4.4.2.1Reactions of acids with metals
Acids react with some metals to produce salts and hydrogen.
acid + metal → salt + hydrogen
Knowledge of these reactions is limited to those of magnesium, zinc and iron with hydrochloric and sulfuric acids. Hydrochloric acid gives chlorides and sulfuric acid gives sulfates, so those six combinations are the ones to be fluent in.
For example: Mg + 2HCl → MgCl₂ + H₂, and Zn + H₂SO₄ → ZnSO₄ + H₂.
The hydrogen is seen as fizzing, and it can be tested for with a lighted splint, which gives a squeaky pop.
4.4.2.1Acid and metal reactions as redoxHigher tier
Higher tier students should be able to explain, in terms of gain or loss of electrons, that these are redox reactions, and identify which species are oxidised and which are reduced in given chemical equations.
Take Mg + 2HCl → MgCl₂ + H₂. The chloride ions are spectators, so the ionic equation is Mg + 2H⁺ → Mg²⁺ + H₂. The magnesium atom loses two electrons to become Mg²⁺, so magnesium is oxidised. Each hydrogen ion gains an electron to become part of a hydrogen molecule, so the hydrogen ions are reduced. Both happen together, which is what makes it a redox reaction.
Common mistake
The acid itself is not "reduced" — name the species. It is the hydrogen ions that gain electrons, and it is the metal atoms that lose them.
4.4.2.2Neutralisation of acids and salt production
Acids are neutralised by alkalis, such as soluble metal hydroxides, and by bases, such as insoluble metal hydroxides and metal oxides, to produce salts and water. They are neutralised by metal carbonates to produce salts, water and carbon dioxide.
The particular salt produced in any reaction between an acid and a base or alkali depends on two things:
• the acid used — hydrochloric acid produces chlorides, nitric acid produces nitrates, sulfuric acid produces sulfates • the positive ions in the base, alkali or carbonate.
So the name of the salt is simply the metal from one reactant followed by the ending from the other. Copper oxide with sulfuric acid gives copper sulfate; sodium hydroxide with nitric acid gives sodium nitrate.
You should be able to predict products from given reactants, and use the formulae of common ions to deduce the formulae of salts.
The formula follows from the charges: the total positive charge must equal the total negative charge, so the numbers of each ion are whatever makes the compound neutral overall.
Worked example
Deduce the formula of calcium nitrate, given Ca²⁺ and NO₃⁻.
- The calcium ion carries 2+, and each nitrate ion carries 1−.
- For the compound to be neutral, two nitrate ions are needed for each calcium ion.
- Write the positive ion first, then the negative ion with the number needed.
- Because nitrate is a group of atoms, put brackets round it before the subscript.
Answer: Ca(NO₃)₂
Common mistake
Brackets matter. Ca(NO₃)₂ is two nitrate groups; CaNO₃₂ means nothing, and CaNO₆ is a different set of atoms entirely. Use brackets whenever more than one of a grouped ion is needed.
- Neutralisation
- the reaction of an acid with an alkali or base to produce a salt and water; hydrogen ions react with hydroxide ions to produce water
- Alkali
- a soluble base, such as a soluble metal hydroxide, whose aqueous solution contains hydroxide ions
- Base
- a substance that neutralises an acid, such as an insoluble metal hydroxide or a metal oxide
4.4.2.3Soluble salts
Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates. The solid is added to the acid until no more reacts, and the excess solid is filtered off to produce a solution of the salt. Salt solutions can be crystallised to produce solid salts.
Adding the solid in excess is the trick that makes this method work: it guarantees that all the acid is used up, so the filtered solution contains only the salt and water, with no acid left to contaminate the crystals.
You should be able to describe how to make pure, dry samples of named soluble salts from information provided.
Required practical 1
Preparation of a pure, dry sample of a soluble salt from an insoluble oxide or carbonate
- Measure a fixed volume of dilute acid into a beaker and warm it gently with a Bunsen burner — warming speeds the reaction up, but do not boil.
- Add the insoluble metal oxide or carbonate a spatula at a time, stirring, until no more reacts and some solid remains at the bottom. That excess is the signal that all the acid has been neutralised.
- Filter the mixture to remove the excess solid, keeping the filtrate — a solution of the salt.
- Pour the filtrate into an evaporating basin and heat it on a water bath or electric heater until crystals begin to form at the edges.
- Leave the solution to cool and crystallise slowly, then pat the crystals dry between filter papers.
Common mistake
Do not evaporate the solution to dryness over a Bunsen burner. Heating too far decomposes many salts and drives off the water of crystallisation, so the product is no longer the salt asked for.
4.4.2.4The pH scale and neutralisation
Acids produce hydrogen ions (H⁺) in aqueous solutions. Aqueous solutions of alkalis contain hydroxide ions (OH⁻).
The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.
A solution with pH 7 is neutral. Aqueous solutions of acids have pH values of less than 7, and aqueous solutions of alkalis have pH values greater than 7.
In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water. This reaction can be represented by the equation:
H⁺(aq) + OH⁻(aq) → H₂O(l)
That one equation is the whole of neutralisation. Whichever acid and whichever alkali are used, the reaction that removes the acidity is always the same — the rest of the ions are spectators and end up as the salt.
You should be able to describe the use of universal indicator, or a wide range indicator, to measure the approximate pH of a solution, and use the pH scale to identify acidic or alkaline solutions.
The method is to add a few drops of universal indicator to the solution — or dip a strip of indicator paper into it — and match the colour against the chart. The word to notice is approximate: universal indicator gives a pH to the nearest whole number at best, which is why a pH probe is used when an accurate value is needed.
Common mistake
Universal indicator is the wrong choice for a titration. It changes colour gradually through a range of pH values, so there is no sharp end point — a single indicator such as phenolphthalein is used instead.
- pH scale
- a scale from 0 to 14 measuring the acidity or alkalinity of a solution; pH 7 is neutral, below 7 is acidic and above 7 is alkaline
4.4.2.5Titrations
Chemistry only — not assessed in Combined Science: Trilogy.
The volumes of acid and alkali solutions that react with each other can be measured by titration using a suitable indicator.
You should be able to describe how to carry out titrations using strong acids and strong alkalis only — sulfuric, hydrochloric and nitric acids only — to find the reacting volumes accurately.
Required practical 2
Determination of the reacting volumes of a strong acid and a strong alkali by titration
- Use a pipette and pipette filler to transfer 25.0 cm³ of the alkali into a conical flask, and add a few drops of a single indicator such as phenolphthalein.
- Rinse the burette with the acid, fill it, and run some through to remove the air bubble below the tap. Record the initial reading to the nearest 0.05 cm³.
- Run the acid into the flask while swirling, quickly at first, then dropwise as the colour starts to change.
- Stop at the first permanent colour change and record the final reading. Subtract to get the titre.
- Repeat until two or more titres are concordant — within 0.10 cm³ of each other — and take the mean of those only.
Higher tier students should also be able to calculate the chemical quantities in titrations involving concentrations in mol/dm³ and in g/dm³ — the calculation set out under using concentrations of solutions.
Common mistake
Read the burette from the bottom of the meniscus, at eye level, and remember that a burette reads downwards — the volume delivered is the final reading minus the initial one, not the other way round.
- Titration
- a technique for measuring the volumes of acid and alkali solutions that react with each other, using a suitable indicator
4.4.2.6Strong and weak acidsHigher tier
A strong acid is completely ionised in aqueous solution. Examples of strong acids are hydrochloric, nitric and sulfuric acids.
A weak acid is only partially ionised in aqueous solution. Examples of weak acids are ethanoic, citric and carbonic acids.
For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.
You should be able to use and explain the terms dilute and concentrated, in terms of amount of substance, and weak and strong, in terms of the degree of ionisation, in relation to acids.
These are two independent ideas, and confusing them is the classic error here. Concentrated and dilute describe how much acid was dissolved in a given volume — the amount of substance. Strong and weak describe what fraction of that acid has ionised to release hydrogen ions. A dilute solution of a strong acid and a concentrated solution of a weak acid are both perfectly ordinary things.
As the pH decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of 10.
You should be able to describe neutrality and relative acidity in terms of the effect on hydrogen ion concentration and the numerical value of pH, using whole numbers only. This is an order of magnitude calculation: each whole step down the pH scale is a tenfold increase in hydrogen ion concentration.
Worked example
Solution A has pH 5 and solution B has pH 2. How many times greater is the hydrogen ion concentration in B?
- Find the difference in pH: 5 − 2 = 3 units.
- Each unit decrease multiplies the hydrogen ion concentration by 10.
- Three units means 10 × 10 × 10.
- So the factor is 10³.
Answer: 1000 times greater in solution B.
Common mistake
A lower pH means more hydrogen ions, not fewer. The scale runs the opposite way to the concentration, so pH 2 is a thousand times more acidic than pH 5, not three times.
- Strong acid
- an acid that is completely ionised in aqueous solution, such as hydrochloric, nitric and sulfuric acids
- Weak acid
- an acid that is only partially ionised in aqueous solution, such as ethanoic, citric and carbonic acids
- Concentrated
- describes a solution containing a large amount of solute in a given volume; a separate idea from strong, which is about the degree of ionisation
Before you practise
- acid + metal → salt + hydrogen; limited to magnesium, zinc and iron with hydrochloric and sulfuric acids.
- acid + base or alkali → salt + water; acid + metal carbonate → salt + water + carbon dioxide.
- The acid gives the ending — chlorides, nitrates, sulfates — and the base gives the metal.
- Required practical 1: add insoluble solid in excess, filter, then crystallise slowly to get a pure dry salt.
- Acids produce H⁺ ions in solution; alkalis contain OH⁻ ions; neutralisation is H⁺(aq) + OH⁻(aq) → H₂O(l).
- pH 7 is neutral, below 7 acidic, above 7 alkaline; universal indicator gives an approximate value, a pH probe an accurate one.
- Required practical 2 (chemistry only): titration with a pipette and burette, using concordant titres within 0.10 cm³.
- Higher tier: strong acids are completely ionised, weak acids only partially — separate from dilute and concentrated, which are about amount of substance.
- Higher tier: each decrease of one pH unit is a tenfold increase in hydrogen ion concentration.