Question
What does a negative value of ΔH indicate? Distinguish between exothermic and endothermic reactions with examples.
Solution — Step by Step
ΔH is the enthalpy change of a reaction — the heat exchanged between the system (the reacting chemicals) and the surroundings at constant pressure. The sign tells you the direction of heat flow, not the amount of heat inside the molecule.
When ΔH is negative, the products have lower enthalpy than the reactants. That “extra” energy has to go somewhere — it flows out as heat into the surroundings. The surroundings get warmer. This is an exothermic reaction.
When ΔH is positive, the products sit at higher enthalpy than the reactants. The reaction pulls heat from the surroundings to make this happen. The surroundings get cooler. This is an endothermic reaction.
Exothermic examples (ΔH is negative):
- Combustion of methane: ,
- Neutralisation of strong acid + strong base: (standard value NCERT expects)
Endothermic examples (ΔH is positive):
- Melting of ice:
- Decomposition of calcium carbonate: ,
| ΔH sign | Heat flow | Reaction type | Surroundings feel |
|---|---|---|---|
| Negative (−) | Out of system | Exothermic | Warmer |
| Positive (+) | Into system | Endothermic | Cooler |
Negative ΔH = exothermic. Positive ΔH = endothermic.
Why This Works
Enthalpy is essentially a measure of the chemical energy stored in bonds. When bonds break, energy is absorbed; when bonds form, energy is released. In an exothermic reaction, the bonds formed in products are stronger (release more energy) than the bonds broken in reactants. The net result is energy leaving the system.
This is why combustion reactions always have large negative ΔH values — burning carbon-hydrogen bonds and replacing them with the much stronger C=O and O-H bonds in CO₂ and H₂O releases enormous energy.
The constant-pressure condition matters because most reactions in chemistry labs and in our bodies happen open to the atmosphere. At constant pressure, the heat exchanged equals ΔH exactly — no correction needed. This is why NCERT Chapter 6 introduces enthalpy specifically for constant-pressure processes.
Alternative Method
Instead of thinking about enthalpy levels, use bond energy to predict the sign of ΔH:
If more energy is released in forming bonds than absorbed in breaking them, the result is negative — exothermic. This method gives approximate values (actual ΔH uses standard enthalpies of formation), but it tells you the sign quickly in MCQ situations.
For neutralisation questions in NEET and CBSE boards: the standard enthalpy of neutralisation of a strong acid with a strong base is always −57.1 kJ/mol. This is because the net reaction is always regardless of which acid or base you use. Memorise this value.
Common Mistake
Students confuse “the system releases heat” with “the system absorbs heat” because the sign convention trips them up. A negative ΔH means the system loses energy — it releases heat to the surroundings. Many students write: “negative ΔH means the surroundings lose heat.” That is backwards. The surroundings gain heat when ΔH is negative. Always ask: which side is the system, which side are the surroundings?
A related error in board exams: students describe endothermic reactions as “reactions that don’t release energy.” That is wrong. Endothermic reactions can still release some energy — they just absorb more than they release. The net absorption makes ΔH positive.