The pH SCALE measures the CONCENTRATION of hydrogen ions (H⁺) in a solution — it indicates how acidic or alkaline a solution is.
Scale range: 0 to 14 (though values slightly outside this range are possible).
pH 0–6: ACIDIC (more H⁺ ions than OH⁻ ions)
pH 7: NEUTRAL (equal H⁺ and OH⁻ — pure water at 25°C)
pH 8–14: ALKALINE (more OH⁻ ions than H⁺ ions)
The pH scale is LOGARITHMIC — each unit change represents a 10× change in H⁺ concentration:
pH 3 has 10× more H⁺ than pH 4.
pH 3 has 100× more H⁺ than pH 5.
Typical pH values:
Hydrochloric acid (conc.): pH ~0–1
Vinegar (ethanoic acid): pH ~3
Coffee: pH ~5
Pure water: pH 7
Baking soda solution: pH ~9
Sodium hydroxide solution: pH ~13–14
Indicators and Measuring pH
An INDICATOR is a substance that changes colour depending on the pH of a solution.
UNIVERSAL INDICATOR:
A mixture of indicators that shows a range of colours across the pH scale.
Red → orange → yellow → green → blue → purple as pH increases from 0 to 14.
Gives an APPROXIMATE pH — tells you the pH range.
LITMUS:
Red in ACID, blue in ALKALI, purple in neutral.
Simple test — tells you acid or alkali, not the pH number.
PHENOLPHTHALEIN:
Colourless in acid, PINK in alkali.
Used in titrations — clear colour change at the end point.
pH PROBE / METER:
Gives a PRECISE numerical pH reading — more accurate than indicators.
Used in industry and for accurate laboratory measurements.
THE RELATIONSHIP BETWEEN pH AND CONCENTRATION:
The MORE H⁺ ions in solution → LOWER pH (more acidic).
The MORE OH⁻ ions in solution → HIGHER pH (more alkaline).
Adding water DILUTES the solution → concentration of H⁺ decreases → pH INCREASES towards 7.
Effect of Neutralisation on pH
Adding a BASE to an ACID gradually increases the pH (makes it less acidic).
Adding an ACID to an ALKALI gradually decreases the pH (makes it less alkaline).
A TITRATION CURVE shows how pH changes as acid is added to alkali:
At the START: pH high (~13) — the solution is strongly alkaline.
As acid is added: pH falls slowly at first.
Near the END POINT: pH drops RAPIDLY — a large pH change for a small addition.
At the END POINT: pH = 7 (if strong acid + strong alkali).
After the end point: pH continues to fall as excess acid is added.
WHY SUCH A SHARP CHANGE NEAR THE END POINT:
Near the equivalence point, there is very little OH⁻ left to absorb the H⁺ from the added acid.
So each small addition of acid produces a large change in H⁺ concentration → large pH change.
⚠️ Common Mistake
Lower pH = MORE ACIDIC = MORE H⁺ ions. Higher pH = MORE ALKALINE = FEWER H⁺ ions (more OH⁻). Students often confuse pH 2 and pH 12 — pH 2 is strongly acidic, pH 12 is strongly alkaline. pH 7 is neutral — pure water. NOT all solutions at pH 7 are water.
📌 Key Note
pH 0–6: acidic. pH 7: neutral. pH 8–14: alkaline. Lower pH = more H⁺ ions. pH scale is logarithmic — each unit = 10× change in H⁺. Universal indicator: shows range of colours. Phenolphthalein: colourless in acid, pink in alkali. pH meter: most accurate.
🎯 Matching Activity — Match the pH to the Description
Match each pH value to the correct description. — drag the symbols on the right to match the component names on the left.
pH 1
Drop here
pH 7
Drop here
pH 13
Drop here
pH 4
Drop here
pH 10
Drop here
Strongly acidic — high concentration of H⁺ ions
Weakly acidic — e.g. black coffee or tomato juice
Weakly alkaline — e.g. baking soda solution
Strongly alkaline — high concentration of OH⁻ ions
Neutral — equal H⁺ and OH⁻ ions — pure water at 25°C
⭐ Higher Tier Only
Strong acids fully dissociate (HCl → H⁺ + Cl⁻) — maximum H⁺ at given concentration. Weak acids partially dissociate (CH₃COOH ⇌ CH₃COO⁻ + H⁺) — less H⁺ at same concentration → higher pH. Strong acid and weak acid at same concentration: strong acid has lower pH. Strong ≠ concentrated.
🎯 Test Yourself
Question 1 of 2
1. A solution has a pH of 2. Another has a pH of 4. How do their H⁺ ion concentrations compare?
2. An acid is diluted by adding water. What happens to its pH?
⭐ How Well Do You Understand This Topic?
Be honest with yourself — this helps you know what to revise!
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