Chemical Reactions & Equilibrium
Thermochemistry
5th Year · 6th Year (Leaving Cert)
- ✓Define and distinguish between exothermic and endothermic reactions, including their energy profile diagrams.
- ✓Define standard enthalpy of formation (ΔH°f), combustion (ΔH°c), and neutralisation (ΔH°neut), and perform calculations involving them.
- ✓Apply Hess's Law to calculate enthalpy changes for reactions that cannot be measured directly (Higher Level).
- ✓Use average bond enthalpies to estimate enthalpy changes for reactions (Higher Level).
- ✓State and apply standard conditions (298 K, 1 atm) in thermochemical calculations.
Key concepts
Chemical reactions that release heat energy to the surroundings. The products have lower energy than the reactants. The temperature of the surroundings increases.
Chemical reactions that absorb heat energy from the surroundings. The products have higher energy than the reactants. The temperature of the surroundings decreases.
The heat change of a reaction carried out at constant pressure. It is the difference between the total energy of the products and the total energy of the reactants.
The enthalpy change that occurs when one mole of a compound is formed from its constituent elements in their standard states under standard conditions (298 K and 1 atm).
The enthalpy change that occurs when one mole of a substance is completely burned in oxygen under standard conditions (298 K and 1 atm).
The enthalpy change that occurs when one mole of water is formed from the reaction of an acid and a base under standard conditions (298 K and 1 atm). For strong acid/strong base, it is approximately -57.3 kJ mol⁻¹.
States that the total enthalpy change for a chemical reaction is independent of the pathway taken, provided the initial and final conditions are the same. It is a direct consequence of the Law of Conservation of Energy.
The energy required to break one mole of a specific type of bond in the gaseous state. Average bond enthalpies are used as they vary slightly depending on the molecule. Bond breaking is endothermic (+ve), bond formation is exothermic (-ve).
Key facts to remember
- 1Exothermic reactions release energy to the surroundings, resulting in a negative enthalpy change (ΔH < 0).
- 2Endothermic reactions absorb energy from the surroundings, resulting in a positive enthalpy change (ΔH > 0).
- 3Standard conditions for thermochemical data are 298 K (25 °C) and 1 atmosphere (101.3 kPa) pressure.
- 4The standard enthalpy of formation (ΔH°f) of an element in its most stable standard state is defined as zero.
- 5Hess's Law is a statement of the conservation of energy and allows for the calculation of enthalpy changes for reactions that are difficult to measure directly.
- 6Bond breaking is an endothermic process (requires energy), while bond formation is an exothermic process (releases energy).
- 7Calculations using average bond enthalpies provide an estimate of the enthalpy change, as bond energies vary slightly between different molecules.
Worked examples
Example 1
Classify the following reactions as exothermic or endothermic: (a) C(s) + O₂(g) → CO₂(g) ; ΔH = -393.5 kJ mol⁻¹ (b) H₂O(l) → H₂O(g) ; ΔH = +44.0 kJ mol⁻¹
Answer
(a) Exothermic (b) Endothermic
The sign of ΔH is crucial for classification.
Example 2
Calculate the standard enthalpy change for the combustion of methane, CH₄(g), given the following standard enthalpies of formation: ΔH°f [CH₄(g)] = -74.8 kJ mol⁻¹ ΔH°f [CO₂(g)] = -393.5 kJ mol⁻¹ ΔH°f [H₂O(l)] = -285.8 kJ mol⁻¹
Answer
The standard enthalpy change for the combustion of methane is -890.3 kJ mol⁻¹.
Ensure the equation is balanced and stoichiometric coefficients are applied correctly.
Example 3
Calculate the enthalpy change for the formation of ethyne (C₂H₂(g)) from its elements: 2C(s) + H₂(g) → C₂H₂(g) ; ΔH = ? Given the following combustion data: (1) C(s) + O₂(g) → CO₂(g) ; ΔH₁ = -393.5 kJ mol⁻¹ (2) H₂(g) + ½O₂(g) → H₂O(l) ; ΔH₂ = -285.8 kJ mol⁻¹ (3) C₂H₂(g) + 2½O₂(g) → 2CO₂(g) + H₂O(l) ; ΔH₃ = -1300.0 kJ mol⁻¹
Answer
The enthalpy change for the formation of ethyne is +227.2 kJ mol⁻¹.
Remember to change the sign of ΔH when reversing an equation and multiply ΔH when multiplying an equation.
Example 4
Estimate the enthalpy change for the reaction: H₂(g) + Cl₂(g) → 2HCl(g) Given the following average bond enthalpies: H-H: 436 kJ mol⁻¹ Cl-Cl: 243 kJ mol⁻¹ H-Cl: 432 kJ mol⁻¹
Answer
The estimated enthalpy change for the reaction is -185 kJ mol⁻¹.
Bond enthalpy calculations provide estimates and are typically for gaseous reactions. Ensure correct signs for bond breaking (+ve) and bond formation (-ve).
Common mistakes
- ✗Confusing the signs of ΔH: A common error is assigning a positive ΔH to an exothermic reaction or a negative ΔH to an endothermic reaction.
- ✗Not balancing chemical equations correctly before performing calculations, leading to incorrect stoichiometric coefficients.
- ✗Forgetting to multiply ΔH values by the stoichiometric coefficients when using Hess's Law or ΔH°f calculations.
- ✗Incorrectly applying the formula for bond enthalpy calculations (e.g., subtracting bonds broken from bonds formed, or using incorrect signs).
- ✗Failing to reverse the sign of ΔH when reversing a chemical equation in Hess's Law problems.
Exam tips
- ★Always include the sign (+ or -) and units (kJ mol⁻¹) with all enthalpy values in your answers.
- ★For Hess's Law problems, clearly show how you manipulate each given equation (e.g., "Eq. 1 × 2", "Reverse Eq. 3") and the corresponding change to its ΔH value.
- ★When using ΔH°f values, remember the formula: ΔH°reaction = ΣnΔH°f(products) - ΣmΔH°f(reactants).
- ★For bond enthalpy calculations, draw out the Lewis structures of reactants and products to clearly identify all bonds broken and formed.
- ★Practise a variety of problems, especially those involving Hess's Law and bond enthalpies, as these often require careful manipulation and attention to detail.
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