Qualitative Analysis and pH Measurement

What these practical techniques are for

Chemists often need to answer two different questions about a sample. Qualitative analysis asks what substance or ion is present? Measuring pH asks how acidic or alkaline is this solution? Both depend on choosing a suitable test, controlling how it is carried out and recording an observation before drawing a conclusion.

OCR places these techniques within PAG 6. The practical work may form part of a larger investigation, but you must understand the reagents, apparatus, observations and reasoning because practical skills are also assessed in written examinations.

What pH tells you

When an acid dissolves in water, it produces hydrogen ions, $\text{H}^+(aq)$. An alkali in water produces hydroxide ions, $\text{OH}^-(aq)$. The pH scale represents how acidic or alkaline an aqueous solution is:

• below 7: acidic;
• 7: neutral;
• above 7: alkaline.

A lower pH means greater acidity, while a higher pH means greater alkalinity. The familiar classroom scale extends to about pH 14, although very acidic solutions can have pH values below 1.

Measuring pH with an indicator

An indicator changes colour according to pH. To estimate the pH of an unknown solution with universal indicator:

1. place a small sample in a clean test tube or spotting tile;
2. add a few drops of universal indicator, using the same amount for samples being compared;
3. compare the resulting colour with the chart supplied for that indicator;
4. record the estimated pH and whether the solution is acidic, neutral or alkaline.

The manufacturer's chart matters because universal-indicator formulations can differ slightly. The method gives a range or approximate whole-number pH rather than a highly precise measurement. A coloured or cloudy sample may also make the indicator colour difficult to judge.

A single indicator changes colour over only part of the pH scale. This makes a suitable single indicator useful for detecting a particular transition, but it cannot normally show pH across the full scale. Universal indicator covers a broad range because it contains a mixture of indicators.

Measuring pH with a meter

A digital pH meter uses a probe to produce a numerical reading. A suitable procedure is to calibrate the meter as instructed, rinse the probe with distilled or deionised water, place it in enough sample to immerse its sensitive end, wait for the reading to become steady and record the result at the resolution shown.

Rinse the probe between samples and blot it gently rather than wiping it, which could damage the sensitive surface or contaminate the next sample. A pH meter can distinguish smaller pH differences than an indicator, but only if it is calibrated and the probe is kept clean. Repeat readings can reveal an anomalous result.

How qualitative analysis identifies an unknown

A useful qualitative test must give a characteristic observation. The reasoning follows a clear chain:

unknown sample + chosen reagent or test → observation → identification

For example, “copper(II) ions are present” is a conclusion. “A blue precipitate formed when sodium hydroxide solution was added” is the evidence supporting it.

Use clean apparatus and a fresh portion of the unknown for each test. Add reagents dropwise at first, mix carefully and record colours, cloudiness, precipitates or gas-test results immediately. Label samples so that results cannot be exchanged accidentally. Suitable eye protection and the laboratory risk-control instructions are essential when handling unknowns, heating samples or testing gases.

Tests for common gases

Gas tests depend on a distinctive chemical effect:

• Hydrogen: hold a burning splint at the mouth of the test tube; hydrogen burns with a squeaky pop.
• Oxygen: insert a glowing splint; it relights.
• Carbon dioxide: bubble the gas through limewater, or shake it with limewater; the limewater turns milky or cloudy because a solid forms.
• Chlorine: expose damp litmus paper to the gas using the instructed small-scale method; the paper is bleached white.

The splint condition is important: oxygen requires a glowing splint, whereas hydrogen is tested with a burning splint. Do not place your face over a reaction vessel or try to identify an unknown gas by smelling it.

Flame tests for metal ions

Some metal ions give characteristic flame colours. A small amount of sample is introduced into a clean, non-luminous Bunsen flame and the colour is observed:

• lithium compounds: crimson;
• sodium compounds: yellow;
• potassium compounds: lilac;
• calcium compounds: orange-red;
• copper compounds: green.

The equipment must be clean because traces of another substance can alter the result. Flame tests are most dependable for single ions; in a mixture, a strong colour—especially sodium's yellow—may mask another ion.

Identifying cations by hydroxide precipitates

Adding sodium hydroxide solution to a solution containing certain metal ions forms an insoluble metal hydroxide. The precipitate colour provides evidence:

• copper(II): blue;
• iron(II): green;
• iron(III): brown;
• aluminium, calcium and magnesium: white.

The three white precipitates cannot be distinguished by colour alone. Aluminium hydroxide dissolves when excess sodium hydroxide is added, whereas the calcium and magnesium hydroxide precipitates do not in this test. This illustrates why the precise procedure—including whether reagent is added in excess—must be recorded.

A precipitation reaction can be represented by a net ionic equation. For copper(II):

$$\text{Cu}^{2+}(aq)+2\text{OH}^-(aq)\rightarrow\text{Cu(OH)}_2(s)$$

The state symbol $(s)$ shows that the observed product is a solid precipitate.

Identifying anions by precipitation

For halide ions, first acidify the sample with dilute nitric acid and then add silver nitrate solution:

• chloride ions form a white precipitate of silver chloride;
• bromide ions form a cream precipitate of silver bromide;
• iodide ions form a yellow precipitate of silver iodide.

For chloride ions, the net ionic equation is:

$$\text{Ag}^+(aq)+\text{Cl}^-(aq)\rightarrow\text{AgCl}(s)$$

For sulfate ions, acidify a fresh sample with dilute hydrochloric acid and add barium chloride solution. A white precipitate provides evidence for sulfate ions:

$$\text{Ba}^{2+}(aq)+\text{SO}_4^{2-}(aq)\rightarrow\text{BaSO}_4(s)$$

A carbonate can be tested using dilute acid. Effervescence occurs as carbon dioxide is released, and the gas turns limewater cloudy. The gas test confirms that the bubbles are carbon dioxide rather than relying on bubbling alone.

Building a justified identification

An unknown may need more than one test. Begin with its physical form and any safe initial observations, then choose tests that distinguish between the remaining possibilities. If an unknown ionic compound gives a lilac flame and its acidified solution gives a cream precipitate with silver nitrate, the two observations provide evidence about its cation and anion respectively.

Negative results also matter, provided the test was performed correctly. If no precipitate forms, this may rule out ions detected by that reagent, but it does not prove that the sample contains no ions at all. Mixtures are harder to interpret because colours can overlap or be masked.

Connecting these skills to titration

Qualitative tests identify species, while quantitative analysis determines an amount or concentration. Measuring pH and choosing an indicator prepare you for acid–alkali titration, where a carefully selected indicator marks an end-point. Universal indicator is useful for surveying a broad pH range, but its several gradual colour changes make it unsuitable for judging a precise titration end-point.