This page looks briefly at enthalpy changes of neutralisation. In common with my experience with most of the other pages in this section, searches for reliable data throw up various values for the same reaction. Don't worry too much about this. It doesn't actually affect the arguments.

Enthalpy change of neutralisation

Defining standard enthalpy change of neutralisation

The standard enthalpy change of neutralisation is the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water.

Notice that enthalpy change of neutralisation is always measured per mole of water formed.

Enthalpy changes of neutralisation are always negative - heat is given out when an acid and and alkali react. For reactions involving strong acids and alkalis, the values are always very closely similar, with values between -57 and -58 kJ mol-1.

That varies slightly depending on the acid-alkali combination (and also on what source you look it up in!).

Why do strong acids reacting with strong alkalis give closely similar values?

We make the assumption that strong acids and strong alkalis are fully ionised in solution, and that the ions behave independently of each other. For example, dilute hydrochloric acid contains hydrogen ions and chloride ions in solution. Sodium hydroxide solution consists of sodium ions and hydroxide ions in solution.

The equation for any strong acid being neutralised by a strong alkali is essentially just a reaction between hydrogen ions and hydroxide ions to make water. The other ions present (sodium and chloride, for example) are just spectator ions, taking no part in the reaction.

The full equation for the reaction between hydrochloric acid and sodium hydroxide solution is:

. . . but what is actually happening is:

If the reaction is the same in each case of a strong acid and a strong alkali, it isn't surprising that the enthalpy change is similar.

Note:  Actually, of course, the enthalpy changes should be the same, not similar, if the assumptions we are making are exactly true! The small differences between strong acid-strong base combinations are almost invariably glossed over at this level. In fact, I can't remember ever seeing this discussed in any source - textbook or web. It isn't uncommon to find a list of enthalpy changes of neutralisation showing some variability in the strong acid-strong alkali cases, and then a few lines later on, this is ignored completely with a statement that in these cases, the enthalpy changes of neutralisation are the same, because . . .

I have decide not to waste time trying to sort out the exact reasons for the problem, because I suspect it will take ages and ages, and it is never going to get asked at this level anyway.

Why do weak acids or weak alkalis give different values?

In a weak acid, such as ethanoic acid, at ordinary concentrations, something like 99% of the acid isn't actually ionised. That means that the enthalpy change of neutralisation will include other enthalpy terms involved in ionising the acid as well as the reaction between the hydrogen ions and hydroxide ions.

And in a weak alkali like ammonia solution, the ammonia is also present mainly as ammonia molecules in solution. Again, there will be other enthalpy changes involved apart from the simple formation of water from hydrogen ions and hydroxide ions.

For reactions involving ethanoic acid or ammonia, the measured enthalpy change of neutralisation is a few kilojoules less exothermic than with strong acids and bases.

For example, one source which gives the enthalpy change of neutralisation of sodium hydroxide solution with HCl as -57.9 kJ mol-1, gives a value of -56.1 kJ mol-1 for sodium hydroxide solution being neutralised by ethanoic acid.

For very weak acids, like hydrogen cyanide solution, the enthalpy change of neutralisation may be much less. A different source gives the value for hydrogen cyanide solution being neutralised by potassium hydroxide solution as -11.7 kJ mol-1, for example.

Questions to test your understanding

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questions on neutralisation enthalpies


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© Jim Clark 2010 (modified July 2013)