The pH Scale, Neutralisation and Strong vs Weak Acids

What makes a solution acidic or alkaline comes down to which ions it releases in water. An acid produces hydrogen ions, $H^{+}$, in aqueous solution. An alkali is a soluble base that produces hydroxide ions, $O{H}^{-}$, in aqueous solution.

For example, hydrochloric acid releases hydrogen ions, and sodium hydroxide releases hydroxide ions:

$HCl\to H^{+}+C{l}^{-}$

$NaOH\to N{a}^{+}+O{H}^{-}$

The more hydrogen ions there are, the more acidic the solution. The more hydroxide ions there are, the more alkaline it is.

An acid produces H⁺ ions in water; an alkali produces OH⁻ ions. These two ions are the reason acids and alkalis behave as they do.

Understanding that acidity is caused by $H^{+}$ ions is the foundation for the pH scale and for neutralisation.

The pH scale measures how acidic or alkaline a solution is. It runs from 0 to 14:

pH range	Meaning	Example
0–6	Acidic	Lemon juice, stomach acid
7	Neutral	Pure water
8–14	Alkaline	Soap, oven cleaner

The lower the pH, the more acidic the solution (more $H^{+}$ ions). The higher the pH, the more alkaline the solution (more $O{H}^{-}$ ions). A pH of exactly 7 is neutral.

pH can be measured in two ways:

• Universal indicator (a wide-range indicator): added to the solution, it turns a colour matched against a chart to give an approximate pH — red for strongly acidic, green for neutral, purple for strongly alkaline.
• pH probe / pH meter: an electronic sensor that gives a precise numerical pH reading.

Universal indicator gives an approximate pH from a colour; a pH probe gives an accurate numerical value. Use "approximate" when describing the indicator method.

When an acid reacts with an alkali, the hydrogen ions from the acid react with the hydroxide ions from the alkali to form water. This is why the mixture becomes neutral: the ions that caused acidity and alkalinity are used up. The ionic equation for all acid–alkali neutralisations is:

$H^{+}+O{H}^{-}\to H_{2}O$

This single equation lies underneath every acid–alkali reaction, whatever the salt formed. For example, when hydrochloric acid neutralises sodium hydroxide the full equation is $HCl+NaOH\to NaCl+H_{2}O$, but the sodium and chloride ions are spectators, and the actual reaction is simply $H^{+}$ joining $O{H}^{-}$ to make water.

Learn $H^{+}+O{H}^{-}\to H_{2}O$ exactly. It is the ionic equation for neutralisation between any acid and any alkali, and it is a common exam recall mark.

As the reaction proceeds, the pH moves towards 7. When just enough alkali has been added to react with all the acid, the solution is neutral.

(Higher Tier only) Strong and weak acids are assessed at Higher Tier.

Acids differ in how far they ionise (split into ions) in water. This is what "strong" and "weak" mean, and it is a different idea from "concentrated" and "dilute".

• A strong acid is fully ionised in aqueous solution: essentially every molecule splits up to release its $H^{+}$ ions. Hydrochloric, nitric and sulfuric acids are strong.
• A weak acid is only partially ionised: at any moment only a small fraction of the molecules have released their $H^{+}$ ions. Ethanoic, citric and carbonic acids are weak.

$\underset{\text{strong: fully ionised}}{\underbrace{HCl\to H^{+}+C{l}^{-}}}\underset{\text{weak: partially ionised}}{\underbrace{C{H}_{3}COOH\rightleftharpoons C{H}_{3}CO{O}^{-}+H^{+}}}$

The reversible arrow ($\rightleftharpoons$) for the weak acid shows that its ionisation does not go to completion. Because a strong acid releases more $H^{+}$ ions, at the same concentration a strong acid has a lower pH than a weak acid.

(Higher Tier) Strength = degree of ionisation. Concentration = amount of acid dissolved per unit volume. A dilute strong acid and a concentrated weak acid are entirely possible.

(Higher Tier only) The relationship between pH and hydrogen ion concentration is assessed at Higher Tier.

The pH scale is linked to the hydrogen ion concentration. As the pH decreases by 1 unit, the hydrogen ion concentration increases by a factor of 10. So each step down the pH scale means ten times more $H^{+}$ ions.

Worked example. A solution at pH 3 is diluted so its pH rises to pH 5. How has the hydrogen ion concentration changed?

1. The pH has increased by 2 units (from 3 to 5).
2. Each 1-unit increase means the $H^{+}$ concentration falls by a factor of 10.
3. Two units means a factor of $10\times 10=100$.

So the hydrogen ion concentration is 100 times smaller at pH 5 than at pH 3.

pH change	Factor change in H⁺ concentration
Down 1 (e.g. 4 → 3)	×10 more H⁺
Down 2 (e.g. 5 → 3)	×100 more H⁺
Up 1 (e.g. 3 → 4)	×10 fewer H⁺
Up 2 (e.g. 3 → 5)	×100 fewer H⁺

(Higher Tier) Lower pH means more hydrogen ions. A drop of one whole pH unit is a tenfold rise in H⁺ concentration.

1. Confusing strong with concentrated (HT)

Strength is about how far an acid ionises; concentration is about how much acid is dissolved. A concentrated weak acid still ionises only partially. Keep the two ideas separate.

2. Getting the pH direction backwards

Low pH is acidic (more H⁺), high pH is alkaline (more OH⁻). pH 2 is more acidic than pH 5, because more hydrogen ions means a lower pH.

3. Writing the wrong ions for acids and alkalis

Acids produce H⁺ ions; alkalis produce OH⁻ ions. Do not swap them. Neutralisation joins these two to make water.

4. Saying universal indicator gives an exact pH

Universal indicator gives only an approximate pH from its colour. For an accurate numerical value you need a pH probe or meter.

5. Forgetting the factor of 10 (HT)

Each 1-unit change in pH is a tenfold change in hydrogen ion concentration, not a change of 1. From pH 4 to pH 2 is two units, so ×100 more H⁺.

6. Missing the neutralisation ionic equation

Many acid–alkali questions ask for the ionic equation. It is always $H^{+}+O{H}^{-}\to H_{2}O$, regardless of which acid and alkali react.