Two words, one minus sign, and a surprising number of lost marks. Almost every mistake here comes from forgetting that the thermometer is sitting in the surroundings, while ΔH describes the system. Get that straight and the signs stop being guesswork.
📘 What you need to know
Enthalpy (H) is the chemical energy stored in a substance, measured at constant pressure.
ΔH = H(products) − H(reactants). You can never measure H itself, only the change.
Exothermic: energy leaves the system. Products are lower in enthalpy, ΔH is negative, the surroundings get hotter.
Endothermic: energy enters the system. Products are higher in enthalpy, ΔH is positive, the surroundings get colder.
Breaking bonds always takes energy in; making bonds always gives energy out. The sign of ΔH is just which of the two won.
Reverse a reaction and ΔH keeps its size but flips its sign.
A negative ΔH makes a reaction energetically favourable — but it does not promise the reaction will actually be fast.
Enthalpy: the energy stored in the chemicals
Bonds are stores of chemical energy. Add up everything stored in a substance under normal laboratory conditions (constant pressure, open to the atmosphere) and you have its enthalpy, symbol H.
Here is the awkward truth: nobody can measure H. There is no absolute zero point for enthalpy that we can start counting from. What we can do is measure how much it changes during a reaction, and that is all chemistry ever needs.
The definition everything rests on
ΔH = H(products) − H(reactants)
Read the order of that subtraction carefully — products first. If the products ended up with less stored energy than the reactants, the subtraction gives a negative answer, and the energy that is no longer in the chemicals has gone out into the surroundings.
Think of enthalpy like money in a bank account. You cannot see the balance, but you can see every deposit and withdrawal. ΔH is the withdrawal slip: negative means the system paid out.
Exothermic reactions
In an exothermic reaction, energy flows out of the chemicals and into everything around them. Trace what happens step by step:
The products hold less enthalpy than the reactants, so ΔH is negative.
That released energy has to go somewhere, so it warms the solvent, the glassware and the air.
The temperature of the surroundings rises — which is what your thermometer records.
The system itself has lost enthalpy, even though the mixture feels hot.
Combustion, neutralisation, respiration, most metal displacement reactions and the setting of cement are all exothermic. If a reaction makes something get warm, that is your clue.
Endothermic reactions
Now run the argument backwards. Energy is pulled in from the surroundings and locked into the new bonds:
The products hold more enthalpy than the reactants, so ΔH is positive.
The energy has to come from somewhere, and it is taken from the solvent and the container.
The temperature of the surroundings falls — the beaker feels cold, sometimes cold enough to frost on the outside.
The system has gained enthalpy.
Thermal decomposition (heating a metal carbonate), photosynthesis, dissolving ammonium nitrate and the reaction in an instant cold pack are all endothermic.
Notice that the vertical axis is enthalpy, not temperature, and the horizontal direction is simply the progress of the reaction. Only the two levels matter.
Why is any reaction exothermic in the first place?
This is the part most notes skip, and it is the part that makes the sign obvious rather than memorised. Every reaction does two jobs:
Breaking the bonds in the reactants. You must supply energy to pull bonded atoms apart, so this step is always endothermic.
Making the bonds in the products. Energy is given out when atoms snap together, so this step is always exothermic.
The overall ΔH is simply the balance of the two. If making the new bonds releases more than breaking the old ones cost, there is energy left over and it escapes as heat. That is an exothermic reaction. Here is the balance sheet for hydrogen burning in chlorine, using bond enthalpies from the data booklet:
You will meet this calculation properly under bond enthalpies. For now it is here to show why a reaction comes out exothermic, rather than asking you to remember that it does.
Spotting it in the lab and in the question
What you observe
Type
Sign of ΔH
Everyday examples
Thermometer reading rises
Exothermic
Negative
Burning fuels, neutralising acid with alkali, adding magnesium to acid
Careful with “gives out heat”. A reaction that is being heated by a Bunsen burner is not automatically endothermic. Heating just gets the reaction started or speeds it up. Ask instead: once it is going, does the mixture warm the surroundings or cool them?
Reverse the reaction and the sign flips
Enthalpy changes are reversible bookkeeping. If a forward reaction releases 184 kJ mol−1, then pushing it backwards must cost the same 184 kJ mol−1. Same magnitude, opposite sign. This is worth internalising now because it is the whole engine behind Hess’s Law later on.
Reversing a reaction
ΔH(reverse) = −ΔH(forward)
Energetically favourable is not the same as fast
An exothermic reaction is going downhill in energy, which is why chemists call it energetically favourable. That does not mean it will happen when you want it to. A lump of coal sitting in air has a hugely negative ΔH available to it and does nothing at all until you light it.
The missing ingredient is rate. A reaction can be energetically favourable and still be so slow it looks dead; we say it is kinetically controlled. Enthalpy tells you where the reaction wants to end up. Kinetics tells you how long you will be waiting.
If a question asks why a favourable reaction does not appear to happen, the answer they want is almost always about the activation energy, which is exactly what the next page is about.
WORKED EXAMPLE
50 cm3 of sodium hydroxide solution is placed in a polystyrene cup at 21.3 °C. Hydrochloric acid is added and the temperature rises to 28.6 °C. State whether the reaction is exothermic or endothermic, give the sign of ΔH, and say what has happened to the enthalpy of the system.
Step 1: What did the thermometer do?ΔT = 28.6 − 21.3 = +7.3 °Cthe thermometer is in the solution, which is the surroundingsStep 2: Surroundings got hotter, so energy left the systemExothermic, ΔH is negativeStep 3: What happened to the system?the system lost enthalpy: the products store less energy than the reactants did
WORKED EXAMPLE
Ammonium nitrate is stirred into water and the temperature drops from 20.8 °C to 15.2 °C. Sketch the shape of the enthalpy level diagram in words and state the sign of ΔH.
Step 1: Read the temperature changeΔT = 15.2 − 20.8 = −5.6 °CStep 2: Surroundings cooled, so they gave energy awaythat energy went into the system, so the system gained enthalpyStep 3: Diagramreactants level low on the left, products level higher on the right, arrow pointing upEndothermic, ΔH is positive
WORKED EXAMPLE
The complete combustion of methane has ΔH = −890 kJ mol−1. Calculate the enthalpy change when (a) 0.500 mol of methane burns, and (b) 1 mol of carbon dioxide and 2 mol of water react to re-form methane and oxygen.
(a) Scale with the amountΔH = 0.500 × (−890) = −445 kJhalf the methane releases half the energy; note the unit is kJ, not kJ mol⁻¹(b) This is the reverse reactionΔH = −(−890) = +890 kJ mol⁻¹(a) −445 kJ (b) +890 kJ mol⁻¹same size, opposite sign, so the reverse reaction is endothermic
🧩 Getting the sign right, every time
Find the temperature change of the mixture. Up or down?
Remember the mixture is the surroundings, not the system.
Warmer surroundings means energy came out of the system, so ΔH is negative. Colder surroundings means the opposite.
Sanity-check against the enthalpy level diagram: does your arrow point down (negative) or up (positive)?
Only now write the number, and attach the sign before you attach the units.
💡 Exam tip
Always write the sign of ΔH, even when it is positive. An unsigned answer is treated as ambiguous.
On enthalpy level diagrams, label both levels, label the axis enthalpy (not “energy released”), and draw the arrow from reactants to products.
Say which temperature changed. “The temperature of the surroundings increases” is the phrase examiners are looking for.
If you are asked for ΔH for a given amount, the unit is kJ. If it is per mole of reaction, it is kJ mol−1.
“Exothermic” and “energetically favourable” earn marks; “the reaction wants to happen” does not.
If a question gives you both bond breaking and bond making energies, subtract in the order in − out and the sign appears by itself.
⚠ Common mix-up
“Exothermic means the reaction gets hot, so ΔH is positive.” The mixture gets hot precisely because the system lost energy. ΔH is negative.
Saying the system gets hotter in an exothermic reaction. The system loses enthalpy; it is the surroundings that heat up.
Thinking bond breaking releases energy. Breaking always takes energy in. Only bond making gives it out.
Drawing the enthalpy arrow upwards for an exothermic reaction. The arrow follows the reaction, and the reaction goes downhill.
Confusing endothermic with “needs heating”. Plenty of exothermic reactions need a match to start them.
Forgetting to flip the sign when a question quietly reverses the equation.
Assuming exothermic means fast. Energetics and rate are separate questions.
Up next: Energy Profile Diagrams — the same two levels, but this time we draw the hill in between them and find out why favourable reactions can still refuse to start.
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