5. Chemical energetics

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.

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.

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.

Use the label that matches the process: standard enthalpy of formation, ΔHf⦵, forms one mole of a compound from its elements in their standard states; standard enthalpy of combustion, ΔHc⦵, completely burns one mole of a substance in oxygen; standard enthalpy of neutralisation, ΔHneut⦵, forms one mole of water when an acid and an alkali react under the stated conditions.

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.

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.

Average bond enthalpies estimate the energy change for a reaction by comparing the energy needed to break the bonds in the reactants with the energy released when new bonds form in the products. Use the balanced equation and count bonds, not whole molecules.

Use this fixed sequence:

1. Balance the equation and draw or inspect the relevant structural changes.
2. Count every reactant bond broken, including stoichiometric multiplicity.
3. Count every product bond formed, including multiplicity.
4. Align the two totals in the bond-energy relationship and substitute the supplied values.
5. Report the result with its energy unit and reaction amount.
6. Check the sign and whether the counted bonds match the equation.

Worked-example scaffold (use the supplied equation and data):

• Reactant bonds broken: [bond type] × [coefficient] = [energy total].
• Product bonds formed: [bond type] × [coefficient] = [energy total].
• Substitute: ΔHᵣ ≈ [broken-bond total] − [formed-bond total] = [result] [unit].
• Interpret the sign as net energy release or input, then state that the result is an estimate.

Average bond enthalpies are model values, so the result is approximate rather than an exact experimental enthalpy. Common errors are using an unbalanced equation, counting unchanged bonds, adding product bond energies instead of subtracting them, confusing bond order, or reporting no unit/sign.

A bond enthalpy is the enthalpy required to break one mole of a specified bond in gaseous molecules. Some values are exact for a particular chemical environment, but many table values are average bond enthalpies taken across different compounds.

The same nominal bond can have slightly different strengths because its surrounding atoms, bond order and molecular environment differ. An average value is therefore useful for estimating a reaction enthalpy but does not exactly describe every bond of that type in every molecule.

When average bond enthalpies are used in ΔHᵣ = Σ(bonds broken) − Σ(bonds formed), the result is an estimate based on the supplied averages. It is appropriate for comparing the direction or approximate size of energy change, not for claiming an exact experimental value.

Do not treat every tabulated bond enthalpy as exact, or interpret a difference from an experimental value automatically as a counting error. Keep the limitation visible: the equation and bond count may be correct while the average-bond model remains approximate.

In a calorimetry experiment, measure the temperature change of a known mass and use q=mcΔT to estimate heat transferred. Divide by moles reacting to obtain an enthalpy change per mole.

Record the sign correctly: if the surroundings warm, the reaction is exothermic and ΔH is negative. Account for the solution’s mass, specific heat capacity and the limiting amount reacted.

If 100 g of solution warms by 5.0 °C and c=4.18 J g⁻¹ K⁻¹, q=2090 J absorbed by the solution; the reaction released −2090 J before dividing by reacting moles.

The measured temperature change is not itself ΔH. Heat capacity, mass, units and the amount of substance all enter the calculation.

Hess’s law states that the enthalpy change for a reaction is independent of the route taken because enthalpy is a state function.

Manipulate known equations until they sum to the target equation. Reverse an equation and change the sign of ΔH; multiply an equation and multiply ΔH by the same factor.

If A→B has ΔH₁ and B→C has ΔH₂, then A→C has ΔH₁+ΔH₂. If a route uses C→B instead, subtract that step’s enthalpy.

Never add enthalpies without first checking that the chemical equations cancel correctly. The target equation, not the diagram’s layout, decides the final sign.

In an energy cycle, arrows represent enthalpy changes between the same states. Choose a route to the target, then combine the signed values along that route.

Write the target reaction explicitly, align intermediate species, and apply reversal/multiplication rules before calculating. A final sign check should match whether the target is exothermic or endothermic.

If direct A→C is unknown, but A→B=−80 kJ mol⁻¹ and B→C=+25 kJ mol⁻¹, then ΔH(A→C)=−55 kJ mol⁻¹.

Do not treat every number around a cycle as positive. A value belongs to its arrow direction; reversing that arrow reverses its sign.