Edexcel IAL Chemistry · Energetics
Energetics
Every chemical reaction is really just a story about breaking old bonds and making new ones — and whether that story releases energy to the surroundings (exothermic) or steals energy from them (endothermic) determines everything else in this chapter.
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
ENERGY
▲
│ REACTANTS ────┐
│ │ ↓ energy released (−ΔH)
│ └──── PRODUCTS
└─────────────────────────────► extent of reaction
Why "thermodynamically possible" isn't the whole story
Exothermic reactions are energetically favourable (reactants are higher in energy, so rolling "downhill" to products makes sense). But if the reaction is too
slow, it might not visibly happen at room temperature — it's
kinetically controlled. Diamond turning into graphite is a classic example: it's exothermic and thermodynamically favourable, but so slow it never visibly happens.
Endothermic Reactions (ΔH is positive)
more
- Energy of the system increases
- Temperature of the surroundings decreases
- ΔH is written as a positive number
2. Enthalpy Change — Definitions
standard conditions
Standard Conditions
Physical states matter — a lot