Standard solutions, primary standards and choosing an indicator

What a standard solution is

• A standard solution is a solution whose concentration is known accurately.

• Everything in your calculation is measured against it, so an error in its concentration is an error in your final answer, by the same percentage.

• This is why the Merit criteria name it explicitly: you must accurately determine the concentration of the standard solution, not read it off a bottle.

Primary standards

A primary standard is a substance pure and stable enough that you can make a standard solution simply by weighing it out. To qualify, a substance must be:

• Available in high purity, so the mass weighed really is the substance.
• Stable in air — not absorbing water or carbon dioxide, and not decomposing.
• Of high molar mass, so weighing errors are a small percentage of the mass.
• Soluble in water.

Substance	Primary standard?	Why
Anhydrous sodium carbonate, Na2CO3	Yes	pure, stable when dried, M = 106 g mol−1
Potassium hydrogen phthalate (KHP)	Yes	very pure, stable, M = 204 g mol−1
Oxalic acid dihydrate	Yes	stable, known water of crystallisation
Sodium hydroxide, NaOH	No	absorbs water and CO2 from the air
Hydrochloric acid	No	a solution; concentration changes as HCl evaporates
Potassium permanganate	No	slowly decomposes; oxidises traces of organic matter

• If your titrant is not a primary standard, it must be standardised — titrated against a solution made from one that is.

Making a standard solution from a primary standard

• Calculate the mass needed for the concentration and volume you want.

• Weigh it accurately on a 4-figure balance, recording the actual mass, not the target mass.

• Dissolve it completely in a beaker with a small amount of distilled water.

• Transfer to a volumetric flask, rinsing the beaker and stirring rod into the flask.

• Make up to the mark, then invert to mix at least ten times.

• Use the actual mass weighed in the calculation. Trying to hit an exact target mass wastes time and gains nothing — any mass near the target works, as long as you record it accurately.

Standardising a titrant

• Standardising means finding the true concentration of a solution by titrating it against a primary standard.
• For example, sodium hydroxide of roughly 0.1 mol L⁻¹ is titrated against a standard solution of sodium carbonate made by weighing:

$$2\text{HCl}(aq) + \text{Na}_2\text{CO}_3(aq) \rightarrow 2\text{NaCl}(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$

• Note this reaction's 2 : 1 ratio, which is one reason sodium carbonate is such a convenient primary standard for this standard's non-1:1 requirement.

Choosing an indicator

• An indicator changes colour over a narrow pH range. The right one changes colour at the pH of the equivalence point of your titration.

Titration	pH at equivalence	Indicator	Colour change
Strong acid + strong base	7	phenolphthalein or methyl orange	—
Weak acid + strong base (e.g. vinegar + NaOH)	above 7	phenolphthalein	colourless → faint pink
Strong acid + weak base (e.g. HCl + ammonia)	below 7	methyl orange	yellow → orange/red
Weak acid + weak base	—	none suitable	do not attempt — no sharp pH jump

• Phenolphthalein changes over about pH 8.3–10; methyl orange over about pH 3.1–4.4.
• The reason a weak acid titration needs phenolphthalein is that the salt formed is basic, so the equivalence point sits above pH 7 and methyl orange would change colour far too early.

Self-indicating titrations

• Some redox titrations need no indicator at all, because a reagent changes colour itself.

• Potassium permanganate is the classic: it is purple, and its product Mn²⁺ is colourless, so the endpoint is the first permanent faint pink from the first drop of unreacted MnO₄⁻.

• Iodine titrations use starch, added near the endpoint, giving a sharp blue-black to colourless change.

• Self-indicating titrations are a good choice for coloured products, since there is no indicator colour to be masked.