Thermochemistry

Subject: Chemistry
Topic: 6
Cambridge Code: 0620 / 0971 / 5070

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Energy in Reactions

Exothermic Reactions

Definition: Release energy (heat) to surroundings

• Temperature increases
• Combustion, neutralization
• ΔH = negative
• Products lower energy than reactants

$\text{Reactants} → \text{Products} + \text{Heat}$

Examples:

• Burning fuel: CH₄ + 2O₂ → CO₂ + 2H₂O + heat
• Neutralization: HCl + NaOH → NaCl + H₂O + heat

Endothermic Reactions

Definition: Absorb energy from surroundings

• Temperature decreases
• Photosynthesis, melting, evaporation
• ΔH = positive
• Products higher energy than reactants

$\text{Reactants} + \text{Heat} → \text{Products}$

Examples:

• Photosynthesis: 6CO₂ + 6H₂O + light → C₆H₁₂O₆ + 6O₂
• Melting ice: Ice + heat → Water

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Enthalpy Change (ΔH)

Enthalpy (H) - Total heat content

Enthalpy change (ΔH): $ΔH = H_{\text{products}} - H_{\text{reactants}}$

Sign Convention

• ΔH < 0: Exothermic (heat released)
• ΔH > 0: Endothermic (heat absorbed)

Units

• J/mol (joules per mole)
• kJ/mol (kilojoules per mole)
• Often written: ΔH = -393 kJ/mol

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Calorimetry

Calorimetry - Measure heat change in reactions

Heat Equation

$q = mc\Delta T$

where:

• q = heat energy (joules)
• m = mass (grams)
• c = specific heat capacity (J/g°C)
• ΔT = temperature change (°C)

Common Specific Heat Capacities

Substance	c (J/g°C)
Water	4.18
Aluminum	0.897
Iron	0.449
Copper	0.385

Example Calculation

Problem: 100 g water heated from 20°C to 30°C (c = 4.18)

$q = 100 \times 4.18 \times (30-20)$ $q = 100 \times 4.18 \times 10 = 4180 \text{ J}$

Bomb Calorimeter

• Insulated chamber
• Combustion occurs in sealed container
• Temperature rise measured
• Accurate for combustion reactions

Simple Calorimeter

• Beaker with thermometer
• Less accurate (heat loss to surroundings)
• Laboratory use

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Hess's Law

Hess's Law - Enthalpy change independent of reaction pathway

$ΔH_{\text{overall}} = \sum ΔH_{\text{equations}}$

Application

Find ΔH of difficult reaction using known reactions

Example: Find ΔH for: C(s) + O₂(g) → CO₂(g)

Given:

1. C(s) + ½O₂(g) → CO(g), ΔH₁ = -110.5 kJ
2. CO(g) + ½O₂(g) → CO₂(g), ΔH₂ = -283.0 kJ

Solution:

• Add equations 1 + 2:
• C(s) + O₂(g) → CO₂(g), ΔH = -110.5 + (-283.0) = -393.5 kJ

Reversing Equations

If reverse equation: ΔH changes sign

Example:

• Forward: H₂ + ½O₂ → H₂O, ΔH = -286 kJ
• Reverse: H₂O → H₂ + ½O₂, ΔH = +286 kJ

Scaling Equations

If multiply by factor n: ΔH also multiplies by n

Example:

• Given: N₂ + 3H₂ → 2NH₃, ΔH = -92 kJ
• Multiply by 2: 2N₂ + 6H₂ → 4NH₃, ΔH = -184 kJ

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Standard Enthalpy Change of Formation (ΔHf°)

ΔHf° - Enthalpy change when 1 mole of compound forms from elements in standard state

Standard state: Most stable form at 25°C, 1 atm

Values

ΔHf° = 0 for elements in standard state

• O₂(g) has ΔHf° = 0
• C(s) has ΔHf° = 0

ΔHf° for compounds: Tabulated values available

Calculating ΔH°reaction

$ΔH°_{\text{reaction}} = \sum (ΔH°_f \text{ products}) - \sum (ΔH°_f \text{ reactants})$

Example: CH₄ + 2O₂ → CO₂ + 2H₂O

Using standard enthalpies of formation

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Breaking and Forming Bonds

Bond Energy

Bond energy - Energy needed to break bond or released when formed

Energy Calculation

$ΔH = \sum \text{(bond energies broken)} - \sum \text{(bond energies formed)}$

Example: H₂ + Cl₂ → 2HCl

• Bonds broken: H-H (436 kJ) + Cl-Cl (244 kJ) = 680 kJ
• Bonds formed: 2 × H-Cl (432 kJ) = 864 kJ
• ΔH = 680 - 864 = -184 kJ (exothermic)

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Fuel Value

Fuel value - Energy released per mass of fuel

$\text{Fuel value} = \frac{\text{Energy released (kJ)}}{\text{Mass of fuel (g)}}$

Examples

Fuel	Value (kJ/g)
Hydrogen	141
Methane	55.5
Ethanol	29.6
Coal	30

Higher value = Better fuel

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Key Points

1. Exothermic: ΔH < 0 (heat released)
2. Endothermic: ΔH > 0 (heat absorbed)
3. Heat equation: q = mcΔT
4. Hess's Law: ΔH independent of pathway
5. ΔHf° for elements = 0
6. Bond energy: Breaking requires energy, forming releases
7. Fuel value: Energy per unit mass

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Practice Questions

1. Classify reactions as exothermic/endothermic
2. Calculate heat using q = mcΔT
3. Apply Hess's Law to calculate ΔH
4. Use ΔHf° values to find ΔH°reaction
5. Calculate from bond energies
6. Compare fuel values

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Revision Tips

• Know exothermic/endothermic clearly
• Practice calorimetry calculations
• Understand Hess's Law
• Learn ΔHf° concept
• Know bond energy calculations
• Know equation manipulation rules
• Practice multiple approaches