
SYSTEM
Heat flow (q)
SURROUNDINGS
CRITICAL CONCEPT: State functions depend only on the initial and final states, not on how the change occurred. This is why we can use Hess’s Law to calculate enthalpy changes.
Memory Tip

State functions = PVT, H, E, S, G (Path Very Tricky, However Every Student Gets)!

Bond Energy

Bond Energy: The average amount of energy required to break all bonds of a particular type in one mole of the substance.

Unit: kJ/mol

ΔH = Σ(Bond energies of bonds broken) – Σ(Bond energies of bonds formed)

Factors affecting bond energy:

  1. Electronegativity difference of bonded atoms
  2. Sizes of the atoms
  3. Bond length

Applications of bond energy:

  • Relative strength of bonds
  • % of ionic character in bond
  • Estimation of ∆H for reactions
Bond Average Bond Energy (kJ/mol)
H-H 436
C-C 347
C-H 413
O=O 498
C=O 799
O-H 463
CRITICAL CONCEPT: If bonds formed are stronger than bonds broken, the reaction is exothermic (∆H negative). If bonds formed are weaker than bonds broken, the reaction is endothermic (∆H positive).
Memory Tip

Stronger bonds = Higher bond energy = More stable = Requires more energy to break!

Internal Energy & First Law of Thermodynamics

Internal Energy (E): “The sum of all the possible kinds of energies of a system.”

Mathematically: E = K.E + P.E + …

First Law of Thermodynamics: “Energy can neither be created nor be destroyed, but can be changed from one form to another.”

ΔE = q + w

Where:

  • ΔE = Change in internal energy
  • q = Heat transferred
  • w = Work done
CRITICAL CONCEPT:
  • At constant volume: ΔE = qv (no work done)
  • At constant pressure: ΔH = qp (enthalpy change equals heat at constant pressure)
  • By definition, standard enthalpy of an element is zero

Internal energy is comprised of:

  1. Kinetic energy: Translational, vibrational, rotational
  2. Potential energy: Intra-molecular and inter-molecular forces
Memory Tip

First Law = Conservation of energy = Energy in = Energy out + Energy stored!

Enthalpy & Its Types

Enthalpy (H): H = E + PV

Where E is internal energy, P is pressure, V is volume

Type of Enthalpy Definition Symbol Sign of ∆H
Enthalpy of Reaction Heat change when reactants form products ∆Hrxn ±
Enthalpy of Formation Heat change when 1 mole compound forms from elements ∆Hf ±
Enthalpy of Combustion Heat change when 1 mole substance burns completely in oxygen ∆Hc –
Enthalpy of Neutralization Heat change when 1 mole H⁺ reacts with 1 mole OH⁻ ∆Hn –
Enthalpy of Solution Heat change when 1 mole solute dissolves in solvent ∆Hsol ±
Enthalpy of Atomization Heat change when 1 mole gaseous atoms form from element ∆Hat +
For strong acid + strong base: ∆Hn = -57.4 kJ/mol
CRITICAL CONCEPT:
  • ∆Hn = -57.4 kJ/mol (Maximum for strong acids and bases)
  • ∆Hn < -57.4 (When one is strong and other is weak)
  • ∆Hn > -57.4 (Not possible)
Memory Tip

Enthalpy = Heat content at constant pressure = H = E + PV!

Measurement of Enthalpy (Calorimetry)

Calorimetry: Experimental measurement of heat changes.

Glass Calorimeter

  • Used for ∆Hn and ∆Hsol
  • Constant pressure
  • Formula: q = m × s × ∆T
  • q = heat, m = mass, s = specific heat, ∆T = temperature change
  • At constant pressure: ∆H = m × s × ∆T

Bomb Calorimeter

  • Used for ∆Hc of food and fuel
  • Constant volume
  • Formula: q = C × ∆T
  • q = heat, C = heat capacity, ∆T = temperature change
  • At constant volume: ∆E = C × ∆T
q = m × s × ∆T (Glass Calorimeter)
q = C × ∆T (Bomb Calorimeter)
CRITICAL CONCEPT: In glass calorimeter, pressure remains constant so q = ∆H. In bomb calorimeter, volume remains constant so q = ∆E. Specific heat capacity (s) has units J·K⁻¹·g⁻¹, while heat capacity (C) has units kJ·K⁻¹.
Memory Tip

Glass calorimeter = Constant P (Pressure), Bomb calorimeter = Constant V (Volume)!

Hess’s Law of Constant Heat Summation

Hess’s Law: “Energy contents of a reaction remain constant whether reaction takes place in single step or in multi-steps.”

If: A → B → C
Then: ∆HA→C = ∆HA→B + ∆HB→C

Key Points:

  • Total enthalpy change is independent of the pathway
  • Allows calculation of enthalpy changes that cannot be measured directly
  • The sum of enthalpy changes in a cyclic process is zero
CRITICAL CONCEPT – Example: Calculate enthalpy of formation of CO from combustion data:
  1. C(graphite) + O₂(g) → CO₂(g) ∆H = -393.7 kJ/mol
  2. CO(g) + ½O₂(g) → CO₂(g) ∆H = -283 kJ/mol
  3. From Hess’s Law: ∆Hf(CO) = -393.7 – (-283) = -110.7 kJ/mol

Applications:

  • Determining enthalpy of formation
  • Calculating bond energies
  • Finding enthalpy of reactions not easily measured
Memory Tip

Hess’s Law = The enthalpy change from A to C is the same whether you go directly or through B!

Born-Haber Cycle

Born-Haber Cycle: A thermodynamic cycle that enables calculation of lattice energy of ionic compounds.

∆Hf = ∆Hat + IE + EA + ∆Hlatt

Where:

  • ∆Hf = Enthalpy of formation
  • ∆Hat = Enthalpy of atomization
  • IE = Ionization energy
  • EA = Electron affinity
  • ∆Hlatt = Lattice energy
Ionic Compound Lattice Energy (kJ/mol)
NaF -895
NaCl -787
NaBr -728
NaI -690
KCl -690
CRITICAL CONCEPT – For NaCl:
  1. Na(s) → Na(g) ∆Hat = +109 kJ/mol
  2. Na(g) → Na⁺(g) + e⁻ IE = +496 kJ/mol
  3. ½Cl₂(g) → Cl(g) ½∆Hat = +122 kJ/mol
  4. Cl(g) + e⁻ → Cl⁻(g) EA = -349 kJ/mol
  5. Na⁺(g) + Cl⁻(g) → NaCl(s) ∆Hlatt = -787 kJ/mol
  6. Sum: ∆Hf = -411 kJ/mol
Memory Tip

Born-Haber cycle = Bookkeeping of energy changes in ionic compound formation!

Intensive vs Extensive Properties

Intensive Properties

  • Independent of amount of substance
  • Examples:
    • Temperature
    • Pressure
    • Density
    • Refractive index
    • Viscosity
    • Boiling point
    • Freezing point
    • Surface tension

Extensive Properties

  • Dependent on amount of substance
  • Examples:
    • Mass
    • Volume
    • Internal energy
    • Enthalpy
    • Entropy
    • Heat capacity
CRITICAL CONCEPT: Intensive properties are useful for identifying substances (like density, melting point). Extensive properties are additive – if you double the amount, you double the property (like mass, volume).

Factors Affecting Heat of Reaction:

  1. Nature or physical state of reactant and products
  2. Allotropic forms of the element
  3. Enthalpies of solution
Memory Tip

INTensive = INdependent of amount, EXTensive = EXTends with amount!

Applications & Importance

Practical Applications of Thermochemistry:

  • Fuel Efficiency: Calculating energy content of fuels
  • Food Industry: Calorific value of food (nutrition labels)
  • Industrial Processes: Optimizing chemical reactions for energy efficiency
  • Battery Technology: Designing batteries with optimal energy density
  • Environmental Science: Studying heat effects in climate change
  • Material Science: Predicting stability of compounds
  • Pharmaceuticals: Designing drug synthesis with optimal energy requirements
  • Construction: Using exothermic reactions in cement setting
CRITICAL CONCEPT: Thermochemistry provides the foundation for understanding energy changes in all chemical processes. It’s essential for designing industrial processes, developing new materials, and understanding biological energy transformations.

Example Reactions:

Reaction Type ∆H (kJ/mol)
H₂(g) + ½O₂(g) → H₂O(l) Exothermic -285.5
C(s) + O₂(g) → CO₂(g) Exothermic -393.7
N₂(g) + O₂(g) → 2NO(g) Endothermic +180.51
H₂O(l) → H₂O(g) Endothermic +44
Memory Tip

From car engines (combustion) to our bodies (respiration) – thermochemistry explains energy in everyday life!