Ch 1 · Enthalpy change, ΔH · AS & A Level Chemistry 9701
Topic 1 of 2 in Chemical energetics — AS & A Level Chemistry 9701 — 53 lessons.
Introduction
Every chemical reaction involves an energy change as bonds in the reactants break and new bonds form in the products. Cambridge 9701 requires you to quantify this change as an enthalpy change, ΔH, measured under standard conditions, and to link it to activation energy on a reaction pathway diagram. This lesson also builds the two calculation methods used throughout energetics: bond energy summations and experimental temperature-change data.
Key ideas — Exothermic and endothermic changes
A reaction is exothermic if it transfers heat energy to the surroundings, so the temperature of the surroundings rises; the enthalpy of the products is lower than that of the reactants, and Δ H is negative. A reaction is endothermic if it takes in heat energy from the surroundings, so the surroundings cool; the products have higher enthalpy than the reactants, and Δ H is positive. Combustion and neutralisation are always exothermic; thermal decomposition is usually endothermic.
🧑🏫 Teacher's explanation
Think of a chemical reaction like a financial transaction. In an exothermic reaction, the system 'pays out' energy to the surroundings, like spending money; the surroundings get warmer because they've received that energy. In an endothermic reaction, the system 'borrows' energy from the surroundings, like taking out a loan; the surroundings get colder because their energy has been pulled into the reaction.
📖 Definitions
🔍 Worked example — Exam-style
Explain why the enthalpy change for the combustion of methane, CH_4(g) + 2O_2(g) arrow CO_2(g) + 2H_2O(l), is negative.
⚠️ Common errors
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🎯 Exam question
Key ideas — Reaction pathway diagrams
A reaction pathway (energy profile) diagram plots enthalpy on the vertical axis against the progress of reaction on the horizontal axis. The reactants and products are drawn as horizontal lines at their respective enthalpy levels, joined by a curve that rises to a maximum before falling (or continuing to rise) to the products. Δ H is the vertical distance between the reactants and products lines, with an arrow and sign to show its direction. The activation energy, E_a, is the vertical distance from the reactants line up to the peak of the curve; it is the minimum energy that colliding particles need for the reaction to occur, and it is always positive, whether the overall reaction is exothermic or endothermic.
🧑🏫 Teacher's explanation
Imagine you are pushing a heavy boulder up a hill. Even if the boulder eventually rolls down into a much lower valley (releasing energy), you still have to exert a significant amount of effort just to get it over the first peak. In chemistry, that 'peak' is the activation energy. Even if a reaction releases energy overall, you must provide a 'kick' of energy to get the molecules reacting.
📖 Definitions
🔍 Worked example — Easy
In an exothermic reaction, the enthalpy of the reactants is 450 kJ mol^-1 and the enthalpy of the products is 320 kJ mol^-1. Calculate the enthalpy change, Δ H.
🔍 Worked example — Standard
A reaction has an activation energy (E_a) of 150 kJ mol^-1 for the forward reaction. If the enthalpy change (Δ H) is -200 kJ mol^-1, calculate the activation energy for the reverse reaction.
🔍 Worked example — Exam-style
A reaction pathway diagram shows a reactant level at 100 kJ mol^-1 and a peak (transition state) at 275 kJ mol^-1. Calculate the activation energy, E_a.
⚠️ Common errors
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🎯 Exam question
Key ideas — Standard conditions and standard enthalpy changes
Standard conditions are a pressure of 1 bar (100 kPa (formerly 101 kPa)) and a stated temperature, usually 298 K, with all solutions at a concentration of 1 mol dm⁻³; the standard symbol ominus shows these conditions apply. Four standard enthalpy changes are defined:
🧑🏫 Teacher's explanation
Imagine you are comparing the cost of two different meal deals. If one shop charges different prices depending on the time of day or the weather, it is hard to compare them fairly. In chemistry, temperature and pressure change how much energy a reaction releases. To make comparisons fair, scientists use 'Standard Conditions'—a set of fixed rules (like 1 bar pressure and 298 K temperature) so everyone is playing by the same rules.
📖 Definitions
🔍 Worked example — Standard
Calculate the Δ H_c^ominus for ethane (C_2H_6) given the following enthalpies of formation: Δ H_f^ominus[C_2H_6] = -84.7 kJ mol^-1, Δ H_f^ominus[CO_2] = -393.5 kJ mol^-1, and Δ H_f^ominus[H_2O] = -285.8 kJ mol^-1.
🔍 Worked example — Exam-style
Given the average bond enthalpies: C-H = 413 kJ mol^-1, C-C = 348 kJ mol^-1, C-O = 358 kJ mol^-1, O-H = 463 kJ mol^-1, and O=O = 460 kJ mol^-1. Calculate Δ H for the reaction: CH_4 + 2O_2 arrow CO_2 + 2H_2O. (Assume CO_2 is O=C=O)
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🎯 Exam question
Key ideas — Energy transfers: breaking and making bonds
Breaking a chemical bond always requires an input of energy, so bond breaking is endothermic; forming a chemical bond always releases energy, so bond making is exothermic. The overall enthalpy change of a reaction is the balance between the energy absorbed to break bonds in the reactants and the energy released on forming bonds in the products:
🧑🏫 Teacher's explanation
Think of chemical bonds like LEGO bricks that are snapped together. To take two bricks apart, you have to put in effort (energy); you can't just pull them apart for free. However, when you snap two bricks together, they 'click' into place, releasing a little bit of energy.
📖 Definitions
🔍 Worked example — Easy
Calculate the enthalpy change for the reaction: H_2(g) + Cl_2(g) arrow 2HCl(g). Given bond energies: H-H = 436 kJ/mol, Cl-Cl = 242 kJ/mol, and H-Cl = 431 kJ/mol.
🔍 Worked example — Standard
In a reaction, 850 kJ/mol of energy is absorbed to break the reactant bonds, and 1200 kJ/mol of energy is released when the product bonds are formed. Calculate Δ H.
🔍 Worked example — Exam-style
Calculate the enthalpy change for the combustion of methane: CH_4(g) + 2O_2(g) arrow CO_2(g) + 2H_2O(l). Use the following average bond enthalpies: C-H = 413 kJ/mol, O=O = 498 kJ/mol, C=O = 799 kJ/mol, O-H = 463 kJ/mol.
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🎯 Exam question
Key ideas — Exact and average bond energies
A bond energy is the energy needed to break one mole of a given covalent bond in the gaseous state. An exact bond energy applies to one specific bond in one specific molecule, for example the C–H bond in methane only. An average bond energy is a mean value taken across many different molecules containing that type of bond, since the exact energy of, say, a C–H bond varies slightly depending on the rest of the molecule. Calculations using average bond energies therefore give only an approximate value of Δ H, whereas calculations using exact bond energies (or Hess cycles from enthalpies of formation) give a more reliable value.
🧑🏫 Teacher's explanation
Think of chemical bonds like springs connecting two balls. To break a spring, you have to pull it apart, which requires an input of energy. When you snap the spring back together (forming a new bond), energy is released.
📖 Definitions
🔍 Worked example — Easy
Calculate the enthalpy change for the reaction: H_2(g) + Cl_2(g) arrow 2HCl(g) using the following average bond energies: H-H = 436 kJ/mol, Cl-Cl = 242 kJ/mol, H-Cl = 431 kJ/mol.
🔍 Worked example — Exam-style
Explain why calculations using average bond energies often differ from experimental values of Δ H.
⚠️ Common errors
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🎯 Exam question
Worked example
Question. In an experiment, 50.0 cm³ of water was heated in a calorimeter by burning 0.360 g of ethanol, C_2H_5OH, M_r = 46.0. The temperature of the water rose from 20.5 °C to 44.5 °C. The specific heat capacity of water is 4.18 J g⁻¹ K⁻¹ and its density is 1.00 g cm⁻³. Calculate the enthalpy change of combustion of ethanol. [4]
🔍 Worked example — Standard
Calculate the enthalpy change of combustion (Δ H_c) in kJ mol⁻¹ if 0.500 g of methanol (CH_3OH, M_r = 32.0) produces 12.5 kJ of heat when burned, and the temperature change of the water was measured via calorimetry.
🔍 Worked example — Exam-style
In an experiment, 50.0 cm³ of water was heated in a calorimeter by burning 0.360 g of ethanol, C_2H_5OH, M_r = 46.0. The temperature of the water rose from 20.5 °C to 44.5 °C. The specific heat capacity of water is 4.18 J g⁻¹ K⁻¹ and its density is 1.00 g cm⁻³. Calculate the enthalpy change of combustion of ethanol in kJ mol⁻¹.
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Practice
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