The pH Scale and Indicators

This lesson covers the pH scale, how it relates to acidity and alkalinity, the use of indicators, and the important distinction between strong and weak acids (Higher tier). This content is part of the Edexcel GCSE Chemistry specification (1CH0).

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The pH Scale

The pH scale is a numerical scale from 0 to 14 that measures how acidic or alkaline a solution is. The scale is based on the concentration of hydrogen ions (H⁺) in solution.

pH Range	Classification	Examples
0–2	Strong acid	Battery acid (pH 0), stomach acid (pH 1–2)
3–6	Weak acid	Vinegar (pH 3), orange juice (pH 4), black coffee (pH 5), milk (pH 6)
7	Neutral	Pure water, sodium chloride solution
8–11	Weak alkali	Baking soda solution (pH 8–9), milk of magnesia (pH 10), ammonia solution (pH 11)
12–14	Strong alkali	Limewater (pH 12), sodium hydroxide solution (pH 13–14)

Key Rules

• pH < 7 → the solution is acidic
• pH = 7 → the solution is neutral
• pH > 7 → the solution is alkaline
• The lower the pH, the more acidic the solution (higher concentration of H⁺ ions)
• The higher the pH, the more alkaline the solution (higher concentration of OH⁻ ions)

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Universal Indicator Colour Chart

Universal indicator produces a continuous range of colours depending on the pH of the solution:

pH	0	1	2	3	4	5	6	7	8	9	10	11	12	13	14
Colour	Dark red	Red	Red-orange	Orange	Yellow-orange	Yellow	Yellow-green	Green	Blue-green	Blue	Blue	Dark blue	Indigo	Violet	Dark purple

Exam Tip: Learn the universal indicator colour chart. A very common question gives you a colour and asks you to state the approximate pH. Remember: red/orange = acidic, green = neutral, blue/purple = alkaline.

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Measuring pH: Indicator Paper vs pH Meter

There are two main methods for measuring pH in the laboratory:

Method 1: Universal Indicator Paper

1. Dip a glass rod into the solution.
2. Touch the wet rod onto a piece of universal indicator paper.
3. Compare the colour of the paper to a standard colour chart.
4. Read off the approximate pH value.

Advantages: Simple, quick, cheap, no calibration required.

Limitations: Only gives a whole-number estimate of pH; relies on comparing colour by eye, which is subjective; some coloured solutions can mask the indicator colour.

Method 2: pH Meter (pH Probe)

1. Calibrate the pH meter using buffer solutions of known pH.
2. Rinse the probe with distilled water and dry gently.
3. Place the probe into the solution.
4. Read the pH value from the digital display.

Advantages: Gives an accurate numerical value (often to 1–2 decimal places); objective (not affected by human colour perception); works with coloured or turbid solutions.

Limitations: More expensive; requires calibration; the probe is fragile and must be maintained.

Exam Tip: If asked which method is better for measuring pH, the answer is usually the pH meter, because it is more accurate and precise. Use these exact words in your answer — they are key terms in the mark scheme.

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Concentration and pH

The concentration of an acid refers to the number of moles of acid dissolved in a given volume of solution (measured in mol/dm³).

• A concentrated acid has a large number of acid particles (and therefore H⁺ ions) per unit volume.
• A dilute acid has a small number of acid particles per unit volume.

When you dilute an acid (add water), the pH increases (moves towards 7) because the concentration of H⁺ ions decreases.

When you dilute an alkali (add water), the pH decreases (moves towards 7) because the concentration of OH⁻ ions decreases.

Important Distinction

Concentration and strength are not the same thing. Concentration refers to how much solute is dissolved in a given volume. Strength refers to whether the acid fully or partially dissociates (see below).

You can have a concentrated weak acid (lots of ethanoic acid dissolved, but only partially dissociated) or a dilute strong acid (a small amount of hydrochloric acid dissolved, but fully dissociated).

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Strong and Weak Acids (Higher Tier)

This section applies to Higher tier students only.

Strong Acids

A strong acid is one that is completely (fully) ionised in aqueous solution. Every molecule of the acid dissociates to release H⁺ ions.

Examples of strong acids:

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)
• Nitric acid (HNO₃)

For hydrochloric acid:

HCl(aq) → H⁺(aq) + Cl⁻(aq)

This is a one-way arrow (→) because the reaction goes to completion — all the HCl molecules dissociate.

Weak Acids

A weak acid is one that is only partially ionised in aqueous solution. Only a small proportion of the acid molecules dissociate to release H⁺ ions at any given time. The rest remain as undissociated molecules.

Examples of weak acids:

• Ethanoic acid (CH₃COOH) — found in vinegar
• Citric acid (C₆H₈O₇) — found in citrus fruits
• Carbonic acid (H₂CO₃) — found in fizzy drinks

For ethanoic acid:

CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq)

This is a reversible arrow (⇌) because the equilibrium lies to the left — most molecules remain undissociated.

Comparing Strong and Weak Acids

For two acids at the same concentration:

Property	Strong Acid (e.g. HCl)	Weak Acid (e.g. CH₃COOH)
Degree of ionisation	Fully ionised	Partially ionised
[H⁺] in solution	High	Low
pH	Lower (e.g. pH 1)	Higher (e.g. pH 3)
Rate of reaction with Mg	Faster	Slower
Electrical conductivity	Higher	Lower
Arrow in equation	Single arrow (→)	Reversible arrow (⇌)

Exam Tip: The most common mistake students make is confusing concentration with strength. Dilute/concentrated describes how much acid is dissolved. Strong/weak describes whether it fully or partially dissociates. You can have a concentrated weak acid and a dilute strong acid. Make sure you use the correct terminology.

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Strong and Weak Bases (Higher Tier)

The same principle applies to bases:

• A strong base (e.g. NaOH) is completely ionised in solution.
• A weak base (e.g. NH₃) is only partially ionised in solution.

NaOH(aq) → Na⁺(aq) + OH⁻(aq) — strong base, fully ionised

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) — weak base, partially ionised

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Summary