Why does IB Chemistry define enthalpy change at constant pressure?

Most reactions you do at school happen in open containers, so the pressure is approximately constant and equal to atmospheric pressure. Under these conditions, the heat you measure corresponds neatly to ΔH, which is why IB Chemistry uses enthalpy change as the standard “heat of reaction” quantity. If pressure changed significantly, some energy would go into expansion work differently, and your heat measurement wouldn’t match the same simple definition. In exams, “constant pressure” is also a clue about what symbol you should use: ΔH rather than other thermodynamic quantities. It’s not there to intimidate you; it’s there to keep the definition consistent. When in doubt, state the definition clearly and link it to typical lab conditions.

Is enthalpy the same thing as energy?

Not exactly, and IB Chemistry likes that distinction. Enthalpy (H) is related to internal energy, but it also accounts for energy associated with a system pushing back the atmosphere (pressure-volume work). That’s why enthalpy change is so convenient for reactions at constant pressure: it matches the heat transfer you can actually measure in a simple calorimetry setup. In many exam questions, you don’t need to derive any of this; you just need to use the definition and apply the correct equations. Still, knowing “enthalpy is a bookkeeping tool for heat at constant pressure” helps your explanations feel grounded. It also helps you avoid treating ΔH as a random formula symbol.

Why can a reaction feel cold even if “energy is released” somewhere?

First, check what you’re actually observing: your hand feels the temperature of the surroundings, not the energy changes inside chemical bonds directly. In IB Chemistry, an endothermic reaction absorbs heat from the surroundings, so it can genuinely feel cold. But even with exothermic reactions, heat might be released slowly, or it might spread out quickly into the air and the container, so the temperature rise is small. Calorimetry also has limitations: heat loss to the surroundings can hide part of the temperature change, making ΔH look less negative than it should be. Another common reason is that multiple processes happen at once, and the net effect determines what you feel. In exam responses, naming at least two realistic factors (rate, heat loss, measurement sensitivity) usually earns strong evaluation marks.

Enthalpy Change Explained for IB Chemistry

IB Chemistry

RevisionDojo Team

•11/21/2025•7 min read

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Enthalpy Change Explained for IB Chemistry

IB Chemistry

RevisionDojo Team

•11/21/2025•7 min read

Enthalpy change is the kind of idea that feels simple right up until the first time you miss a negative sign and your “exothermic” reaction suddenly needs a heater. In IB Chemistry, that tiny sign flip can turn a solid explanation into a lost mark.

The good news: enthalpy change is not mysterious. It’s just a careful way of describing heat transfer during reactions (under constant pressure), so you can calculate, compare, and explain reaction energetics with confidence in IB Chemistry exams and labs.

Enthalpy change in IB Chemistry (in one clean definition)

Enthalpy change (ΔH) is the heat energy absorbed or released by a system during a chemical reaction at constant pressure.

That “constant pressure” detail matters because most school experiments happen in open containers, so pressure is approximately atmospheric. In IB Chemistry, this is why ΔH is the default energy change you’ll meet in thermochemistry: it matches how reactions are commonly measured.

If you want the syllabus-aligned version with practice built in, start with R1.1 Measuring enthalpy change.

Quick checklist: what you must remember about ΔH

Units: usually kJ mol⁻¹ (watch J vs kJ)
Sign: negative for heat released, positive for heat absorbed
Measured: often using calorimetry and temperature change
Used in: energy profile diagrams, Hess’s law, bond enthalpy estimates

RevisionDojo students often drill this with the R1.1 Questionbank and then tighten their explanations using Study Notes and Flashcards.

Exothermic vs endothermic: the story your diagram is telling

In IB Chemistry, examiners love questions where you must connect a diagram, a temperature change, and the sign of ΔH.

Exothermic reactions (ΔH < 0)

Exothermic reactions release heat to the surroundings.

What that means in practice:

Surroundings warm up
Products have lower enthalpy than reactants
Energy profile ends lower than it starts

Combustion is the classic example, and you can revise it alongside enthalpy ideas here: Enthalpy of Combustion Explained.

Endothermic reactions (ΔH > 0)

Endothermic reactions absorb heat from the surroundings.

What to look for:

Surroundings cool down
Products have higher enthalpy than reactants
Energy profile ends higher than it starts

Energy profile hiking joke

Standard enthalpy change (ΔH°): why the little circle matters

Many data-booklet values in IB Chemistry come as standard enthalpy changes, written ΔH°.

Standard conditions (IB style) are:

298 K
100 kPa
Solutions at 1.0 mol dm⁻³

Standard conditions let you compare experiments fairly, like comparing race times on the same track. If you’re using formation data, this link is the natural next step: Standard Enthalpy of Formation Explained.

Calorimetry calculations: turning temperature into ΔH

A lot of IB Chemistry enthalpy questions are really two mini-problems: calculate heat transfer, then convert to enthalpy change per mole.

Step 1: calculate heat transfer

Use:

q = mcΔT

Where:

q = heat energy (J)
m = mass of solution/water (g)
c = specific heat capacity (usually 4.18 J g⁻¹ K⁻¹ for water)
ΔT = temperature change (K or °C)

Step 2: convert to molar enthalpy change

Then use:

ΔH = q / n

Where n is moles of the limiting reagent (or moles of product specified, like water formed in neutralisation).

To improve your experimental evaluation, read R1.1.5 Practical considerations. It’s the difference between “heat loss occurred” and a high-mark answer that explains how it affects ΔH.

Why enthalpy change keeps coming back in IB Chemistry

ΔH is not a one-off definition. It’s a tool you keep reusing:

Hess’s law and energy cycles (pathway doesn’t matter)
Bond enthalpy estimates (breaking vs forming bonds)
Explaining why some reactions need continuous heating

For Hess’s law practice, use Hess’s Law Explained Simply and then test yourself with the Hess’s Law Questionbank.

For bond energetics links, pair this topic with Bond Enthalpy Explained for IB Chemistry or the syllabus notes on R1.2.1 Bond enthalpy.

Common IB Chemistry mistakes (and how to avoid them)

Mixing up ΔH and activation energy: activation energy is the “hill,” ΔH is the “start-to-finish change.”
Forgetting the per mole step: q is total heat, ΔH is usually per mole.
Sign confusion: temperature rise in surroundings usually means ΔH < 0.
Unit slips: data in kJ, calculation in J, final answer in confusion.

If you want a structured path across the whole topic, the IB Chemistry resources hub pulls together Study Notes, Flashcards, Lessons, and the Questionbank in one place.

Bring it home: make ΔH an easy mark in IB Chemistry

Enthalpy change is one of those IB Chemistry topics that rewards calm thinking: define it clearly, track the sign carefully, and treat units like they matter (because they do). Once ΔH is stable in your mind, calorimetry, energy profiles, and Hess’s law stop feeling like separate chapters and start feeling like the same story told with different tools.

To lock it in, use RevisionDojo’s Study Notes and Flashcards for the definitions, then move to the Questionbank for exam-style repetition, and finish with AI Chat to diagnose exactly why your sign or units went wrong. That workflow turns IB Chemistry thermochemistry into something you can actually trust under time pressure.