5.1 Enthalpy change, ΔH
- Syllabus
- 9701–2028–2029
- Topic
- 5.1
- Level
- AS
Learning objectives
5.1.1Chemical reactions are accompanied by enthalpy• Understand: chemical reactions are accompanied by enthalpy changes and these changes can be exothermic (ΔH is negative) or endothermic (ΔH is positive)5.1.2Construct/interpret a reaction pathway diagram• Construct/interpret: a reaction pathway diagram, in terms of the enthalpy change of the reaction and of the activation energy5.1.3Terms• Define/use:: (a) standard conditions (this syllabus assumes that these are 298 K and 101 kPa) shown by ⦵; (b) enthalpy change with particular reference to: reaction, ΔHr , formation, ΔHf , combustion, ΔHc , neutralisation, ΔHneut5.1.4Energy transfers occur during chemical• Understand: energy transfers occur during chemical reactions because of the breaking and making of chemical bonds5.1.5Bond energies and ΔHr• Use bond energies (ΔH positive, i.e. bond breaking) to calculate enthalpy change of reaction, ΔHr5.1.6Some bond energies are exact and some bond• Understand: some bond energies are exact and some bond energies are averages5.1.7Enthalpy changes from appropriate experimental• Calculate enthalpy changes from appropriate experimental results, including the use of the relationships q = mcΔT and ΔH = –mcΔT/n
Enthalpy change measures energy transferred in a chemical reaction
An enthalpy change, ΔH, is the energy transferred during a chemical reaction under stated conditions and for a stated amount of reaction. It is a property of the reaction as written, not an undefined amount of heat detached from the equation.
In an exothermic reaction, energy is transferred from the reacting system to the surroundings, so ΔH is negative and the surroundings warm. The products are at lower enthalpy than the reactants for the stated reaction direction.
In an endothermic reaction, energy is transferred from the surroundings to the reacting system, so ΔH is positive and the surroundings cool. The products are at higher enthalpy than the reactants for the stated reaction direction.
Always attach the sign to the reaction direction and amount shown. Do not treat a negative ΔH as energy created, a positive ΔH as a faster reaction, or the sign convention as independent of whether the reaction is written forwards or reversed.
Reaction pathway diagrams show energy changes from reactants to products
Read a reaction pathway diagram with energy or enthalpy on the vertical axis and reaction progress on the horizontal axis. The reactants and products are placed at their starting and ending energy levels for the reaction direction shown.
Find ΔH from the signed vertical difference between products and reactants: products lower than reactants gives a negative, exothermic ΔH; products higher gives a positive, endothermic ΔH. Reverse the reaction and the sign reverses.
The peak represents the activated complex. The vertical distance from the reactants to the peak is the forward activation energy; for the reverse direction, measure from the products to the same peak. A catalyst provides an alternative pathway with lower activation energy but does not change the reactant/product energy difference or ΔH.
Use an ordered read-off: identify axes and reaction direction, mark reactant and product levels, calculate the signed product-minus-reactant gap for ΔH, then measure the relevant reactant-or-product-to-peak gap for activation energy. Keep energy changes and reaction rate as separate conclusions.
Do not call the peak-to-product distance the forward activation energy, infer a faster reaction from ΔH alone, or claim that a catalyst changes equilibrium energy levels. Use only distances with clearly identified endpoints and direction.
Standard enthalpy definitions require a specified state and amount
A standard enthalpy change is measured for substances in their standard states under the syllabus standard conditions: 298 K and 101 kPa. The standard-state symbol is written as ⦵. Always state the physical state and the amount of reaction defined by the balanced equation.
The standard enthalpy change of reaction, ΔHᵣ⦵, is the enthalpy change for the reaction as written under standard conditions. Reversing the equation reverses the sign; multiplying the equation changes the enthalpy change by the same factor.
| Standard enthalpy change | Definition for the stated equation |
|---|---|
| formation, ΔHf⦵ | formation of one mole of a compound from its elements in their standard states |
| combustion, ΔHc⦵ | complete combustion of one mole of a substance in oxygen |
| neutralisation, ΔHneut⦵ | formation of one mole of water when an acid and an alkali react |
Before using a named value, check the equation, stoichiometric amount, physical states and standard conditions. Do not omit state symbols or silently change the amount of substance represented by the definition.
Energy transfers can be tracked by breaking bonds and forming bonds
Breaking a chemical bond separates bonded atoms and requires energy, so it is an energy input. In a reaction, identify the bonds present in the reactants that must be broken before the atoms can be rearranged.
Forming a chemical bond releases energy as the new bonded arrangement is established. Identify the new bonds in the products and treat their formation as an energy output.
The reaction enthalpy is the net result of these transfers: energy required to break bonds is balanced against energy released when bonds form. More energy released on forming products gives an exothermic result; more energy required to break reactant bonds gives an endothermic result.
Use a balanced equation to track which bonds change, but do not claim that the sign of ΔH alone identifies a reaction mechanism or rate. The bond-energy calculation route belongs to the next objective; this card establishes the causal energy account.
Calculate ΔHᵣ as bond-breaking input minus bond-forming output
ΔHr≈∑E(bonds broken)−∑E(bonds formed)
Balance the equation, draw enough structure to count every bond, multiply each bond count by its supplied bond energy, then subtract the formed-bond total from the broken-bond total. Bond breaking values are positive inputs; the subtraction accounts for energy released by bond formation.
For CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g), use C–H 413, O=O 498, C=O in CO₂ 805 and O–H 464 kJ mol⁻¹.
| Account | Bonds counted | Total / kJ mol⁻¹ |
|---|---|---|
| broken | 4(C–H) + 2(O=O) | 4(413) + 2(498) = 2648 |
| formed | 2(C=O) + 4(O–H) | 2(805) + 4(464) = 3466 |
ΔHr≈2648−3466=−818 kJ mol−1
The negative result means the energy released on forming product bonds exceeds the energy required to break reactant bonds. It is an estimate because the supplied bond energies represent gaseous bonds and may be averages.
Count bonds, not molecules, and include coefficients. Do not add the formed-bond total, change the formula sign to force an expected answer, or use liquid-water data in a calculation whose average-bond model represents gaseous molecules.
Exact and average bond energies describe different levels of specificity
Bond energy is the enthalpy required to break one mole of a covalent bond in gaseous molecules; bond breaking therefore has positive ΔH. An exact bond energy refers to one specified bond in one defined gaseous species, whereas an average bond energy combines values for the same bond type in different molecular environments.
| Value | What it represents | Consequence in a calculation |
|---|---|---|
| exact bond energy | a specified bond in a specified gaseous molecule | specific to that bond and environment |
| average bond energy | a mean for that bond type across several gaseous compounds | useful beyond one compound, but gives an approximate ΔHᵣ |
A C–H bond, for example, does not have exactly the same strength in every molecule because neighbouring atoms and electron distribution alter its environment. The table value is often averaged so it can be reused across compounds.
A difference between a bond-energy estimate and an experimental enthalpy does not automatically prove that the bond count is wrong. First check the equation and arithmetic, then retain the model limitation: average bond energies cannot reproduce every exact molecular environment.
Calorimetry converts ΔT into ΔH per mole with the opposite sign
q=mcΔTΔH=−nmcΔT
| Symbol | Meaning | Typical unit |
|---|---|---|
| m | mass of solution or water heated | g |
| c | specific heat capacity | J g⁻¹ K⁻¹ |
| ΔT | final temperature − initial temperature | K or °C interval |
| n | amount of reaction specified, usually limiting amount | mol |
Find the mass that actually changes temperature; for dilute aqueous solutions, total volume in cm³ is often treated as the same numerical mass in g using density 1.00 g cm⁻³. Calculate q for the surroundings, identify the reacting amount n from concentrations and stoichiometry, apply the opposite sign for the reaction, and convert J mol⁻¹ to kJ mol⁻¹.
Mixing 50.0 cm³ of 1.00 mol dm⁻³ HCl with 50.0 cm³ of 1.00 mol dm⁻³ NaOH raises the temperature by 6.80 °C. With m = 100 g and c = 4.18 J g⁻¹ K⁻¹, q = 100 × 4.18 × 6.80 = 2842 J. The amount neutralised is 0.0500 mol, so ΔHneut = −2842 ÷ 0.0500 ÷ 1000 = −56.8 kJ mol⁻¹.
Heat lost to the surroundings, heat absorbed by the cup, incomplete combustion or reaction, evaporation, and the water-like density and heat-capacity assumptions can move an experimental value away from the accepted value. A lid, insulation and temperature extrapolation reduce some losses.
ΔT is not ΔH, and q for the solution has the opposite sign to ΔH for the reaction. Use the total heated mass, the limiting reacting amount and the stoichiometric definition—not whichever reagent volume or mole value is easiest to copy.