Neutralisation: An Introduction

Introduction

If you add lemon juice to milk of magnesia, the sourness of one and the chalky alkalinity of the other both seem to disappear. Acids and bases are chemical opposites, and when they meet they react so that each cancels out the other's properties. This reaction is called neutralisation . It is one of the most important reactions in chemistry, used everywhere from indigestion remedies to farming and the treatment of factory waste. This page introduces the idea and the particle picture behind it.

Core explanation

What happens. When an acid reacts with a base, the acidic properties of the acid and the basic properties of the base are both reduced. If exactly the right amounts react, the mixture is neither acidic nor alkaline. This reaction is called neutralisation , and a base can be thought of as a "neutraliser" of acids.

Following the pH. Imagine starting with a beaker of dilute hydrochloric acid at about pH 1 and slowly adding dilute sodium hydroxide solution:

- At first, the pH rises only a little, because plenty of acid is still left. - As more alkali is added, the pH climbs faster. - When just enough alkali has been added to react with all the acid, the pH reaches 7 — the neutral point . - If more alkali is added beyond that, the alkali is now in excess and the pH rises above 7, often up to 12 or 13.

The reverse is also true: adding acid to an alkali lowers the pH towards 7 and then below it.

The particle picture. Acids in water contain hydrogen ions, H⁺; alkalis contain hydroxide ions, OH⁻. When they meet, these ions join together to form water molecules:

H⁺(aq) + OH⁻(aq) → H₂O(l)

As the H⁺ ions are used up, the solution becomes less acidic. When the number of OH⁻ ions added equals the number of H⁺ ions present, both have been removed and only neutral water remains, along with the other ions from the acid and alkali. Those left-over ions form a salt , which you will study next.

Not only alkalis. Insoluble bases, such as copper(II) oxide or magnesium oxide, and carbonates, such as calcium carbonate, also neutralise acids. Carbonates produce carbon dioxide gas as well, so you see fizzing.

Neutralisation is a chemical change. New substances are formed — water and a salt — and there is an energy change: neutralisation releases heat, so the mixture warms up slightly.

Step-by-step reasoning

To predict what happens when an alkali is added to an acid:

1. Start with the acid's pH, which is below 7. 2. Each portion of alkali removes some H⁺ ions by forming water. 3. The pH rises as the H⁺ ions are used up. 4. At the neutral point, H⁺ and OH⁻ have exactly cancelled, and the pH is 7. 5. Any further alkali is in excess and pushes the pH above 7.

Visual explanation

Picture a see-saw. On one side sit red counters labelled H⁺; on the other, blue counters labelled OH⁻. Each time a red and a blue counter meet, they pair up and leave as a water molecule. When the numbers are equal, all pairs have left and the see-saw is level — the neutral point.

Real-world analogy

Neutralisation is like mixing hot and cold water in a bath. Hot water on its own is too hot; cold water is too cold. Mix them in the right proportions and you get a comfortable middle temperature. Add too much of one and the bath tips towards that extreme again.

Real-world example

Indigestion is caused by too much stomach acid. Antacid tablets contain bases such as magnesium hydroxide or calcium carbonate. When swallowed, they neutralise some of the excess hydrochloric acid in the stomach, raising the pH slightly and relieving the burning feeling.

Why?

Why does the pH rise slowly at first and then quickly near the neutral point? The pH scale is based on factors of ten. At the start, removing some H⁺ hardly changes the huge number of ions present. Near neutral, only a tiny amount of H⁺ is left, so removing the last of it makes a large difference to the pH.

Common misconception

"Neutralisation always produces a solution of pH 7." Only exactly matched amounts of an acid and alkali of similar strength give pH 7. If either is in excess, or if a strong acid reacts with a weak base, the final pH may be above or below 7.

Worked example

Question: A student adds dilute sodium hydroxide to dilute hydrochloric acid until the pH is 11. Describe what has happened to the H⁺ ions and which reactant is in excess.

Reasoning: The acid started below pH 7. The alkali's OH⁻ ions reacted with all the H⁺ ions to form water. Because the pH is now above 7, more OH⁻ was added than needed.

Answer: All the H⁺ ions have been removed as water; sodium hydroxide is in excess.

Quick check

1. Which ions combine during the neutralisation of an acid by an alkali, and what do they form? Answer: Hydrogen ions and hydroxide ions combine to form water.

Exam focus

Be able to define neutralisation as the reaction between an acid and a base, write the ionic equation H⁺ + OH⁻ → H₂O, and describe the pH change when an alkali is added to an acid: rising towards 7 and then above 7 if alkali is in excess.

Advanced insight

A graph of pH against the volume of alkali added is called a titration curve. For a strong acid and a strong alkali, the curve is flat at first and then rises very steeply around pH 7, jumping several pH units for a single drop of alkali. This steep region is why indicators can pinpoint the neutral point so precisely.

Summary

Neutralisation is the reaction between an acid and a base, in which each cancels the other's properties. As an alkali is added to an acid, the pH rises to 7 at the neutral point and above 7 if alkali is in excess. The key particle reaction is H⁺ + OH⁻ → H₂O. Insoluble bases and carbonates also neutralise acids. Neutralisation is a chemical change that releases heat.

Practice questions

1. Define neutralisation. Answer: Neutralisation is the reaction between an acid and a base in which they cancel out each other's properties. 2. Write the ionic equation for neutralisation between an acid and an alkali. Answer: H⁺(aq) + OH⁻(aq) → H₂O(l). 3. Describe how the pH changes when dilute hydrochloric acid is added slowly to sodium hydroxide solution until the acid is in excess. Answer: The pH starts high, around 13, falls towards 7 at the neutral point and then drops below 7 once acid is in excess. 4. Give one everyday example of neutralisation and name the base involved. Answer: Taking an antacid for indigestion; the base is magnesium hydroxide or calcium carbonate.