Summary — What This Chapter Covers

• Enthalpy level diagrams show whether reactants or products have more stored chemical energy, telling you if a reaction is exothermic (ΔH negative) or endothermic (ΔH positive).
• Reaction profile diagrams add the missing piece — activation energy and the unstable "transition state" at the peak of the curve.
• Standard conditions (100 kPa, 298 K, standard states) let us fairly compare enthalpy changes between different reactions.
• Four named enthalpy changes to know cold: reaction, formation, combustion, neutralisation — each with a precise one-mole definition.
• Calorimetry is how we measure enthalpy changes experimentally, using q = mcΔT for solutions and burning fuels to heat water for combustion.
• Hess's Law lets us calculate enthalpy changes we can't measure directly, by building "indirect route" energy cycles through elements or combustion products.
• Bond enthalpies give an alternative way to estimate ΔH: energy in to break bonds, energy out to make new ones.

1. Enthalpy Level Diagrams

Think of enthalpy as the total amount of chemical energy "stored" inside a substance — like the energy stored in a stretched spring or a full battery. We can never measure the absolute amount of this stored energy, but we can measure the change in it when a reaction happens. That change is called the enthalpy change, symbol ΔH (Δ = "change in", H = "enthalpy").

Here's the key mental model: picture enthalpy as a height on a vertical axis. Reactants sit at one height, products sit at another. The reaction is basically a trip from one height to the other, and ΔH is simply how far you moved and in which direction.

Exothermic Reactions (ΔH is negative)

In an exothermic reaction, the products end up with less energy than the reactants started with. Where did that "missing" energy go? It didn't disappear — it was released as heat into the surroundings. That's why exothermic reactions make the test tube feel warm: the chemical system is losing energy, and the surroundings are gaining it.

• Energy of the system decreases
• Temperature of the surroundings increases (you can feel/measure this)
• ΔH is written as a negative number

Endothermic Reactions (ΔH is positive)

Endothermic is the mirror image: products end up with more energy than the reactants. That extra energy has to come from somewhere — it's absorbed from the surroundings, which is why endothermic reactions often feel cold (think of the cooling packs used for sports injuries, which rely on an endothermic dissolving process).

• Energy of the system increases
• Temperature of the surroundings decreases
• ΔH is written as a positive number

Reaction Profile Diagrams

Enthalpy level diagrams only show the start and end points — they don't tell you anything about how the reaction gets from reactants to products. That's where reaction profile diagrams come in. They add a curved "hill" between reactants and products, and that hill introduces two new ideas:

A crucial link: because exothermic reactions start with reactants that are already closer in energy to the transition state (higher up the diagram), they generally have a lower activation energy than endothermic reactions, where the reactants sit further below the transition state.

2. Enthalpy Change — Definitions

To compare enthalpy changes fairly between different experiments (maybe done in different labs, on different days, at different pressures), chemists agree on a fixed set of standard conditions. Think of it like agreeing to always measure people's height while they're standing on flat ground, not on a hill — otherwise the comparison is meaningless.

The Four Named Enthalpy Changes

These four definitions are the backbone of the whole chapter — they show up constantly in Hess's Law questions, and examiners love asking you to identify which one an equation represents. Learn the exact wording, especially the phrase "one mole", because that's what makes each definition precise.

A fifth definition, used less at AS level but still worth knowing: Standard enthalpy of atomisation — the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state, under standard conditions.

3. Using Calorimetry

Calorimetry is simply the practical technique of measuring enthalpy changes by tracking temperature. The core idea across every version of this experiment: heat energy released or absorbed by a reaction changes the temperature of a known mass of liquid (usually water), and since we know water's specific heat capacity, we can work backwards to calculate exactly how much energy was transferred.

Reactions in Solution (e.g. displacement, neutralisation)

You carry out the reaction directly inside an insulated cup (polystyrene cup, vacuum flask, or metal can), stir continuously, and record temperature over time. One reagent is used in excess so the other is fully consumed — this makes the "number of moles reacted" calculation straightforward.

To actually calculate an enthalpy change per mole, we divide the total energy transferred by the number of moles of the limiting reagent:

Temperature Correction Graphs

Real reactions aren't always instant — sometimes it takes a few seconds for the maximum temperature to be reached, and during that delay heat is already leaking to the surroundings. This means the "true" maximum temperature is never actually recorded directly. The fix: plot temperature against time, then extrapolate the cooling section of the graph backwards to the exact moment the second reactant was added.

You use both extrapolated lines to read off ΔT at the moment of mixing, which gives a far more accurate (and larger) value than just reading the actual peak temperature from the raw data.

Enthalpy of Combustion

Here the logic flips slightly: instead of the reaction happening in the water, we burn a fuel underneath a container of water and measure how much the water heats up. The fuel is weighed before and after burning so we know exactly how much mass was burnt.

4. Using Hess Cycles

Some enthalpy changes simply can't be measured directly in a lab. For example, you can't easily react carbon and hydrogen together to directly form propane under standard conditions. So how do chemists find ΔH_f⦵ for propane? They use Hess's Law.

Building a Cycle Using ΔH_f (Formation Data)

When you're given enthalpy of formation data, you build a cycle where elements sit at the bottom, with arrows pointing up to both reactants and products (because formation always goes from elements upward to a compound).

1. Write the balanced equation for the reaction
2. Write the elements (with correct moles + state symbols) underneath
3. Draw upward arrows from elements to both reactants and products
4. Label each arrow with the ΔH_f value × number of moles
5. Go around the cycle: reverse the sign of any value if you travel against the arrow's direction

Building a Cycle Using ΔH_c (Combustion Data)

Combustion data is often easier to measure experimentally than formation data, so it's frequently used instead. Here the logic flips: since burning things always converts them down to combustion products (CO₂ and H₂O), the arrows point downward to the combustion products, which sit at the bottom of the cycle this time.

1. Write the equation for formation of the compound
2. Write the combustion products (usually CO₂ + H₂O) below the equation
3. Draw downward arrows from every substance to its combustion products
4. Label arrows with ΔH_c × moles
5. Travel around the cycle, flipping signs against arrow direction

5. Bond Enthalpies

Zoom right into a reaction at the molecular level, and every reaction is really just a rearrangement of atoms — old bonds break, new bonds form. Since bonds are forces of attraction, pulling atoms apart always costs energy (endothermic), while atoms coming together to form new bonds releases energy (exothermic).

Why "average"? Take methane, CH₄. You might assume all four C–H bonds are identical, but breaking the first C–H bond is actually easier than breaking the second, because the remaining hydrogens get pulled in more strongly toward the carbon once one hydrogen leaves. Since we can't practically measure each individual bond's energy, we take the total dissociation energy for the whole molecule and divide by the number of bonds — then compare that average against similar compounds to settle on an accepted value.

If more energy is released forming new bonds than was spent breaking old ones, the overall reaction is exothermic and the products are more stable than the reactants. If the opposite is true, the reaction is endothermic and the products are less stable.

Worked Example — Haber Process

N₂(g) + 3H₂(g) ⇌ 2NH₃(g), using average bond energies: N≡N = 945, H–H = 436, N–H = 391 kJ/mol

Δ_rH⦵ = (+2253) + (−2346) = −93 kJ/mol

What to Memorise

Concepts Checklist

• I can explain the difference between exothermic and endothermic in terms of energy of system vs. surroundings
• I can draw and label both an enthalpy level diagram and a reaction profile diagram
• I understand what activation energy and transition state mean, and can identify them on a curve
• I can explain why exothermic reactions generally have lower activation energy than endothermic ones
• I know the standard conditions (pressure, temperature, physical state) and what the ⦵ symbol means
• I can state the precise definitions of ΔH of reaction, formation, combustion, and neutralisation
• I know ΔH_f⦵ of an element in its standard state is always zero
• I can use q = mcΔT to calculate heat transferred in a calorimetry experiment
• I can use ΔH = q ÷ n to convert energy transferred into an enthalpy change per mole
• I understand why temperature correction graphs are used and how to extrapolate them
• I can identify the main sources of error in calorimetry (heat loss, incomplete combustion)
• I can state Hess's Law and explain it using an analogy (e.g. the mountain path)
• I can construct a Hess cycle using enthalpies of formation (elements at the bottom, arrows up)
• I can construct a Hess cycle using enthalpies of combustion (combustion products at the bottom, arrows down)
• I remember to reverse the sign of a ΔH value when travelling against an arrow in a cycle
• I understand why bond breaking is endothermic and bond making is exothermic
• I know the difference between bond dissociation enthalpy and average bond enthalpy
• I can calculate ΔH from bond enthalpies (bonds broken + bonds formed) and remember the correct signs
• I can draw full displayed structures before a bond enthalpy calculation to avoid missing bonds
• I always check for coefficients in balanced equations and multiply bond/enthalpy values accordingly

Exam Tips & Common Mistakes

What examiners are actually looking for:

• Precise definitions using the exact phrasing (e.g. "one mole", "standard conditions", "excess oxygen") — vague wording loses marks even if the general idea is right.
• Clear, methodical working in multi-step calculations — don't skip straight to an answer.
• Correct sign conventions throughout, especially in final answers.
• Recognising which named enthalpy change type an equation represents (sometimes more than one applies at once).
• Awareness of experimental limitations — heat loss, assumptions about specific heat capacity of solutions being the same as water, incomplete combustion.