Chapter 3: Acids, Bases and Salts
When Antoine Lavoisier looked for what makes an acid sour and sharp, he thought he had found it. Every acid he could burn or break apart seemed to hold one element in common, and he gave that element a name built from the Greek words for ‘acid’ and ‘forming’. We still use his name for it: oxygen. For a generation, chemists taught that oxygen was the acid-maker.
One acid refused to behave. Muriatic acid — the acid we now call hydrochloric acid — was fiercely sour, attacked metals, and turned every indicator the way an acid should. Yet nobody could squeeze oxygen out of it. Chemists assumed the oxygen was simply hidden, locked in too tightly to find.
Around 1810, Humphry Davy in London decided to stop assuming. Working with the powerful electric batteries of his laboratory, he pulled the acid apart every way he knew and found no oxygen at all. What he found instead was hydrogen, and a choking greenish gas that he argued was an element in its own right. He named it chlorine.
Lavoisier’s rule had to go. The element common to every acid is not oxygen but hydrogen — and the story of acids, bases and salts begins from there.
Diagram
Fig. 3.0-1 — Two Views Of One Acid Two panels side by side. Left panel: an early nineteenth-century laboratory bench with a stoppered glass bottle labelled ‘muriatic acid’, a voltaic pile of stacked metal discs connected by wires to a glass vessel of the acid, and a gas jar collecting a pale greenish gas; label the bottle, the voltaic pile and the gas jar. Right panel: a modern reagent bottle labelled ‘hydrochloric acid, HCl’, with two callout arrows to the labels ‘hydrogen’ and ‘chlorine’, and a third callout reading ‘no oxygen’ struck through with a diagonal line.
3.1 The Three Classes Of Compounds
Elements do not join in any proportion they please. When two or more elements combine, they combine in a fixed ratio by mass, and the substance that results is a compound. Millions of compounds are known, and chemists have found ways of sorting them into families whose members behave alike.
One of the oldest and most useful of these sorts divides a large group of compounds into three classes — acids, bases and salts. Acids and bases are opposites in the way they behave, and a salt is what you are left with when an acid and a base have acted on each other and cancelled out.
The first clues came from the senses. Acids taste sour, as lemon juice and vinegar do. Bases taste bitter and feel soapy or slippery between the fingers. Salts, as a class, do neither.
These are useful hints, not laboratory tests. Mineral acids and strong alkalis burn skin and destroy tissue, so no acid or base in a laboratory is ever tasted, and none is handled without gloves. A chemist needs a way of telling an acid from a base that does not put a person at risk.
Telling Them Apart Safely
That way is an indicator — a substance that answers the question by changing colour. Dip a strip of blue litmus into an acid and it turns red; dip red litmus into an alkali and it turns blue. The colour change is sharp, it happens at once, and the chemist never has to come near the solution.
Indicators are dealt with fully in Section 3.5.2, where the common ones and their colours in acidic and basic media are set out together. Alongside colour, each class has its own set of chemical reactions, and it is these reactions — not taste — that finally decide what a compound is.
What The Three Classes Have In Common
The properties of the three classes were placed on a scientific footing by Lavoisier and by Davy, whose work you met at the start of this chapter. Stripped to their essentials, their findings are these.
- Every acid contains hydrogen as a constituent element, and in solution that hydrogen appears as the hydrogen ion, \(\mathrm{H^{+}}\), which is simply a proton.
- Every base is either a metallic oxide, containing the oxide ion \(\mathrm{O^{2-}}\), or a metallic hydroxide, containing the hydroxide ion \(\mathrm{OH^{-}}\), or else ammonia solution.
- Every salt is a compound formed by the interaction of an acid and a base, in which the hydrogen ions of the acid have been replaced by a metallic ion, or by the non-metallic ammonium ion, \(\mathrm{NH_4^{+}}\).
Examples of acids are hydrochloric acid, sulphuric acid and nitric acid. Examples of bases are copper oxide, sodium hydroxide and ammonium hydroxide. Examples of salts are sodium chloride, calcium carbonate and ammonium phosphate.
Read those three statements together and the pattern behind the whole chapter appears: an acid brings a replaceable hydrogen, a base brings something to replace it with, and the salt is the record of the exchange. Table 3.1 gathers the three classes side by side.
3.2 Acids And The Hydronium Ion
3.2.1 What An Acid Is
The word acid comes from the Latin word acidus, which means sour. The name is a good one. The sharp taste of lemon juice, of vinegar, of curd left standing and of an unripe tamarind is in every case the taste of an acid that the food contains.
Sourness, though, is a taste and not a test. To define an acid, chemists ask what an acid does when it is dissolved in water. The answer is always the same: it sets hydrogen free as a positively charged ion, and it produces no other kind of positive ion at all.
Two parts of that definition carry the weight. The acid must contain hydrogen, and that hydrogen must be able to leave the molecule in water. The hydronium ion is the form the departed hydrogen actually takes, and the next subsection explains how it comes about.
3.2.2 Formation Of The Hydronium Ion
When hydrogen chloride gas dissolves in water, the molecule splits. The bonding pair of electrons stays with the chlorine, so the chlorine leaves as a chloride ion carrying a negative charge, and the hydrogen leaves as a bare hydrogen ion carrying a positive charge. This splitting of a molecule into ions in water is called ionisation.
hydrogen chloride + water → hydrogen ion + chloride ion
\[\mathrm{HCl} \xrightarrow{\ \mathrm{H_2O}\ } \mathrm{H^{+}} + \mathrm{Cl^{-}}\]
Now look closely at what that hydrogen ion is. A hydrogen atom has one proton and one electron. Take the electron away and nothing is left but the proton — a particle with no electron cloud around it at all. Such a naked proton cannot exist on its own in water. It is seized at once by the lone pair of electrons on the oxygen atom of a nearby water molecule, and the two join to form a single ion, the hydronium ion.
hydrogen ion + water molecule → hydronium ion
\[\mathrm{H^{+}} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}}\]
This hydrated ion is stable and can exist independently in the solution. Its structure is shown in Fig. 3.2-1: three hydrogen atoms bonded to one oxygen atom, with the positive charge belonging to the ion as a whole.
Diagram
Fig. 3.2-1 — Structure Of The Hydronium Ion Electron-dot diagram of H₃O⁺ drawn as a pyramid: a central oxygen atom with three hydrogen atoms bonded to it and one lone pair of electrons on the oxygen. Show shared electron pairs as dot-and-cross pairs, mark the third O—H bond as the co-ordinate bond formed with the incoming H⁺, and enclose the whole ion in square brackets with a + sign outside the bracket. Label the oxygen, the hydrogens, the lone pair and the co-ordinate bond.
Every Acid Does The Same Thing
Hydrogen chloride is not a special case. Sulphuric acid, nitric acid and acetic acid all dissolve in water and all hand over hydrogen to water molecules, producing hydronium ions.
\[\mathrm{H_2SO_4} + 2\,\mathrm{H_2O} \rightleftharpoons 2\,\mathrm{H_3O^{+}} + \mathrm{SO_4^{2-}}\]
\[\mathrm{HNO_3} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{NO_3^{-}}\]
\[\mathrm{CH_3COOH} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{CH_3COO^{-}}\]
Notice that in every equation the negative ion is different — sulphate, nitrate, acetate — while the positive ion is the same in all three. Hydrogen is common to all acids, and the hydronium ion is common to all their solutions.
That single fact explains a great deal. The sour taste, the action on litmus, the attack on active metals, the fizzing with a carbonate — these are not properties of hydrochloric acid or of nitric acid separately. They are the properties of the hydronium ions that every acid puts into water, which is why acids as different as vinegar and battery acid behave alike.
Writing \(\mathrm{H_3O^{+}}\) in every equation soon becomes clumsy. For simplicity, chemists write the hydrogen ion as \(\mathrm{H^{+}}\) and take it as understood that in water it is really present as the hydronium ion. Both forms are used in this book, and they mean the same thing.
3.3 Classification Of Acids
Acids can be sorted in more than one way, and each way answers a different question. Where does the acid come from? What elements does its molecule contain? How many hydrogen ions can one molecule give away? The three classifications that follow answer those three questions in turn.
3.3.1 On The Basis Of Source
Organic acids are acids obtained usually from plants. Along with hydrogen, every one of them contains carbon.
Examples are oxalic acid, \(\mathrm{(COOH)_2}\), acetic acid, \(\mathrm{CH_3COOH}\), and formic acid, \(\mathrm{HCOOH}\).
Organic acids are weak acids. Their molecules do not ionise completely in solution — only a small fraction of them break up — so the solution contains unionised molecules as well as ions. That is why a squeeze of lemon on the tongue stings but does no damage.
Many of these acids are already familiar to you from the foods and materials in which they occur, as Table 3.2 shows.
| Organic acid | Where it occurs |
|---|---|
| Acetic acid | Vinegar |
| Citric acid | Citrus fruits (oranges and lemons) |
| Butyric acid | Rancid butter |
| Formic acid | Sting of bees and ants |
| Lactic acid | Sour milk, curd |
| Malic acid | Apples |
| Oleic acid | Olive oil |
| Oxalic acid | Tomatoes |
| Stearic acid | Fats |
| Tartaric acid | Grapes, tamarind and apples |
| Uric acid | Urine |
Table 3.2 Some naturally occurring organic acids
Inorganic acids, also called mineral acids, are acids obtained usually from minerals. As a rule they do not contain carbon. The one standing exception is carbonic acid, \(\mathrm{H_2CO_3}\), which is counted as an inorganic acid even though it contains carbon.
Examples are hydrochloric acid, \(\mathrm{HCl}\), sulphuric acid, \(\mathrm{H_2SO_4}\), nitric acid, \(\mathrm{HNO_3}\), and phosphoric acid, \(\mathrm{H_3PO_4}\).
These are strong acids. They ionise completely in solution, producing a high concentration of hydronium ions, so their solutions contain only ions and no unionised molecules of the acid remain.
Carbonic acid is the exception here too. It is a weak mineral acid: it ionises only partly, so its solution holds ions as well as molecules. It turns blue litmus pink rather than red, and it is non-corrosive, which is why it can be dissolved under pressure in soft drinks.
3.3.2 On The Basis Of Molecular Composition
The second classification looks only at which elements the acid molecule contains.
Oxyacids are acids which contain oxygen along with hydrogen and some other element. Nitric acid, \(\mathrm{HNO_3}\), and sulphuric acid, \(\mathrm{H_2SO_4}\), are oxyacids.
Hydracids contain hydrogen and a non-metallic element, and no oxygen at all. Hydrochloric acid, \(\mathrm{HCl}\), and hydrobromic acid, \(\mathrm{HBr}\), are hydracids.
The distinction is worth holding on to, because the two kinds are made in different ways. You will see in Section 3.4 that a hydracid can be built directly from its two elements, while an oxyacid is usually made by dissolving an oxide in water.
3.3.3 On The Basis Of Basicity
The third classification counts. When one molecule of an acid ionises in water, how many hydronium ions does it yield? That number is the acid’s basicity.
Monobasic Acids
Monobasic acids are acids which on ionisation in water produce one hydronium ion per molecule of the acid. Hydrochloric acid is the plainest example.
\[\mathrm{HCl} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}} + \mathrm{Cl^{-}}\]
Other monobasic acids are \(\mathrm{HBr}\), \(\mathrm{HNO_3}\), \(\mathrm{HI}\) and \(\mathrm{CH_3COOH}\). A monobasic acid ionises in one step, and because it has only one hydrogen to give away, it can form only one salt with an alkali.
\[\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}\]
Here a warning is needed, and it is the point on which students most often slip. Basicity does not depend on the number of hydrogen atoms in one molecule of the acid. It depends on the number of ionisable hydrogen atoms — those that can actually leave as \(\mathrm{H^{+}}\).
Acetic acid, \(\mathrm{CH_3COOH}\) or \(\mathrm{C_2H_4O_2}\), contains four hydrogen atoms in its molecule. Even so, it is monobasic, because its molecule ionises by liberating only one hydronium ion. The other three hydrogens are bonded to carbon and stay where they are.
\[\mathrm{CH_3COOH} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{CH_3COO^{-}}\]
Dibasic Acids
Dibasic acids are acids which on ionisation in water produce two hydronium ions per molecule of the acid. Sulphuric acid, \(\mathrm{H_2SO_4}\), oxalic acid, \(\mathrm{(COOH)_2}\), sulphurous acid, \(\mathrm{H_2SO_3}\), and carbonic acid, \(\mathrm{H_2CO_3}\), are dibasic.
A dibasic acid does not lose both hydrogens at once. It ionises in two steps, giving up one hydrogen in each.
\[\mathrm{H_2SO_4} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{HSO_4^{-}}\]
\[\mathrm{HSO_4^{-}} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{SO_4^{2-}}\]
\[\mathrm{H_2CO_3} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{HCO_3^{-}}\]
\[\mathrm{HCO_3^{-}} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{CO_3^{2-}}\]
Because a dibasic acid has two replaceable hydrogen ions, it can form two different salts with an alkali. Replace one hydrogen and you get an acid salt; replace both and you get a normal salt. These two kinds of salt are classified fully in Section 3.15.
\[\mathrm{NaOH} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{H_2O}\]
\[2\,\mathrm{NaOH} + \mathrm{H_2SO_4} \rightarrow \mathrm{Na_2SO_4} + 2\,\mathrm{H_2O}\]
Sodium hydrogen sulphate, \(\mathrm{NaHSO_4}\), is the acid salt and sodium sulphate, \(\mathrm{Na_2SO_4}\), the normal salt. Which one you get depends simply on how much alkali you supply.
Phosphorous acid, \(\mathrm{H_3PO_3}\), is a useful case to test yourself on. Its molecule holds three hydrogen atoms, yet it is a dibasic acid. In the oxyacids of phosphorus, only the hydrogen atoms attached to oxygen atoms are replaceable; a hydrogen bonded directly to the phosphorus atom is not. Fig. 3.3-1 shows why.
Diagram
Fig. 3.3-1 — Structure Of Phosphorous Acid Structural formula of H₃PO₃: a central phosphorus atom with a double bond to one oxygen atom drawn upwards, two OH groups drawn to the left and below, and one hydrogen atom bonded directly to the phosphorus on the right. Label the two OH hydrogens ‘replaceable’ and the hydrogen joined straight to phosphorus ‘not replaceable’.
Tribasic Acids
Tribasic acids are acids which on ionisation in water produce three hydronium ions per molecule of the acid. Phosphoric acid, \(\mathrm{H_3PO_4}\), and citric acid, \(\mathrm{C_3H_5O(COOH)_3}\), are tribasic. They ionise in three steps.
\[\mathrm{H_3PO_4} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{H_2PO_4^{-}}\]
\[\mathrm{H_2PO_4^{-}} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{HPO_4^{2-}}\]
\[\mathrm{HPO_4^{2-}} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{PO_4^{3-}}\]
Three replaceable hydrogen ions mean three possible salts with an alkali — two acid salts and one normal salt.
\[\mathrm{NaOH} + \mathrm{H_3PO_4} \rightarrow \mathrm{NaH_2PO_4} + \mathrm{H_2O}\]
\[2\,\mathrm{NaOH} + \mathrm{H_3PO_4} \rightarrow \mathrm{Na_2HPO_4} + 2\,\mathrm{H_2O}\]
\[3\,\mathrm{NaOH} + \mathrm{H_3PO_4} \rightarrow \mathrm{Na_3PO_4} + 3\,\mathrm{H_2O}\]
The three products are monosodium dihydrogen phosphate, disodium hydrogen phosphate and trisodium phosphate. Counting is not restricted to three: silicic acid, \(\mathrm{H_4SiO_4}\), has basicity four.
3.3.4 Concentration And Strength Of An Acid
Two words are used loosely in everyday speech and precisely in chemistry, and confusing them costs marks and, in a laboratory, sometimes worse. The words are concentrated and strong.
Concentration of an acid means the amount of acid present in a definite amount of its aqueous solution. It is nothing more than a statement about how much water is present.
An acid which contains a very small amount of water, or none at all, is called a concentrated acid. An acid which contains far more water than its own mass is called a dilute acid.
Strength is a different quantity altogether. The strength of an acid is the measure of the concentration of hydronium ions it produces in its aqueous solution. Put simply, the more \(\mathrm{H^{+}}\) ions an acid releases in water, the stronger it is.
A strong acid ionises almost completely, so nearly every molecule contributes a hydronium ion. A weak acid ionises only partly, so most of its molecules stay whole and only a few hydronium ions appear. Fig. 3.3-2 shows the difference at the level of the particles.
Diagram
Fig. 3.3-2 — Strong And Weak Acids In Solution Two beakers side by side, each containing the same number of acid particles. Left beaker labelled ‘strong acid’: show almost all particles drawn as separated H⁺ and anion pairs, with only one or two whole molecules. Right beaker labelled ‘weak acid’: show mostly whole undissociated molecules with only two or three separated H⁺ and anion pairs. Label the free H⁺ ions in both beakers and add the note that both beakers hold the same total amount of acid.
Now the two ideas can be put together. Dilute hydrochloric acid is a stronger acid than highly concentrated acetic acid, because dilute hydrochloric acid releases more \(\mathrm{H^{+}}\) ions. Concentration tells you how much water is in the bottle; strength tells you how readily the acid gives up its hydrogen. They are independent of each other.
Chemists measure this by the degree of ionisation, written \(\alpha\), which states what fraction of the acid’s molecules have actually ionised in solution.
\[\alpha = \frac{\text{number of acid molecules ionised}}{\text{total number of acid molecules present in aqueous solution}} \times 100\]
If the degree of ionisation for an acid, or for an alkali, in aqueous solution is greater than 30 per cent, it is strong; if it is less than 30 per cent, it is weak.
Diluting An Acid Safely
Diluting a concentrated acid is one of the few laboratory operations that can injure someone in a single careless second. The rule is fixed: pour the acid into the water, in small amounts, stirring constantly. Water is never added to acid.
The reason lies in the heat. Mixing acid and water is strongly exothermic. If water is poured onto concentrated acid, all that heat is released in the thin top layer, which can boil at once and spit acid out of the container; the container itself may crack. Adding acid slowly to a large volume of water spreads the heat through the whole of the water, which warms gently instead.
3.4 Preparation Of Acids
An acid is not always waiting to be scooped out of a bottle. Four standard routes are used to make one, and which route you choose depends on the acid you want. Two of them build the acid from elements, one oxidises a non-metal, and one drives a ready-made acid out of one of its own salts.
3.4.1 By Synthesis
Binary acids — acids whose molecules contain just two elements — can be made by combining those two elements directly, under suitable conditions.
hydrogen + chlorine → hydrogen chloride
\[\mathrm{H_2}\text{(g)} + \mathrm{Cl_2}\text{(g)} \rightarrow 2\,\mathrm{HCl}\text{(g)}\]
\[\mathrm{H_2}\text{(g)} + \mathrm{Br_2}\text{(g)} \rightarrow 2\,\mathrm{HBr}\text{(g)}\]
\[\mathrm{H_2}\text{(g)} + \mathrm{S}\text{(s)} \rightarrow \mathrm{H_2S}\text{(g)}\]
Each product is a gas, and each becomes an acid only when it is dissolved in water — hydrogen chloride gas dissolved in water is hydrochloric acid. Notice that all three are hydracids, which is exactly what you would expect: no oxygen is present anywhere in the reaction.
3.4.2 By The Action Of Water On Acidic Oxides
Certain oxides of non-metals dissolve in water and the solution turns out to be acidic. Such an oxide is called an acidic oxide, or an acidic anhydride — the second name means the acid with its water taken away, so adding the water back gives the acid again.
sulphur trioxide + water → sulphuric acid
\[\mathrm{SO_3} + \mathrm{H_2O} \rightarrow \mathrm{H_2SO_4}\]
\[\mathrm{SO_2} + \mathrm{H_2O} \rightarrow \mathrm{H_2SO_3}\]
\[\mathrm{CO_2} + \mathrm{H_2O} \rightarrow \mathrm{H_2CO_3}\]
\[\mathrm{P_2O_5} + 3\,\mathrm{H_2O} \rightarrow 2\,\mathrm{H_3PO_4}\]
\[\mathrm{N_2O_5} + \mathrm{H_2O} \rightarrow 2\,\mathrm{HNO_3}\]
These give sulphuric acid, sulphurous acid, carbonic acid, phosphoric acid and nitric acid in turn. Because water is being added to an oxide, every acid made this way must contain oxygen: only oxyacids can be prepared by this method.
Nitrogen dioxide behaves oddly and has a name of its own for it. When \(\mathrm{NO_2}\) reacts with water, two different acids are formed at once — nitrous acid and nitric acid — so it is called a mixed acid anhydride, or double acid anhydride.
\[2\,\mathrm{NO_2} + \mathrm{H_2O} \rightarrow \mathrm{HNO_2} + \mathrm{HNO_3}\]
3.4.3 By The Oxidation Of Non-Metals
A non-metal can also be converted straight into its acid by oxidising it. Concentrated nitric acid is the oxidising agent used, and it is itself reduced to nitrogen dioxide in the process.
sulphur + concentrated nitric acid → sulphuric acid + water + nitrogen dioxide
\[\mathrm{S} + 6\,\mathrm{HNO_3} \rightarrow \mathrm{H_2SO_4} + 2\,\mathrm{H_2O} + 6\,\mathrm{NO_2}\]
\[\mathrm{P} + 5\,\mathrm{HNO_3} \rightarrow \mathrm{H_3PO_4} + \mathrm{H_2O} + 5\,\mathrm{NO_2}\]
Sulphur gives sulphuric acid and phosphorus gives phosphoric acid. Nitrogen dioxide is a dense brown gas and is poisonous, so a reaction of this kind is carried out only in a fume cupboard.
3.4.4 By Displacement Of A More Volatile Acid
The last method uses no elements at all. Take the normal salt of an acid that vaporises easily, add an acid that does not, and warm the mixture: the volatile acid escapes as a vapour and can be condensed and collected, while the non-volatile acid stays behind in the salt.
A volatile acid is one that vaporises easily — at room temperature or on gentle heating — while a non-volatile acid does not. Sulphuric acid is the non-volatile acid used in practice, and both hydrochloric acid and nitric acid are prepared from their salts in this way.
sodium chloride + concentrated sulphuric acid → sodium hydrogen sulphate + hydrogen chloride
\[\mathrm{NaCl} + \mathrm{H_2SO_4} \xrightarrow{\ \Delta\ } \mathrm{NaHSO_4} + \mathrm{HCl}\]
\[\mathrm{NaNO_3} + \mathrm{H_2SO_4} \xrightarrow{\ \Delta\ } \mathrm{NaHSO_4} + \mathrm{HNO_3}\]
The acid that is driven off is the more volatile one, which is why the method works in this direction and not the other. You cannot use hydrochloric acid on a sulphate to make sulphuric acid, because there is nothing to carry the sulphuric acid away.
3.5 Properties Of Acids
3.5.1 Physical Properties
Taste. Acids have a sour taste. This is a property of the class, not a test to be used: mineral acids such as sulphuric acid and nitric acid are highly corrosive, and no acid in a laboratory is ever tasted.
Physical state. Some acids are solids at room temperature and some are liquids, as Table 3.3 sets out. Among the liquids, most are volatile and vaporise easily — at room temperature or on heating to about \(100^\circ\text{C}\) — while sulphuric acid is non-volatile.
| Acid | Formula | State and behaviour |
|---|---|---|
| Boric acid | \(\mathrm{H_3BO_3}\) | Solid |
| Oxalic acid | \(\mathrm{(COOH)_2}\) | Solid |
| Tartaric acid | \(\mathrm{C_4H_6O_6}\) | Solid |
| Citric acid | \(\mathrm{C_6H_8O_7}\) | Solid |
| Phosphoric acid | \(\mathrm{H_3PO_4}\) | Solid |
| Acetic acid | \(\mathrm{CH_3COOH}\) | Liquid, volatile |
| Formic acid | \(\mathrm{HCOOH}\) | Liquid, volatile |
| Carbonic acid | \(\mathrm{H_2CO_3}\) | Liquid, volatile |
| Hydrochloric acid | \(\mathrm{HCl}\) | Liquid, volatile |
| Nitric acid | \(\mathrm{HNO_3}\) | Liquid, volatile |
| Sulphurous acid | \(\mathrm{H_2SO_3}\) | Liquid, volatile |
| Sulphuric acid | \(\mathrm{H_2SO_4}\) | Liquid, non-volatile |
Table 3.3 Physical state of some common acids
Effect on skin. All mineral acids have a corrosive action on the skin and cause painful burns and blisters. Each leaves its own mark: concentrated sulphuric acid chars the skin black, concentrated nitric acid stains it yellow, and concentrated hydrochloric acid gives it an amber colour. Carbonic acid, though a mineral acid, is not corrosive, and neither are the organic acids. Acids are handled with gloves and eye protection, and any splash on skin is washed at once under running water.
Conduction. Acids conduct electricity in the aqueous state. They are electrolytes, and the reason is one you already know — an acid in water is full of ions, and it is the movement of those ions that carries the current.
3.5.2 Action On Indicators
Indicators are complex substances that acquire distinct colours in acidic and in basic media. The three in everyday laboratory use are litmus, methyl orange and phenolphthalein.
Litmus is a natural colouring matter, obtained from certain lichens — small plants that grow on rocks and bark — by dissolving the extract in distilled water. Methyl orange and phenolphthalein are synthetic indicators, prepared in the laboratory. A universal indicator is a mixture of several different pH indicators combined, and it is explained in Section 3.13.2.
Table 3.4 gives the colours you should be able to recall without hesitating.
| Indicator | Neutral solution | Acidic solution | Basic medium |
|---|---|---|---|
| Litmus | Purple | Red | Blue |
| Methyl orange | Orange | Red or pink | Yellow |
| Phenolphthalein | Colourless | Colourless | Pink |
| Universal indicator | Green | Red | Blue or purple |
Table 3.4 Colour shifts of common acid–base indicators
Read the phenolphthalein row carefully. It is colourless in a neutral solution and colourless in an acid, so a colourless result tells you only that the solution is not basic. Litmus and methyl orange, by contrast, give a different colour in each of the three cases.
Indicators need not come from a laboratory shelf. Red cabbage extract, a natural indicator, remains red in acidic solutions but turns green in basic ones. Stranger still are the olfactory indicators — substances whose smell, rather than whose colour, changes with the nature of the solution. Onion, vanilla and clove oil are examples. When a basic solution such as sodium hydroxide solution is added to a cloth strip treated with onion extract, the onion smell can no longer be detected; an acidic solution does not destroy the smell.
3.5.3 Reaction With Active Metals
Drop a piece of zinc into dilute hydrochloric acid and bubbles rise from the metal almost at once. The gas is hydrogen, and what is left in solution is a salt of the metal.
zinc + dilute hydrochloric acid → zinc chloride + hydrogen
Both dilute hydrochloric acid and dilute sulphuric acid behave this way. With very active metals such as potassium, sodium and calcium the reaction is explosive and is never attempted as a preparation; with less active metals such as magnesium and zinc it is moderate and controllable.
\[\mathrm{Mg}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{MgCl_2}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{Zn}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{Fe}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
The gas is confirmed by bringing a burning splint to the mouth of the tube: hydrogen burns with a squeaky pop, and therefore the gas collected is hydrogen. Fig. 3.5-1 shows the arrangement.
Diagram
Fig. 3.5-1 — Action Of A Dilute Acid On Zinc Conical flask on a bench holding granulated zinc covered by dilute hydrochloric acid, closed with a two-holed stopper carrying a thistle funnel dipping below the liquid and a bent delivery tube. The delivery tube leads under a gas jar inverted over water in a trough, collecting hydrogen by downward displacement of water. Label the granulated zinc, the dilute hydrochloric acid, the thistle funnel, the delivery tube, the trough of water and the gas jar of hydrogen.
Nitric acid is the exception, and the reason is worth understanding. Nitric acid is a very strong oxidising agent, so as fast as hydrogen is formed it is oxidised to water — which defeats the purpose of the reaction. Nitric acid is therefore not used to prepare hydrogen. Only magnesium and manganese can produce hydrogen with very dilute nitric acid, of about 1 per cent strength.
\[\mathrm{Mn} + 2\,\mathrm{HNO_3}\text{(dilute, 1%)} \rightarrow \mathrm{Mn(NO_3)_2} + \mathrm{H_2}\text{(g)}\]
3.5.4 Reaction With Bases
An acid and a base destroy each other’s characteristic properties. The hydrogen ions of the acid and the hydroxide or oxide ions of the base combine to form water, and the remaining ions stay in solution as a salt. Nothing else is produced — the word only in that sentence is what makes the reaction a neutralisation.
acid + base → salt + water
\[\mathrm{CuO}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{CuSO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]
\[\mathrm{NaOH}\text{(aq)} + \mathrm{HNO_3}\text{(aq)} \rightarrow \mathrm{NaNO_3}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]
The first uses a basic oxide and the second an alkali, and both give salt and water alone. Neutralisation is taken up again in Section 3.19.5, where it is written as an ionic equation and used as a method of preparing salts.
3.6 Action Of Dilute Acids On Salts
A dilute acid does not only attack metals and bases. Poured on certain salts, it breaks them up and drives out a gas. Which gas appears depends entirely on which salt you started from, and that makes this reaction one of the most useful tools a chemist has for identifying an unknown solid.
The idea behind all four cases is the same. The acid supplies hydrogen ions, these take hold of the negative ion of the salt, and the unstable acid so formed falls apart into water and a gas that escapes.
3.6.1 Carbonates And Hydrogen Carbonates
Acids liberate carbon dioxide on reaction with metallic carbonates and hydrogen carbonates. Drop dilute hydrochloric acid on a piece of marble and it fizzes strongly as the gas comes off.
calcium carbonate + dilute hydrochloric acid → calcium chloride + water + carbon dioxide
\[\mathrm{CaCO_3}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{CaCl_2}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{CO_2}\text{(g)}\]
\[\mathrm{Ca(HCO_3)_2}\text{(aq)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{CaCl_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + 2\,\mathrm{CO_2}\text{(g)}\]
Notice that the hydrogen carbonate gives twice as much carbon dioxide, because each formula unit carries two \(\mathrm{HCO_3^{-}}\) groups.
Pass the gas through lime water and the lime water turns milky, and therefore the gas is carbon dioxide. This is the standard test, and Fig. 3.6-1 shows how it is set up.
Diagram
Fig. 3.6-1 — Testing The Gas From A Carbonate Test tube on a stand containing marble chips covered with dilute hydrochloric acid, closed by a stopper with a bent delivery tube. The delivery tube dips into a second test tube containing clear lime water, which is shown turning milky. Label the marble chips, the dilute hydrochloric acid, the delivery tube and the lime water, and mark the effervescence at the surface of the acid.
Two limits apply to this reaction, and both are asked about often.
If the salt produced is insoluble, the reaction does not proceed. The insoluble product forms a crust over the solid and seals it off from fresh acid. So lead carbonate does not react with hydrochloric acid or with sulphuric acid, and calcium carbonate does not react with sulphuric acid — in each case the salt that would form is insoluble.
Among the hydrogen carbonates, only those of sodium and potassium are common bicarbonates that exist in the solid state. The rest are known only in solution.
3.6.2 Sulphites And Hydrogen Sulphites
Acids react with sulphites and hydrogen sulphites of metals to liberate sulphur dioxide.
calcium sulphite + dilute hydrochloric acid → calcium chloride + water + sulphur dioxide
\[\mathrm{CaSO_3}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{CaCl_2}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
\[\mathrm{NaHSO_3}\text{(s)} + \mathrm{HCl}\text{(aq)} \rightarrow \mathrm{NaCl}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
Sulphur dioxide is recognised by its sharp, choking smell of burning sulphur. It irritates the eyes and the lungs, so the tube is never held to the face and the reaction is carried out in a fume cupboard.
3.6.3 Sulphides
Acids react with metal sulphides to liberate hydrogen sulphide.
zinc sulphide + dilute hydrochloric acid → zinc chloride + hydrogen sulphide
\[\mathrm{ZnS}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + \mathrm{H_2S}\text{(g)}\]
\[\mathrm{FeS}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2S}\text{(g)}\]
Hydrogen sulphide smells strongly of rotten eggs. It is poisonous, and dangerous partly because the nose stops noticing the smell after a short exposure, so it too is handled only in a fume cupboard.
Table 3.5 collects the three gases together, since telling them apart is the whole point of these reactions.
3.6.4 Chlorides And Nitrates
Chlorides and nitrates are, as a rule, unmoved by dilute acids. There is a small group of exceptions, and they are exceptions for a clear reason: in each of them the product that forms is insoluble and drops out of solution, which pulls the reaction forward.
\[\mathrm{BaCl_2}\text{(aq)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow 2\,\mathrm{HCl}\text{(aq)} + \mathrm{BaSO_4}\!\downarrow\]
\[\mathrm{Pb(NO_3)_2}\text{(aq)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow 2\,\mathrm{HNO_3}\text{(aq)} + \mathrm{PbSO_4}\!\downarrow\]
\[\mathrm{Pb(NO_3)_2}\text{(aq)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow 2\,\mathrm{HNO_3}\text{(aq)} + \mathrm{PbCl_2}\!\downarrow\]
The downward arrow marks a solid precipitate. Barium sulphate and lead sulphate are insoluble, and so is lead chloride in cold water.
Apart from these, chlorides and nitrates react only with concentrated sulphuric acid, and only on warming, when they give up hydrogen chloride and nitric acid respectively. This is the displacement route you met in Section 3.4.4, and the product depends on the temperature.
\[\mathrm{NaCl} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ \text{below }200^\circ\text{C}\ } \mathrm{NaHSO_4} + \mathrm{HCl}\]
\[2\,\mathrm{NaCl} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ \text{above }200^\circ\text{C}\ } \mathrm{Na_2SO_4} + 2\,\mathrm{HCl}\]
\[\mathrm{KNO_3} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ \text{below }200^\circ\text{C}\ } \mathrm{KHSO_4} + \mathrm{HNO_3}\]
\[2\,\mathrm{KNO_3} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ \text{above }200^\circ\text{C}\ } \mathrm{K_2SO_4} + 2\,\mathrm{HNO_3}\]
Below \(200^\circ\text{C}\) only one hydrogen of the sulphuric acid is used and the acid salt is left behind; above \(200^\circ\text{C}\) both are used and the normal salt is formed. The temperature, not the reagents, decides which.
3.7 Uses Of Some Common Acids
Acids are not confined to the laboratory. They preserve food, clean metal, treat the eye, feed crops and flavour drinks, and in each case the use follows directly from the chemistry. A weak, non-corrosive acid can be swallowed or put on the skin; a strong, corrosive one is put to work on metal and rock instead.
Table 3.6 lists the acids you should be able to place.
| Acid | Use |
|---|---|
| Boric acid | Eye-wash and antiseptic |
| Citric acid | Food preservation; preparation of vitamin C |
| Oxalic acid | Ink stain remover |
| Carbonic acid | Flavoured drinks |
| Tartaric acid | Baking powder |
| Acetic acid | Table vinegar, and in cooking |
| Hydrochloric acid | Cleaning of metal items, called pickling |
| Benzoic acid | Preservation of food, and making of perfumes and medicines |
| Nitric acid | Explosives |
| Phosphoric acid | Fertilizers |
Table 3.6 General uses of some acids
Look down the list and a pattern appears. The organic acids — citric, acetic, tartaric, benzoic — are weak, and every one of them is used in or on the body. The strong mineral acids are kept for work on materials: hydrochloric acid strips the oxide film from steel before it is coated, and phosphoric acid ends up in the fertilizer that supplies phosphorus to a crop.
3.8 Bases And Alkalis
3.8.1 What A Base Is
An acid puts hydronium ions into water. A base is the compound that takes them out again. Bring the two together and the hydronium ions are used up, water is formed, and what remains is a salt — nothing more.
Copper oxide, \(\mathrm{CuO}\), is a base of the first kind; magnesium hydroxide, \(\mathrm{Mg(OH)_2}\), is a base of the second. Ammonium hydroxide is added to the definition because it behaves in every way like a base while containing no metal at all.
\[\mathrm{CuO} + 2\,\mathrm{HCl} \rightarrow \mathrm{CuCl_2} + \mathrm{H_2O}\]
\[\mathrm{Mg(OH)_2} + \mathrm{H_2SO_4} \rightarrow \mathrm{MgSO_4} + 2\,\mathrm{H_2O}\]
\[\mathrm{Al(OH)_3} + 3\,\mathrm{HNO_3} \rightarrow \mathrm{Al(NO_3)_3} + 3\,\mathrm{H_2O}\]
\[\mathrm{NH_4OH} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl} + \mathrm{H_2O}\]
In every one of these, the products are a salt and water and nothing else. Hold on to that word only — the next subsection shows what it is doing there.
3.8.2 Basic Oxides And Basic Hydroxides
A basic oxide is a metallic oxide which contains the oxide ion, \(\mathrm{O^{2-}}\), and reacts with an acid to form salt and water only.
The word only is not decoration. Leave it out, and compounds that are quite unlike ordinary basic oxides would slip into the class. Lead(IV) oxide is a metallic oxide, and it does react with hydrochloric acid to give lead(II) chloride, which is a salt, and water.
\[\mathrm{PbO_2}\text{(s)} + 4\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{PbCl_2}\text{(aq)} + \mathrm{Cl_2}\text{(g)} + 2\,\mathrm{H_2O}\text{(l)}\]
But look at the products: chlorine is given off as well. Because something other than salt and water is formed, lead(IV) oxide is excluded, and \(\mathrm{PbO_2}\) is not a base.
A basic hydroxide is a metallic hydroxide which contains the hydroxide ion, \(\mathrm{OH^{-}}\), and reacts with an acid to give salt and water only. Sodium hydroxide, \(\mathrm{NaOH}\), and aluminium hydroxide, \(\mathrm{Al(OH)_3}\), are basic hydroxides.
\[\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}\]
3.8.3 Alkalis
Most bases are solids that will not dissolve. All metallic oxides and hydroxides are insoluble in water, except those of sodium and potassium, and that of calcium, which is slightly soluble. The few that do dissolve get a name of their own.
An alkali, in short, is a base that is soluble in water.
\[\mathrm{NaOH}\text{(aq)} \rightleftharpoons \mathrm{Na^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)}\]
Every alkali releases hydroxide ions into solution as the only negative ions, and it is those hydroxide ions that give an alkaline solution its properties — the bitter taste, the soapy feel, and the turning of red litmus blue. This is the exact mirror of the acids: there, one ion, \(\mathrm{H_3O^{+}}\), was responsible for all the common properties.
The alkalis you will meet are these four.
- Sodium hydroxide, \(\mathrm{NaOH}\), known as caustic soda.
- Potassium hydroxide, \(\mathrm{KOH}\), known as caustic potash.
- Calcium hydroxide, \(\mathrm{Ca(OH)_2}\), known as slaked lime.
- Ammonium hydroxide, \(\mathrm{NH_4OH}\).
The first two are strong alkalis, because they are completely dissociated in their aqueous solutions. The last two are weak alkalis, because they are only feebly dissociated or ionised in their solutions.
Every Alkali Is A Base, But Not Every Base Is An Alkali
This single sentence causes more trouble than any other in the chapter, and Fig. 3.8-1 is worth studying until it is obvious. Ferric hydroxide, \(\mathrm{Fe(OH)_3}\), and cupric hydroxide, \(\mathrm{Cu(OH)_2}\), are bases — they neutralise acids to give salt and water. They are not alkalis, because they are insoluble in water and so can never put hydroxide ions into solution.
Diagram
Fig. 3.8-1 — Bases And Alkalis A large oval labelled ‘BASES: metallic oxides and hydroxides that give salt and water only with an acid’, containing a smaller oval inside it labelled ‘ALKALIS: bases soluble in water’. Place NaOH, KOH, Ca(OH)₂ and NH₄OH inside the inner oval; place CuO, Fe(OH)₃, Cu(OH)₂ and Mg(OH)₂ in the outer region between the two ovals; place PbO₂ outside the large oval altogether with the note ‘not a base — gives chlorine as well’.
Table 3.7 sets the two side by side.
3.9 Classification Of Bases
Bases are sorted in the same two ways acids were: by how completely they break up in water, and by how many ions one molecule of them can supply. The parallel is exact, and holding the two side by side makes both easier to remember.
3.9.1 On The Basis Of Strength
A strong alkali undergoes almost complete dissociation in aqueous solution to produce a high concentration of hydroxide ions. Sodium hydroxide and potassium hydroxide are strong alkalis.
\[\mathrm{NaOH}\text{(aq)} \rightleftharpoons \mathrm{Na^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)}\]
\[\mathrm{KOH}\text{(aq)} \rightleftharpoons \mathrm{K^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)}\]
A weak alkali undergoes only partial dissociation or ionisation in aqueous solution, and so produces a low concentration of hydroxide ions in the solution. Ammonium hydroxide and calcium hydroxide are weak alkalis.
\[\mathrm{NH_4OH}\text{(aq)} \rightleftharpoons \mathrm{NH_4^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)}\]
\[\mathrm{Ca(OH)_2}\text{(aq)} \rightleftharpoons \mathrm{Ca^{2+}}\text{(aq)} + 2\,\mathrm{OH^{-}}\text{(aq)}\]
Compare these with the equations of Section 3.3.4 and the symmetry is plain. An acid is strong when nearly every molecule releases its \(\mathrm{H^{+}}\); an alkali is strong when nearly every formula unit releases its \(\mathrm{OH^{-}}\). In both cases strength is about the fraction that dissociates, and has nothing to do with how much water is present.
3.9.2 On The Basis Of Acidity
The second classification counts hydroxide ions, exactly as basicity counted hydronium ions.
Both halves of that definition describe the same number, which is useful, because many bases never dissolve at all. For a water-insoluble hydroxide, the acidity is simply the number of replaceable \(\mathrm{OH^{-}}\) ions in one molecule of the base.
A monoacidic base dissociates in the molten, or fused, state, or in aqueous solution, to produce one hydroxide ion per molecule of the base. Sodium hydroxide, potassium hydroxide and ammonium hydroxide are monoacidic.
\[\mathrm{NaOH} \rightarrow \mathrm{Na^{+}} + \mathrm{OH^{-}}\]
A diacidic base dissociates in the molten state, or in aqueous solution, to produce two hydroxide ions per molecule of the base. Calcium hydroxide and copper hydroxide are diacidic.
\[\mathrm{Cu(OH)_2}\text{(fused)} \rightleftharpoons \mathrm{Cu^{2+}} + 2\,\mathrm{OH^{-}}\]
A triacidic base dissociates in aqueous solution to produce three hydroxide ions per molecule of the base, or equivalently contains three replaceable hydroxyl ions in one molecule. Aluminium hydroxide, \(\mathrm{Al(OH)_3}\), and ferric hydroxide, \(\mathrm{Fe(OH)_3}\), are triacidic.
Table 3.8 places the two counting ideas together, because examination questions very often ask for exactly this distinction.
3.10 Preparation Of Bases
Seven routes lead to a base, and they fall into two groups. In the first, a metal or a metal compound is oxidised or broken down to leave the oxide behind. In the second, water or an alkali is used to bring the hydroxide out directly. FC 3.1 sets out the choices.
Flowchart
FC 3.1 — Routes To A Base Start from a metal → burn in oxygen → basic oxide; or, if the metal is reactive (Na, K, Ca), add water → alkali + hydrogen Start from a metal compound → soluble oxide + water → alkali; or metal sulphide + oxygen, heated → basic oxide + sulphur dioxide; or carbonate or nitrate, heated → basic oxide; or salt solution + alkali → insoluble hydroxide precipitated Start from ammonia gas → dissolve in water → ammonium hydroxide
From The Metal Itself
Metals react with oxygen to give bases. Burn the metal in air and the oxide is left as the product.
sodium + oxygen → sodium oxide
\[4\,\mathrm{Na} + \mathrm{O_2} \rightarrow 2\,\mathrm{Na_2O}\]
\[2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO}\]
The most reactive metals — sodium, potassium and calcium — do not need oxygen gas at all. They take oxygen from water, and the base formed is the hydroxide, dissolved and ready.
sodium + water → sodium hydroxide + hydrogen
\[2\,\mathrm{Na} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH} + \mathrm{H_2}\text{(g)}\]
\[2\,\mathrm{K} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{KOH} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{Ca} + 2\,\mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2} + \mathrm{H_2}\text{(g)}\]
These reactions are violent, and the hydrogen released can catch fire from the heat of the reaction itself. They are demonstrated by a teacher behind a safety screen, using a very small piece of metal, and never carried out by a student.
From An Oxide
Where the oxide of a metal is soluble, adding water converts it straight into the alkali. This works for the oxides of sodium, potassium and calcium.
\[\mathrm{Na_2O} + \mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH}\]
\[\mathrm{K_2O} + \mathrm{H_2O} \rightarrow 2\,\mathrm{KOH}\]
\[\mathrm{CaO} + \mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2}\]
This is the exact counterpart of Section 3.4.2, where the oxide of a non-metal plus water gave an acid. A basic oxide plus water gives an alkali. The oxide decides which way the solution will go.
By Double Decomposition
An insoluble hydroxide cannot be made by dissolving anything, so it is thrown out of solution instead. Add an alkali to the aqueous solution of a salt of the metal, and the metallic hydroxide precipitates while the other salt stays dissolved.
iron(III) chloride + sodium hydroxide → iron(III) hydroxide + sodium chloride
\[\mathrm{FeCl_3}\text{(aq)} + 3\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_3}\!\downarrow + 3\,\mathrm{NaCl}\text{(aq)}\]
\[\mathrm{CuSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Cu(OH)_2}\!\downarrow + \mathrm{Na_2SO_4}\text{(aq)}\]
\[\mathrm{AlCl_3}\text{(aq)} + 3\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{Al(OH)_3}\!\downarrow + 3\,\mathrm{NH_4Cl}\text{(aq)}\]
The two solutions simply exchange partners, and the pair that cannot stay in solution falls out as a solid.
From A Sulphide Or A Salt, By Heating
Heating a metal sulphide in oxygen converts it to the oxide, driving off sulphur as sulphur dioxide.
\[2\,\mathrm{ZnS} + 3\,\mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 2\,\mathrm{SO_2}\]
\[2\,\mathrm{PbS} + 3\,\mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO} + 2\,\mathrm{SO_2}\]
Certain salts break down on heating and leave the basic oxide behind. A metal carbonate gives the oxide and carbon dioxide.
\[\mathrm{CaCO_3} \xrightarrow{\ \Delta\ } \mathrm{CaO} + \mathrm{CO_2}\text{(g)}\]
\[\mathrm{CuCO_3} \xrightarrow{\ \Delta\ } \mathrm{CuO} + \mathrm{CO_2}\text{(g)}\]
The carbonates of sodium and of potassium are the exceptions: they do not decompose on heating.
A metal nitrate gives the oxide, nitrogen dioxide and oxygen.
\[2\,\mathrm{Ca(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{CaO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
\[2\,\mathrm{Zn(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Sodium and potassium nitrates are again the exceptions, and do not give the metal oxide when decomposed.
Ammonium Hydroxide
The last route needs no metal at all. Ammonium hydroxide is prepared by simply dissolving ammonia gas in water, in which ammonia is extremely soluble.
\[\mathrm{NH_3}\text{(g)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{NH_4OH}\text{(aq)}\]
The product is a weak alkali, for the reason given in Section 3.9.1 — only a small fraction of it dissociates to give hydroxide ions.
3.11 Properties Of Bases And Alkalis
3.11.1 Physical Properties
Taste. Bases have a sharp and bitter taste. Like the sourness of acids, this is a property of the class and never a test.
Action on indicators. Alkalis change the colour of indicators, and each indicator changes in its own way — litmus from red to blue, methyl orange from orange to yellow, and phenolphthalein from colourless to pink. Insoluble bases do not affect indicators at all, because an indicator can only respond to ions in solution and an insoluble base puts none there.
Feel. Alkalis are soapy substances; they are slippery to touch. The slipperiness is not an accident of texture but a reaction taking place on your fingers. Alkalis react with oils and fats to form soap, and skin carries a natural film of oil, so a drop of caustic soda or caustic potash on the hand begins converting that film into a soapy solution. That is also why an alkali burn feels smooth rather than sharp, and why it must be washed off at once under running water: the reaction does not stop at the oil film.
Conduction. Alkalis are strong electrolytes and conduct electricity in the aqueous or molten state, since in both states their ions are free to move.
Effect on skin. Alkalis show a mild corrosive action, giving a slight burn on the skin. Strong alkalis are handled with gloves, and eye protection is worn whenever their solutions are poured.
3.11.2 Chemical Properties
Absorption Of Carbon Dioxide
Strong alkalis absorb carbon dioxide from the air to form carbonates. This is why a bottle of caustic soda left open loses its strength — the surface slowly turns to sodium carbonate.
sodium hydroxide + carbon dioxide → sodium carbonate + water
\[2\,\mathrm{NaOH} + \mathrm{CO_2} \rightarrow \mathrm{Na_2CO_3} + \mathrm{H_2O}\]
\[2\,\mathrm{KOH} + \mathrm{CO_2} \rightarrow \mathrm{K_2CO_3} + \mathrm{H_2O}\]
Neutralisation Of Acids
Bases and alkalis neutralise acids to form salt and water. Both a soluble and an insoluble base do this, which is the defining reaction of the whole class.
\[\mathrm{Ca(OH)_2} + 2\,\mathrm{HCl} \rightarrow \mathrm{CaCl_2} + 2\,\mathrm{H_2O}\]
\[\mathrm{Fe(OH)_2} + 2\,\mathrm{HCl} \rightarrow \mathrm{FeCl_2} + 2\,\mathrm{H_2O}\]
Precipitation Of Insoluble Hydroxides
Added to a solution of a salt of a heavy metal — copper, iron, zinc and the like — an alkali precipitates the insoluble hydroxide of that metal. You met this reaction in Section 3.10 as a way of making a base. Read it now from the other side: the colour of the precipitate tells you which metal was in the solution.
\[\mathrm{CuSO_4}\text{(aq)} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{(NH_4)_2SO_4}\text{(aq)} + \mathrm{Cu(OH)_2}\!\downarrow\]
\[\mathrm{ZnSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + \mathrm{Zn(OH)_2}\!\downarrow\]
\[\mathrm{FeSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2SO_4}\text{(aq)} + \mathrm{Fe(OH)_2}\!\downarrow\]
| Salt solution | Alkali added | Hydroxide precipitated | Colour of the precipitate |
|---|---|---|---|
| Copper(II) sulphate | Ammonium hydroxide | \(\mathrm{Cu(OH)_2}\) | Pale blue |
| Zinc sulphate | Sodium hydroxide | \(\mathrm{Zn(OH)_2}\) | Gelatinous white |
| Iron(II) sulphate | Sodium hydroxide | \(\mathrm{Fe(OH)_2}\) | Dirty green |
Table 3.9 Colours of some hydroxides precipitated by an alkali
Now add more alkali and watch. Most precipitates simply sit there, but the hydroxides of zinc, aluminium and lead dissolve again in excess sodium hydroxide or potassium hydroxide. These three are amphoteric — they behave as a base towards an acid and as an acid towards a strong alkali, and so react with both.
\[\mathrm{Zn(OH)_2} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + 2\,\mathrm{H_2O}\]
The product is sodium zincate, and it is soluble, which is why the white precipitate disappears. Other hydroxides do not behave in this way, and that difference is one of the sharpest tests in the laboratory.
Action On Ammonium Salts
When alkalis are warmed with an ammonium salt, ammonia gas is given out. The alkali takes the hydrogen from the ammonium ion, and what is left leaves as a gas.
ammonium chloride + sodium hydroxide → sodium chloride + water + ammonia
\[\mathrm{NH_4Cl} + \mathrm{NaOH} \xrightarrow{\ \Delta\ } \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{NH_3}\text{(g)}\]
\[2\,\mathrm{NH_4Cl} + \mathrm{Ca(OH)_2} \xrightarrow{\ \Delta\ } \mathrm{CaCl_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{NH_3}\text{(g)}\]
\[\mathrm{(NH_4)_2SO_4} + 2\,\mathrm{KOH} \xrightarrow{\ \Delta\ } \mathrm{K_2SO_4} + 2\,\mathrm{H_2O} + 2\,\mathrm{NH_3}\text{(g)}\]
Ammonia is recognised at once by its pungent smell and by the dense white fumes it gives when a rod moistened with concentrated hydrochloric acid is held at the mouth of the tube. The gas is irritating to the eyes and throat, so it is never smelt directly; the vapour is wafted gently towards the nose with the hand.
3.12 Uses Of Some Common Bases
Every use in Table 3.10 follows from a property established earlier in this chapter. Alkalis attack oils and fats, so they make soap. Bases neutralise acids, so they treat acidity. Some absorb water and carbon dioxide from the air, so they dry other substances.
| Base | Use |
|---|---|
| Sodium hydroxide | Manufacture of soap |
| Potassium hydroxide | Manufacture of salts and soaps; used in alkaline batteries |
| Calcium hydroxide (slaked lime) | Manufacture of bleaching powder; softening of hard water |
| Magnesium hydroxide | As an antacid, to neutralise acidity |
| Aluminium hydroxide | As a foaming agent in fire extinguishers |
| Ammonium hydroxide | To remove grease stains from clothes |
| Magnesia, \(\mathrm{MgO}\) | In making refractory bricks |
| Calcium oxide | Drying agent; basic flux |
Table 3.10 General uses of some bases
Two entries need a word of explanation. A refractory brick is one that can stand very high temperatures without melting or cracking, which is what the lining of a furnace must do; magnesium oxide melts only at an extremely high temperature and serves well. A flux is a substance added to a furnace charge to combine with the impurities and carry them off as a fusible slag, and a basic flux such as calcium oxide is chosen when the impurity to be removed is acidic.
Notice how sodium hydroxide and ammonium hydroxide do the same kind of job at very different strengths. The strong alkali converts fat into soap outright in a factory; the weak one is mild enough to loosen a grease stain from cloth at home without destroying the fabric.
3.13 The pH Scale
3.13.1 pH And Hydrogen Ion Concentration
Litmus, methyl orange and phenolphthalein will tell you whether a solution is acidic, basic or neutral. What they will not tell you is how acidic it is. A drop of lemon juice and a drop of concentrated hydrochloric acid both turn blue litmus red, yet one is safe to swallow and the other would destroy the tissue of your mouth.
To measure the difference, a number is needed. In 1909 the Danish biochemist Sorensen devised a scale for exactly this — the pH scale, which represents the hydronium ion concentration of a given aqueous solution.
\[\mathrm{pH} = \log_{10}\frac{1}{[\mathrm{H^{+}}]} = -\log_{10}[\mathrm{H^{+}}]\]
Here \([\mathrm{H^{+}}]\) is the concentration of hydrogen ions in the solution, measured in moles per litre. The two forms are the same statement written two ways, since taking the reciprocal of a number changes the sign of its logarithm.
Where does the scale’s range come from? Water itself ionises, though only very slightly. In pure water the concentration of \(\mathrm{H_3O^{+}}\) ions and of \(\mathrm{OH^{-}}\) ions is each found to be \(1\times10^{-7}\ \text{mol litre}^{-1}\), so every aqueous solution contains both. The product of these two concentrations is
\[(1\times10^{-7}) \times (1\times10^{-7}) = 1\times10^{-14}\ \text{mol}^2\ \text{litre}^{-2} \text{ at } 25^\circ\text{C}\]
and this product is called the ionic product of water, \(K_w\). Add acid to water and the concentration of \(\mathrm{H_3O^{+}}\) rises while that of \(\mathrm{OH^{-}}\) falls, so that \(K_w\) stays the same; add an alkali and the reverse happens. The nature of a solution therefore depends entirely on the concentration of \(\mathrm{H_3O^{+}}\) ions in it — which is what pH measures.
The normal pH scale runs from 0 to 14, as Fig. 3.13-1 shows.
Diagram
Fig. 3.13-1 — The pH Scale A horizontal number line marked 0 to 14 at equal intervals, with a vertical arrow at 7 labelled NEUTRAL. An arrow running leftwards from 7 towards 0 labelled INCREASING ACIDIC NATURE, and an arrow running rightwards from 7 towards 14 labelled INCREASING ALKALINE NATURE. Mark the numbers clearly and label the whole line ‘pH’.
Three statements sum the scale up.
- The number 7 represents the neutrality of the solution.
- Numbers less than 7 — pH 6, 5, 4 and so on — indicate that the acidic nature is increasing as the numbers decrease.
- Numbers greater than 7 — pH 8, 9, 10 and so on — show that alkalinity increases as the numbers increase.
Because pH is built from a logarithm, each step of one unit means a tenfold change in the hydrogen ion concentration. A solution of pH 3 is not a little more acidic than one of pH 5; it is a hundred times more acidic.
3.13.2 Universal Indicator And pH Paper
For a colourless liquid, a reasonably accurate pH value can be obtained by adding a universal indicator. It is a mixture of indicator dyes which together give a spectrum of colours, each colour belonging to a different concentration of hydrogen ions.
In a neutral solution, at pH 7, a universal indicator produces a green colour. As the solution becomes more basic and the pH rises from 7 to 14, the colour changes progressively from blue to indigo to violet. As the solution becomes more acidic and the pH falls from 7 to 1, the colour changes from yellow to pink, and then to red. Fig. 3.13-2 sets the colours against the scale.
Diagram
Fig. 3.13-2 — Colours Of The Universal Indicator A horizontal band divided into coloured segments, running left to right, placed above a pH axis numbered 0 to 14 and below a second axis showing hydrogen ion concentration from 10⁻¹ to 10⁻¹⁴ moles per litre. Colour the band red at pH 1, pink at pH 3, yellow at pH 5, green at pH 7, blue at pH 9, indigo at pH 11 and violet at pH 13, blending between them. Label the left half ACIDIC NATURE INCREASES with a leftward arrow and the right half BASIC NATURE INCREASES with a rightward arrow.
A pH paper is a piece of paper used to find out whether a solution is acidic, basic or neutral. A drop of the liquid is placed on the paper and the colour that appears is matched against a printed chart.
One warning belongs here, and it catches students out regularly. The presence of an H atom or an OH group in a formula does not tell you whether a substance is an acid or a base. Ethanol, \(\mathrm{C_2H_5OH}\), contains both, and is neither. What decides the matter is the ability of the substance to release \(\mathrm{H^{+}}\) or \(\mathrm{OH^{-}}\) ions in aqueous solution, and pH measures the result.
3.13.3 Importance Of pH In Everyday Life
In the body. Our body works within a narrow pH range of 7.0 to 7.8, and plants and animals generally survive only within narrow ranges of their own. A shift of a few tenths of a unit in blood pH is a serious medical event.
In agriculture. Every crop grows better in a particular pH range. Rice grows better in slightly acidic soil, sugarcane in neutral soil, and citrus fruits in alkaline soil. A farmer who knows the pH of a field knows which crop it will favour, and what must be added to shift it.
Acid rain. If the pH of rain water is less than 5.6, the rain is said to be acid rain. Such water contains acids like sulphuric acid, nitric acid and carbonic acid, formed from atmospheric pollutants — mainly sulphur dioxide and oxides of nitrogen. Acid rain reduces the fertility of soil, destroys marine life, and causes damage to monuments.
In medicine. Certain diseases are diagnosed on the basis of the pH value of blood and of urine, since a body that has stopped regulating its acidity leaves that record in both.
In the digestive system. Hydrochloric acid is produced in the stomach, where it helps in the digestion of food. If it becomes excessive the pH falls, and pain and irritation follow. To get rid of this, antacids such as milk of magnesia, \(\mathrm{Mg(OH)_2}\), are generally used to adjust the pH. An antacid is chosen to be a weak, insoluble base — strong enough to neutralise the excess acid, mild enough to be swallowed.
In preventing tooth decay. Substances like chocolates and sweets are degraded by bacteria present in the mouth, and the acids produced lower the pH. When the pH falls to 5.5, tooth decay starts. Tooth enamel, which is calcium phosphate, is the hardest substance in our body, and even that gets corroded. Saliva produced by the salivary glands is slightly alkaline and helps to raise the pH to some extent, but toothpaste is used to neutralise the excess acid in the mouth.
A bee sting. A bee sting leaves acid in the body. If baking soda, which is a base, is applied on the stung area, it neutralises the acid and gives relief.
3.14 Salts
3.14.1 What A Salt Is
Everything in this chapter so far has been leading here. An acid carries replaceable hydrogen; a base carries a metal or the ammonium group ready to take its place. Let the two meet, and the compound left behind is a salt.
The words partial and total are both in the definition because both are possible, and which one you get is under your control. Section 3.6.4 showed sodium nitrate warmed with concentrated sulphuric acid, and the temperature decided the outcome. Look at the same two reactions again, this time asking how many hydrogen atoms of the sulphuric acid were replaced.
Partial replacement — one of the two hydrogen atoms is replaced:
\[\mathrm{NaNO_3} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ <200^\circ\text{C}\ } \mathrm{NaHSO_4} + \mathrm{HNO_3}\]
Complete replacement — both hydrogen atoms are replaced:
\[2\,\mathrm{NaNO_3} + \mathrm{H_2SO_4}\text{(conc.)} \xrightarrow{\ >200^\circ\text{C}\ } \mathrm{Na_2SO_4} + 2\,\mathrm{HNO_3}\]
Sodium hydrogen sulphate still holds one replaceable hydrogen; sodium sulphate holds none. Both are salts.
Diagram
Fig. 3.14-1 — Replacement Of Hydrogen By A Metal Three boxes in a row joined by arrows. First box: the formula H₂SO₄ with both hydrogen atoms circled and labelled ‘ionisable hydrogen’. Second box: NaHSO₄, with one circled hydrogen replaced by Na and labelled ‘partial replacement — acid salt’. Third box: Na₂SO₄, with both replaced by Na and labelled ‘total replacement — normal salt’. Below the row, a caption strip reading ‘the metal takes the place of the acid’s hydrogen’.
3.14.2 The Ionic Definition Of A Salt
The definition above describes how a salt is made. A second definition describes what a salt is once it is dissolved, and this one is often the more useful of the two.
Sodium chloride is the plainest example. It consists of \(\mathrm{Na^{+}}\) and \(\mathrm{Cl^{-}}\) ions, and in water these separate.
\[\mathrm{NaCl} \xrightarrow{\ \text{water}\ } \mathrm{Na^{+}} + \mathrm{Cl^{-}}\]
The sodium ion is positive, but it is not a hydrogen ion; the chloride ion is negative, but it is not a hydroxyl ion. That is exactly what makes the compound a salt and not an acid or a base. An acid would have supplied \(\mathrm{H^{+}}\); an alkali would have supplied \(\mathrm{OH^{-}}\); a salt supplies neither.
It is prepared by neutralising sodium hydroxide with hydrochloric acid.
\[\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}\]
3.15 Classification Of Salts
There are six kinds of salts: normal salts, acid salts, basic salts, double salts, mixed salts and complex salts. The first three are told apart by how much replacement has taken place; the last three by what the salt breaks into when it dissolves.
3.15.1 Normal Salts
Normal salts are the salts formed by the complete replacement of the ionisable hydrogen atoms of an acid by a metallic or an ammonium ion. Hydrochloric acid yields sodium chloride, and sulphuric acid gives sodium sulphate.
\[\mathrm{HCl} + \mathrm{NaOH} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}\]
\[\mathrm{H_2SO_4} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2SO_4} + 2\,\mathrm{H_2O}\]
The test of a normal salt is simple: it has no ionisable hydrogen atoms left. Sodium carbonate, \(\mathrm{Na_2CO_3}\), sodium phosphate, \(\mathrm{Na_3PO_4}\), and sodium acetate, \(\mathrm{CH_3COONa}\), are all normal salts.
3.15.2 Acid Salts
Acid salts are formed by the partial replacement of the ionisable hydrogen atoms of a polybasic acid by a metal or an ammonium ion. Only a polybasic acid can give one, for the plain reason that a monobasic acid has a single hydrogen and cannot part with some of it.
Dibasic sulphuric acid yields sodium hydrogen sulphate by either of two routes.
\[\mathrm{H_2SO_4} + \mathrm{NaCl} \xrightarrow{\ \text{below }200^\circ\text{C}\ } \mathrm{NaHSO_4} + \mathrm{HCl}\]
\[\mathrm{H_2SO_4} + \mathrm{NaOH} \rightarrow \mathrm{NaHSO_4} + \mathrm{H_2O}\]
Acid salts are usually formed when the amount of base taken is insufficient for the complete neutralisation of the acid. Sodium hydrogen sulphite, \(\mathrm{NaHSO_3}\), disodium hydrogen phosphate, \(\mathrm{Na_2HPO_4}\), and sodium dihydrogen phosphate, \(\mathrm{NaH_2PO_4}\), are further examples.
Because an acid salt still holds replaceable hydrogen, it ionises in water solution to give hydronium ions, and therefore shows all the properties of an acid.
3.15.3 Basic Salts
Basic salts are formed by the partial replacement of the hydroxyl group of a di-acidic or a tri-acidic base by an acid radical. This is the mirror image of an acid salt: there, part of the acid’s hydrogen survived; here, part of the base’s hydroxyl survives.
\[\mathrm{Pb(OH)_2} + \mathrm{HCl} \rightarrow \mathrm{Pb(OH)Cl} + \mathrm{H_2O}\]
\[\mathrm{Mg(OH)_2} + \mathrm{HCl} \rightarrow \mathrm{Mg(OH)Cl} + \mathrm{H_2O}\]
The products are basic lead chloride and basic magnesium chloride. Basic copper chloride, \(\mathrm{Cu(OH)Cl}\), and basic copper nitrate, \(\mathrm{Cu(OH)NO_3}\), belong to the same family.
Look at any of these formulae and you will find three parts: a metallic cation, a hydroxyl ion left over from the base, and an anion from the acid. That is what makes a salt basic.
3.15.4 Complex Salts
Complex salts are those which on dissociation give one simple ion and one complex ion. A complex ion is a group of atoms bound together that stays intact in solution and carries a charge as a unit, and it is always written inside square brackets.
\[\mathrm{Na[Ag(CN)_2]} \rightleftharpoons \mathrm{Na^{+}} + \mathrm{[Ag(CN)_2]^{-}}\]
This salt is sodium argentocyanide. Notice that the silver does not appear as a free \(\mathrm{Ag^{+}}\) ion anywhere in the solution; it stays locked inside the complex ion. Other complex salts are silver amino chloride, \(\mathrm{[Ag(NH_3)_2]Cl}\), tetraammine copper(II) sulphate, \(\mathrm{[Cu(NH_3)_4]SO_4}\), potassium mercuric iodide, \(\mathrm{K_2[HgI_4]}\), and sodium zincate, \(\mathrm{Na_2ZnO_2}\) — the last of which you have already made, in Section 3.11.2, by dissolving zinc hydroxide in excess alkali.
3.15.5 Double Salts And Mixed Salts
Double salts are formed by the union of two simple salts that dissolve in water and crystallise together. Potash alum, \(\mathrm{K_2SO_4\!\cdot\!Al_2(SO_4)_3\!\cdot\!24H_2O}\), Mohr’s salt, \(\mathrm{FeSO_4\!\cdot\!(NH_4)_2SO_4\!\cdot\!6H_2O}\), and dolomite, \(\mathrm{CaCO_3\!\cdot\!MgCO_3}\), are double salts.
Mixed salts are those salts that contain more than one basic or acid radical. Sodium potassium carbonate, \(\mathrm{NaKCO_3}\), has two different basic radicals; bleaching powder, \(\mathrm{CaOCl_2}\), has two different acid radicals.
Table 3.11 gathers all six.
3.16 Methods Of Preparation Of Soluble Salts
A soluble salt is always prepared in solution. Whatever route you take, the salt ends up dissolved in water, and it is recovered by evaporating the water and letting the salt crystallise out. Fig. 3.16-1 shows the finishing steps, which are the same for every method in this section.
Diagram
Fig. 3.16-1 — Recovering A Soluble Salt From Its Solution Three stages left to right, joined by arrows. Stage 1: a funnel with filter paper resting in a conical flask, separating unreacted solid from the salt solution; label ‘filtration’. Stage 2: an evaporating dish on a wire gauze and tripod over a burner, the solution being heated to the point of crystallisation; label ‘evaporation to the crystallisation point’. Stage 3: the same dish set aside to cool, with crystals forming in it; label ‘cooling and crystallisation’. Label the filter paper, the residue, the filtrate, the tripod and the crystals.
One warning goes with this. Some soluble salts, such as the zincates and the aluminates, exist only in solution and cannot be obtained in pure form by crystallisation at all.
3.16.1 Direct Combination
Two elements are heated together and the salt forms directly. A metal combines with a non-metal, and nothing else is produced.
sodium + chlorine → sodium chloride
\[2\,\mathrm{Na}\text{(molten)} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{NaCl}\]
\[2\,\mathrm{Fe} + 3\,\mathrm{Cl_2} \rightarrow 2\,\mathrm{FeCl_3}\]
\[2\,\mathrm{Al} + 3\,\mathrm{Cl_2} \rightarrow 2\,\mathrm{AlCl_3}\]
This method is reserved for salts in which the metal shows its higher valency, such as iron(III) chloride, because chlorine is an oxidising agent and drives the metal to that state.
3.16.2 Simple Displacement
A dilute acid is poured on an active metal. The metal displaces the hydrogen of the acid and takes its place, so a salt of the metal is left in solution while hydrogen escapes.
\[\mathrm{Zn} + \mathrm{H_2SO_4}\text{(dil.)} \rightarrow \mathrm{ZnSO_4} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{Fe} + 2\,\mathrm{HCl}\text{(dil.)} \rightarrow \mathrm{FeCl_2} + \mathrm{H_2}\text{(g)}\]
Only the soluble salts of active metals can be made this way, since a metal below hydrogen in reactivity will not displace it. Compare the two iron salts and the point of Section 3.16.1 becomes clear: chlorine gives iron(III) chloride, while dilute hydrochloric acid gives iron(II) chloride.
3.16.3 Decomposition Of Salts By Dilute Acids
The reactions themselves were set out in Section 3.6, where the interest was in the gas evolved. Here the interest is in what stays behind in the beaker, which is a new salt in solution.
Carbonates and hydrogen carbonates give the salt, water and carbon dioxide.
\[\mathrm{NaHCO_3} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{CO_2}\text{(g)}\]
\[\mathrm{Ca(HCO_3)_2} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Ca(NO_3)_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{CO_2}\]
\[\mathrm{CuCO_3} + 2\,\mathrm{HCl} \rightarrow \mathrm{CuCl_2} + \mathrm{H_2O} + \mathrm{CO_2}\text{(g)}\]
\[\mathrm{MgCO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{MgSO_4} + \mathrm{H_2O} + \mathrm{CO_2}\text{(g)}\]
Sulphites and hydrogen sulphites give the salt, water and sulphur dioxide.
\[\mathrm{Na_2SO_3} + 2\,\mathrm{HCl} \rightarrow 2\,\mathrm{NaCl} + \mathrm{H_2O} + \mathrm{SO_2}\]
\[\mathrm{CuSO_3} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Cu(NO_3)_2} + \mathrm{H_2O} + \mathrm{SO_2}\]
Sulphides give the salt and hydrogen sulphide.
\[\mathrm{K_2S} + \mathrm{H_2SO_4} \rightarrow \mathrm{K_2SO_4} + \mathrm{H_2S}\]
\[\mathrm{ZnS} + \mathrm{H_2SO_4} \rightarrow \mathrm{ZnSO_4} + \mathrm{H_2S}\]
The gas leaves the vessel of its own accord, which is convenient: it takes itself out of the mixture and leaves the salt solution clean.
3.16.4 Neutralisation Of An Alkali By Titration
An acid and a soluble base are mixed until each has exactly used up the other. Because both the reactants and the products are soluble, there is nothing to filter off, and the only way of knowing when to stop is to measure — which is what a titration does.
\[\mathrm{HNO_3} + \mathrm{NaOH} \rightarrow \mathrm{NaNO_3} + \mathrm{H_2O}\]
\[\mathrm{HCl} + \mathrm{NH_4OH} \rightarrow \mathrm{NH_4Cl} + \mathrm{H_2O}\]
The procedure itself, with the burette and the indicator, is described in Section 3.19.7.
3.16.5 Neutralisation Of An Insoluble Base
Here the base will not dissolve, and that turns out to be an advantage. Warm dilute acid is added to the insoluble oxide, hydroxide or carbonate a little at a time, until no more of it dissolves and the excess settles at the bottom. There is then no acid left over, and the excess base is simply filtered off.
\[2\,\mathrm{HCl} + \mathrm{MgO} \rightarrow \mathrm{MgCl_2} + \mathrm{H_2O}\]
\[2\,\mathrm{HNO_3} + \mathrm{PbO} \rightarrow \mathrm{Pb(NO_3)_2} + \mathrm{H_2O}\]
\[\mathrm{H_2SO_4} + \mathrm{Cu(OH)_2} \rightarrow \mathrm{CuSO_4} + 2\,\mathrm{H_2O}\]
\[2\,\mathrm{HNO_3} + \mathrm{Pb(OH)_2} \rightarrow \mathrm{Pb(NO_3)_2} + 2\,\mathrm{H_2O}\]
No titration is needed, and no indicator either. The excess solid is its own signal that the acid has been used up.
3.16.6 Action Of Alkalis On Certain Metals, Oxides And Hydroxides
Aluminium and zinc are the awkward pair of this chapter. Their oxides and hydroxides are amphoteric, and the metals themselves dissolve in a strong alkali as well as in an acid, giving a salt in each case. With sodium hydroxide or potassium hydroxide the salts formed are the aluminates and the zincates.
\[2\,\mathrm{Al} + 2\,\mathrm{NaOH} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaAlO_2} + 3\,\mathrm{H_2}\]
\[\mathrm{Al_2O_3} + 2\,\mathrm{NaOH} \rightarrow 2\,\mathrm{NaAlO_2} + \mathrm{H_2O}\]
\[\mathrm{Al(OH)_3} + \mathrm{NaOH} \rightarrow \mathrm{NaAlO_2} + 2\,\mathrm{H_2O}\]
\[\mathrm{Zn} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + \mathrm{H_2}\]
\[\mathrm{ZnO} + 2\,\mathrm{KOH} \rightarrow \mathrm{K_2ZnO_2} + \mathrm{H_2O}\]
\[\mathrm{Zn(OH)_2} + 2\,\mathrm{KOH} \rightarrow \mathrm{K_2ZnO_2} + 2\,\mathrm{H_2O}\]
Note that the metal releases hydrogen while the oxide and the hydroxide release only water. These are the salts mentioned above that exist in solution alone and cannot be crystallised.
FC 3.2 puts the six methods into the order in which you would actually think about them.
Flowchart
FC 3.2 — Choosing A Method For A Soluble Salt What is available? → two elements only → heat them together, direct combination → an active metal + dilute acid → simple displacement → a carbonate, sulphite or sulphide + dilute acid → decomposition, gas escapes → an insoluble oxide, hydroxide or carbonate + warm dilute acid → add until excess remains, filter, no indicator needed → a soluble base (alkali) + acid → titration, indicator needed → aluminium or zinc, or their oxides and hydroxides, + strong alkali → aluminate or zincate, solution only → in every case except the last: evaporate to the crystallisation point → cool → crystals
3.17 Solubility Of Salts
Before a salt can be prepared, one question has to be settled: will it dissolve? A soluble salt is built in solution and crystallised out of it; an insoluble salt has to be thrown down as a precipitate. The whole of Section 3.18 depends on knowing which is which, and Table 3.12 is the reference to learn.
| Salts | Exceptions |
|---|---|
| All the compounds of ammonium, sodium and potassium are soluble in water. | — |
| All nitrates and nitrites are soluble. | — |
| All chlorides, bromides and iodides are soluble in water. | \(\mathrm{Hg_2Cl_2}\); \(\mathrm{AgCl}\); \(\mathrm{PbCl_2}\), which is soluble in hot water but not in cold water. |
| All sulphates are soluble. | \(\mathrm{CaSO_4}\); \(\mathrm{PbSO_4}\); \(\mathrm{Ag_2SO_4}\); \(\mathrm{BaSO_4}\). |
| All carbonates, sulphides, sulphites and phosphates are insoluble. | Those of ammonium, sodium and potassium. |
Table 3.12 Solubility of salts in water
Three further points complete the picture. All metallic oxides and hydroxides are insoluble, except those of sodium and potassium, and ammonium hydroxide; calcium hydroxide is slightly soluble. Sodium hydrogen carbonate and potassium hydrogen carbonate are only sparingly soluble, which is what makes the preparation in Section 3.19.4 possible at all.
Read the table as a set of rules with named exceptions, not as a list to be memorised blindly. The first line alone settles a great many questions: if the salt contains ammonium, sodium or potassium, it dissolves, and there are no exceptions to worry about.
3.18 Preparation Of Insoluble Salts
An insoluble salt cannot be crystallised out of a solution, because it was never dissolved in one. It has to be made as a solid from the start — either by joining two substances directly, or by assembling it inside a solution and letting it fall out as a precipitate.
3.18.1 By Direct Combination
A metal and a non-metal are heated together and the insoluble salt is formed at once.
lead + sulphur → lead sulphide
\[\mathrm{Pb} + \mathrm{S} \xrightarrow{\ \Delta\ } \mathrm{PbS}\]
\[\mathrm{Fe} + \mathrm{S} \xrightarrow{\ \Delta\ } \mathrm{FeS}\]
3.18.2 By Combining An Acidic Oxide With A Basic Oxide
An acidic oxide and a basic oxide, heated together, give the salt directly and no water at all. This is neutralisation stripped down to its bare bones — the acid without its water meeting the base without its water.
\[\mathrm{SO_2} + \mathrm{CaO} \xrightarrow{\ \Delta\ } \mathrm{CaSO_3}\]
\[\mathrm{CO_2} + \mathrm{CaO} \xrightarrow{\ \Delta\ } \mathrm{CaCO_3}\]
3.18.3 By Precipitation Or Double Decomposition
This is the method that matters most. Two soluble salts are dissolved separately, and the solutions are mixed. The ions change partners, and if one of the new pairs cannot stay in solution, it drops to the bottom of the beaker as a solid.
\[\mathrm{BaCl_2} + \mathrm{H_2SO_4} \rightarrow \mathrm{BaSO_4}\!\downarrow + 2\,\mathrm{HCl}\]
\[\mathrm{AgNO_3} + \mathrm{HCl} \rightarrow \mathrm{AgCl}\!\downarrow + \mathrm{HNO_3}\]
\[\mathrm{Pb(NO_3)_2} + \mathrm{Na_2CO_3} \rightarrow \mathrm{PbCO_3}\!\downarrow + 2\,\mathrm{NaNO_3}\]
\[\mathrm{Na_2SO_4} + \mathrm{BaCl_2} \rightarrow \mathrm{BaSO_4}\!\downarrow + 2\,\mathrm{NaCl}\]
\[\mathrm{CuSO_4} + \mathrm{H_2S} \rightarrow \mathrm{CuS}\!\downarrow + \mathrm{H_2SO_4}\]
\[\mathrm{Pb(NO_3)_2} + 2\,\mathrm{HCl} \rightarrow \mathrm{PbCl_2}\!\downarrow + 2\,\mathrm{HNO_3}\]
The precipitate is filtered off, washed with distilled water to remove the soluble salt clinging to it, and dried.
Silver Chloride And Barium Sulphate
Silver chloride is usually prepared by the reaction of aqueous silver nitrate with dilute hydrochloric acid, and a precipitate of silver chloride is obtained. Barium sulphate is prepared by reacting aqueous barium chloride with dilute sulphuric acid, which gives a white precipitate.
Carbonates
Most of the carbonates are insoluble, and they can all be prepared by the same trick: boil a solution of the metal salt with sodium carbonate solution. The carbonate ion of the sodium salt attaches itself to the other metal ion, leaving the sodium ion behind in solution, so the metal carbonate is precipitated.
\[\mathrm{ZnSO_4} + \mathrm{Na_2CO_3} \rightarrow \mathrm{Na_2SO_4} + \mathrm{ZnCO_3}\!\downarrow\]
\[\mathrm{BaCl_2} + \mathrm{Na_2CO_3} \rightarrow 2\,\mathrm{NaCl} + \mathrm{BaCO_3}\!\downarrow\]
3.18.4 An Insoluble Salt From Another Insoluble Salt
Suppose the substance you have is itself insoluble. Nothing can be exchanged, because nothing is in solution. The way round it is to go through a soluble stage: the insoluble salt is first converted into a soluble salt, and that solution is then used to prepare the salt you want. FC 3.3 shows the route.
Flowchart
FC 3.3 — Making One Insoluble Salt From Another Insoluble lead carbonate or insoluble lead oxide → add dilute nitric acid → soluble lead nitrate in solution → add dilute sulphuric acid or sodium sulphate solution → white precipitate of lead sulphate → filter, wash, dry; the same lead nitrate solution, treated instead with hydrochloric acid or sodium chloride solution → precipitate of lead chloride
Dilute nitric acid is the reagent for the first step, and the reason lies in Table 3.12: all nitrates are soluble, so whatever lead compound you start from, lead nitrate will dissolve.
\[\mathrm{PbO} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Pb(NO_3)_2} + \mathrm{H_2O}\]
\[\mathrm{PbCO_3} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Pb(NO_3)_2} + \mathrm{H_2O} + \mathrm{CO_2}\text{(g)}\]
The resulting solution is then treated with sulphuric acid or with sodium sulphate solution, and a white precipitate of lead sulphate is obtained.
\[\mathrm{Pb(NO_3)_2} + \mathrm{H_2SO_4} \rightarrow \mathrm{PbSO_4}\!\downarrow + 2\,\mathrm{HNO_3}\]
\[\mathrm{Pb(NO_3)_2} + \mathrm{Na_2SO_4} \rightarrow \mathrm{PbSO_4}\!\downarrow + 2\,\mathrm{NaNO_3}\]
Why not save a step and pour dilute sulphuric acid straight onto lead carbonate? Because it does not work. An insoluble crust of lead sulphate forms on the surface of the lead carbonate and seals it off, so the reaction stops almost as soon as it starts. The two-step route exists precisely to avoid that crust.
Lead chloride is prepared from insoluble lead oxide or lead carbonate by the same method. First convert the \(\mathrm{PbO}\) or \(\mathrm{PbCO_3}\) into soluble lead nitrate with dilute nitric acid, and then add hydrochloric acid or sodium chloride solution.
\[\mathrm{Pb(NO_3)_2} + 2\,\mathrm{NaCl} \rightarrow \mathrm{PbCl_2}\!\downarrow + 2\,\mathrm{NaNO_3}\]
3.19 Laboratory Preparation Of Some Salts
The methods of the last three sections are put to work here on seven actual preparations. Read each one with two questions in mind: why was this method chosen for this salt, and how does the chemist know when to stop?
3.19.1 Iron(III) Chloride By Direct Combination
Iron(III) chloride is prepared by synthesis — the direct combination of its elements.
Iron is taken in a combustion tube. Dry chlorine gas, obtained by passing the gas through concentrated sulphuric acid, is passed through the combustion tube, and the air inside the apparatus is first expelled. The arrangement is shown in Fig. 3.19-1.
Diagram
Fig. 3.19-1 — Preparation Of Iron(III) Chloride Left to right: two wash bottles containing concentrated sulphuric acid through which chlorine enters, connected to a horizontal combustion tube clamped on a stand and holding iron wire, heated below by a burner; the far end of the combustion tube leads down into a conical receiver in which solid iron(III) chloride collects, and the receiver is fitted with a guard tube of fused calcium chloride. Label the chlorine inlet, the concentrated sulphuric acid, the iron wire, the heat, the solid iron(III) chloride and the fused calcium chloride.
Once the reaction starts, the iron turns red hot, because the reaction is exothermic, and the external heating is then suspended. Iron(III) chloride is formed, volatilises, and condenses in the receiver as a brown scale.
\[2\,\mathrm{Fe}\text{(heated)} + 3\,\mathrm{Cl_2}\text{(dry)} \rightarrow 2\,\mathrm{FeCl_3}\]
Everything in this apparatus is arranged to keep water out. Iron(III) chloride is highly deliquescent — it absorbs moisture from the air and dissolves in it — so the product is kept dry with the help of fused calcium chloride, which acts as the drying agent.
There is no short cut through the hydrated salt. Anhydrous \(\mathrm{FeCl_3}\) cannot be prepared by simply heating hydrated ferric chloride, \(\mathrm{FeCl_3\!\cdot\!6H_2O}\), because on heating it gives iron(III) oxide, water and hydrogen chloride instead.
\[2,[\mathrm{FeCl_3\!\cdot\!6H_2O}] \xrightarrow{\ \Delta\ } \mathrm{Fe_2O_3} + 9\,\mathrm{H_2O} + 6\,\mathrm{HCl}\]
Anhydrous iron(III) chloride is used as a catalyst in certain reactions. Chlorine is poisonous, so the whole preparation is carried out in a fume cupboard.
3.19.2 Zinc Sulphate And Iron(II) Sulphate By Displacement
Both these salts are prepared by displacement — the action of a dilute acid, here dilute sulphuric acid, on an active metal.
Zinc Sulphate, Or White Vitriol
\[\mathrm{Zn}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{ZnSO_4}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{ZnSO_4} + 7\,\mathrm{H_2O} \rightarrow \mathrm{ZnSO_4\!\cdot\!7H_2O}\]
The procedure has four steps, and Fig. 3.19-2 follows them through.
- Take dilute sulphuric acid — 1 volume of acid to 5 volumes of water, the acid poured into the water — in a beaker and heat it on a wire gauze. Add some granulated zinc, a little at a time, with constant stirring.
- Effervescence takes place because of the liberation of hydrogen gas. Add zinc till the metal settles at the base of the beaker. When effervescence stops, it indicates that all the acid has been used up.
- The excess of zinc is filtered off.
- Collect the solution in a china dish. Evaporate the solution to get crystals. Filter, wash them with water, and dry them between the folds of a filter paper.
Diagram
Fig. 3.19-2 — Preparation Of Zinc Sulphate Crystals Four panels. (a) A beaker of dilute sulphuric acid with granulated zinc at the bottom, heated on a wire gauze and tripod, with bubbles of hydrogen rising; label the dilute sulphuric acid, the zinc metal and the heat. (b) A funnel with filter paper above a beaker, separating the excess zinc and impurities from the clear zinc sulphate solution; label the excess zinc and the pure zinc sulphate solution. (c) A china dish of the solution on a tripod over a burner, evaporated till crystals begin to appear. (d) The dish set aside to cool, with needle-shaped crystals of zinc sulphate in it.
The white, needle-shaped crystals obtained are of hydrated zinc sulphate, commonly known as white vitriol.
Iron(II) Sulphate, Or Green Vitriol
\[\mathrm{Fe}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
\[\mathrm{FeSO_4} + 7\,\mathrm{H_2O} \rightarrow \mathrm{FeSO_4\!\cdot\!7H_2O}\]
Filter the hot solution formed and allow the filtrate to evaporate at room temperature. Pale green crystals of hydrated iron(II) sulphate are formed. The solution is not boiled down as the zinc solution was, because iron(II) salts are slowly oxidised on strong heating in air.
3.19.3 Lead Chloride And Calcium Carbonate By Double Decomposition
An insoluble salt can be prepared by mixing the solutions of two appropriate salts, each containing one of its ions. The mixing results in the formation of an insoluble solid precipitate, which can be filtered, washed and dried.
Lead Chloride
It is prepared by adding either dilute hydrochloric acid or sodium chloride solution to a solution of lead nitrate, as in Fig. 3.19-3.
\[\mathrm{Pb(NO_3)_2} + 2\,\mathrm{HCl} \rightarrow \mathrm{PbCl_2}\!\downarrow + 2\,\mathrm{HNO_3}\]
Diagram
Fig. 3.19-3 — Preparation Of Lead Chloride Four panels. (1) Sodium chloride solution being poured from a dropper into a beaker of lead nitrate solution, with a white solid settling; label both solutions. (2) A funnel with filter paper over a beaker, the white residue on the paper and sodium nitrate solution passing through; label the residue and the filtrate. (3) A china dish on a tripod over a burner, the precipitate with a little water being heated until it dissolves. (4) The dish cooling, with needle-shaped crystals of lead chloride forming. Label the crystals of PbCl₂.
The solution of lead nitrate and sodium chloride is then filtered. The heavy, white precipitate of lead chloride accumulates as residue on the filter paper. It is washed repeatedly with cold water. The precipitate is then taken in a china dish with some water and heated till it dissolves, and the solution is cooled. Pure, needle-shaped crystals of lead chloride are thus obtained.
Two facts make that last step work. Lead chloride dissolves in hot water though not in cold, so heating brings it into solution and cooling brings it back out — purer than before, because the impurities stay dissolved. Lead compounds are poisonous, so hands are washed after handling them and no solution is pipetted by mouth.
The chlorides of lead, silver and mercury(I), and the sulphates of barium, lead and calcium, are also prepared using this method.
Calcium Carbonate
It is prepared by adding sodium carbonate solution to a hot solution of calcium chloride in a beaker, till the former is in excess.
\[\mathrm{CaCl_2} + \mathrm{Na_2CO_3} \rightarrow \mathrm{CaCO_3}\!\downarrow + 2\,\mathrm{NaCl}\]
The white precipitate formed by the interchange of radicals is filtered and washed repeatedly with cold water and collected. The washed precipitate is dried to obtain an amorphous powder — that is, a powder with no crystal shape of its own.
The carbonates of all the metals, except those of sodium, potassium and ammonium, are prepared by the above method.
3.19.4 Sodium Hydrogen Carbonate, An Acid Salt
The principle is to pass carbon dioxide gas into a cold solution of sodium carbonate.
\[\mathrm{Na_2CO_3} + \mathrm{CO_2} + \mathrm{H_2O} \rightarrow 2\,\mathrm{NaHCO_3}\]
- Dissolve nearly 5 grams of anhydrous sodium carbonate in about \(25\ \text{cm}^3\) of distilled water in a flask.
- Cool the solution by keeping the flask in a freezing mixture.
- Pass carbon dioxide gas into the solution.
- Crystals of sodium bicarbonate precipitate out after some time.
- Filter the crystals and dry them in the folds of a filter paper.
Diagram
Fig. 3.19-4 — Preparation Of Sodium Hydrogen Carbonate A conical flask containing sodium carbonate solution, standing in a trough packed with ice, and closed with a stopper carrying an inlet tube that dips into the solution and an outlet tube. Carbon dioxide enters through the inlet and excess gas leaves by the outlet; crystals of sodium hydrogen carbonate are shown settling at the bottom of the flask. Label the CO₂ inlet, the outlet, the ice, the sodium carbonate solution and the crystals.
The cold is essential rather than decorative. Sodium hydrogen carbonate is only sparingly soluble, and it is least soluble of all in cold water, so cooling is what forces the crystals out of the solution.
Sodium hydrogen carbonate is used in health salts. These are a mixture of dry citric acid or tartaric acid and sodium hydrogen carbonate. On adding water, the acid reacts with the \(\mathrm{NaHCO_3}\) to liberate carbon dioxide, which forms an effervescent drink and releases gas pressure in the stomach.
3.19.5 Neutralisation And Its Ionic Equation
Take the familiar reaction in aqueous solution and write out every ion that is present.
\[\mathrm{Na^{+}} + \mathrm{OH^{-}} + \mathrm{H^{+}} + \mathrm{Cl^{-}} \rightarrow \mathrm{Na^{+}} + \mathrm{Cl^{-}} + \mathrm{H_2O}\]
The sodium ion and the chloride ion appear on both sides, unchanged. They were dissolved before the reaction and they are dissolved after it, so they have taken no part in anything. Cancel them from both sides, and what is left is the reaction that actually happened.
\[\mathrm{H^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)} \rightleftharpoons \mathrm{H_2O}\text{(l)}\]
\[\mathrm{H_3O^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)} \rightleftharpoons 2\,\mathrm{H_2O}\text{(l)}\]
The same cancellation works for any strong acid with any strong alkali. Potassium hydroxide with sulphuric acid, for instance, leaves exactly the same ionic equation behind.
\[2\,\mathrm{KOH} + \mathrm{H_2SO_4} \rightarrow \mathrm{K_2SO_4} + 2\,\mathrm{H_2O}\]
This is why all such neutralisations release the same amount of heat for the same quantity of acid: the reaction taking place is always the joining of \(\mathrm{H^{+}}\) and \(\mathrm{OH^{-}}\) to make water. Neutralisation reactions are generally exothermic in nature.
3.19.6 Copper(II) Sulphate From An Insoluble Base
Copper(II) sulphate is prepared by the neutralisation of an insoluble base — the action of dilute sulphuric acid on copper hydroxide, copper oxide or copper carbonate.
\[\mathrm{Cu(OH)_2} + \mathrm{H_2SO_4} \rightarrow \mathrm{CuSO_4} + 2\,\mathrm{H_2O}\]
\[\mathrm{CuO} + \mathrm{H_2SO_4} \rightarrow \mathrm{CuSO_4} + \mathrm{H_2O}\]
\[\mathrm{CuCO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{CuSO_4} + \mathrm{H_2O} + \mathrm{CO_2}\text{(g)}\]
\[\mathrm{CuSO_4} + 5\,\mathrm{H_2O} \rightarrow \mathrm{CuSO_4\!\cdot\!5H_2O}\]
- Take dilute sulphuric acid in a beaker and heat it on a wire gauze.
- Add copper hydroxide, or black cupric oxide, or copper carbonate, in small quantities at a time, stirring the solution until no more of it dissolves and the excess compound settles to the bottom.
- Filter it hot, and collect the filtrate in a china dish. Evaporate the filtrate by heating to the point of crystallisation, and then allow it to cool.
- Collect the bright blue crystals of copper(II) sulphate pentahydrate, \(\mathrm{CuSO_4\!\cdot\!5H_2O}\), known as blue vitriol, and dry them.
The settled excess in step 2 is the signal that matters: while any base remains undissolved, no free acid can be left, so the solution filtered off in step 3 contains the salt and nothing else. Other salts which can be prepared by the same method are zinc sulphate, lead nitrate, calcium chloride, calcium nitrate and magnesium sulphate.
3.19.7 Sodium Sulphate Crystals By Titration
Sodium sulphate is prepared by the neutralisation of an alkali, caustic soda, with dilute sulphuric acid.
\[2\,\mathrm{NaOH} + \mathrm{H_2SO_4} \rightarrow \mathrm{Na_2SO_4} + 2\,\mathrm{H_2O}\]
This method is useful for preparing the soluble salts of sodium, potassium and ammonium only, since these are the metals whose oxides and hydroxides dissolve.
Here the earlier trick is unavailable. Both the reactants and the products are soluble, so an excess of either cannot be removed by filtration. It is therefore necessary to find out, on a small scale, the ratio of the solutions of the two reactants required for complete neutralisation before the preparation is carried out. That measurement is a titration, shown in Fig. 3.19-5.
Diagram
Fig. 3.19-5 — Neutralisation Of An Acid With A Base By Titration A burette graduated from 0 downwards, clamped vertically on a stand and filled with dilute sulphuric acid, its tap above the neck of a conical flask standing on the base of the stand. The conical flask contains sodium hydroxide solution with a drop of phenolphthalein, shown pink, and drops are falling from the burette into it. Label the burette, the dilute sulphuric acid, the tap, the conical flask and the sodium hydroxide solution.
- \(25\ \text{cm}^3\) of sodium hydroxide solution is measured and transferred to a conical flask. A drop of the indicator phenolphthalein is added, and the solution turns pink.
- A burette is filled with dilute sulphuric acid and the reading on the burette is noted. The acid is run in drops into the conical flask until the last drop turns the solution colourless. This is known as the end point.
- Note the burette reading at the end point. The difference between the two readings is the volume of the acid required.
The acid and the alkali, in the same ratio as found above, are now mixed in an evaporating basin. The resulting solution contains sodium sulphate and water, and on evaporating this solution, crystals of \(\mathrm{Na_2SO_4\!\cdot\!10H_2O}\) are obtained.
Two variations are worth noting. In place of sodium hydroxide, sodium carbonate can be used in this reaction; methyl orange is then used as the indicator. And to prepare the acid salt \(\mathrm{NaHSO_4}\), add the same volume of sulphuric acid again to the product formed in preparing the normal salt — that is, for a given amount of alkali, the addition of twice the normal amount of acid results in the formation of the acid salt.
\[\mathrm{NaOH} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{H_2O}\]
3.20 General Properties Of Salts
3.20.1 Conduction, Volatility And Solubility
Salts are electrovalent compounds — they are built of ions held together by the attraction between opposite charges. From that one fact the first two properties follow.
They conduct electricity in their molten state as well as in their aqueous solutions, because in both states the compound has dissociated into ions and those ions are free to move. In the solid, the same ions are locked in place in the crystal, and the solid does not conduct.
They are non-volatile solids that form crystals. The attraction between the ions is strong in every direction, so a salt does not vaporise easily and its particles pack into the regular shapes you see in a crystal of blue vitriol or of common salt.
Solubility. Most of the salts are soluble in water, and their degree of solubility in water varies with temperature, as Fig. 3.20-1 shows.
Diagram
Fig. 3.20-1 — Variation In Solubility Of Some Salts With Temperature Line graph. Horizontal axis: temperature in °C, from 0 to 100. Vertical axis: solubility in grams, from 0 to 100. Draw seven labelled curves: potassium nitrate rising very steeply from a low value at 0 °C to the top of the scale; potassium bromide and potassium chlorate rising steeply; potassium chloride and potassium sulphate rising gently; sodium chloride almost flat at about 36 g across the whole range; calcium sulphate almost flat and very low near the axis. Label each curve on the curve itself.
Read the graph before reading on, and notice that the curves are not parallel. Potassium nitrate becomes far more soluble as the water is heated, so a hot saturated solution of it deposits a great many crystals on cooling. Sodium chloride is almost as soluble in cold water as in hot, which is why common salt cannot be purified by cooling a hot solution — it has to be crystallised by evaporating the water instead.
3.20.2 The Nature Of A Salt Solution And Hydrolysis
A salt is made from an acid and a base, and you might expect its solution to be neutral every time. Test the solutions and you find otherwise: some turn litmus blue, some turn it red, and some leave it purple. The parents decide the outcome.
A salt of a strong base and a weak acid gives an alkaline solution, with a pH more than 7. Sodium carbonate and potassium acetate are examples.
\[\mathrm{Na_2CO_3} + 2\,\mathrm{H_2O} \rightleftharpoons \mathrm{H_2CO_3} + 2\,\mathrm{NaOH}\]
\[\mathrm{CH_3COOK} + \mathrm{H_2O} \rightleftharpoons \mathrm{CH_3COOH} + \mathrm{KOH}\]
Follow the reasoning. Water splits the salt back towards its parents. The strong alkali that reappears is fully ionised and floods the solution with \(\mathrm{OH^{-}}\) ions; the weak acid that reappears holds on to its hydrogen and contributes almost no \(\mathrm{H^{+}}\). Hydroxide ions therefore win, and the solution is alkaline.
A salt of a strong acid and a weak base gives an acidic solution, with a pH less than 7. Iron(III) chloride, copper sulphate, aluminium chloride and ammonium sulphate are examples.
\[\mathrm{FeCl_3} + 3\,\mathrm{H_2O} \rightleftharpoons 3\,\mathrm{HCl} + \mathrm{Fe(OH)_3}\]
\[\mathrm{CuSO_4} + 2\,\mathrm{H_2O} \rightleftharpoons \mathrm{H_2SO_4} + \mathrm{Cu(OH)_2}\]
The argument is the same one, reversed: the strong acid ionises fully and the weak base barely at all, so hydrogen ions are in excess.
A salt of a strong acid and a strong base gives a neutral solution. Sodium chloride, sodium sulphate, potassium sulphate and potassium nitrate are of this kind. Here the \(\mathrm{H^{+}}\) ions of the acid have completely combined with the \(\mathrm{OH^{-}}\) ions of the base, and nothing is left over to tip the balance either way.
A salt of a weak acid and a weak base may be acidic, alkaline or neutral, depending on the dissociated ions and the undissociated molecules present. Ammonium acetate, \(\mathrm{CH_3COONH_4}\), and ammonium carbonate, \(\mathrm{(NH_4)_2CO_3}\), belong here.
3.21 Water Of Crystallisation
Pick up a crystal of blue vitriol and it is perfectly dry to the touch. Heat it, and water comes off. The water was inside the crystal all along, built into its structure in a fixed amount.
This water is in loose chemical combination with the salt, and it can be driven out by heating the crystals above \(100^\circ\text{C}\). A salt that contains it is said to be hydrated; the same salt with the water removed is anhydrous.
Copper sulphate crystals on heating give out their water of crystallisation, and the change of colour is unmistakable.
\[\mathrm{CuSO_4\!\cdot\!5H_2O} \underset{\text{cool}}{\overset{\text{heat}}{\rightleftharpoons}} \mathrm{CuSO_4} + 5\,\mathrm{H_2O}\]
The bright blue crystal becomes a white powder. Add water to that powder and the blue returns, which shows that nothing has been destroyed — only the water has been taken out and put back.
The water vapour produced can be tested with cobalt chloride paper, as in Fig. 3.21-1. The liquid water turns the blue \(\mathrm{CoCl_2}\) paper pale pink, almost colourless.
\[\mathrm{CoCl_2} + 6\,\mathrm{H_2O} \rightarrow \mathrm{CoCl_2\!\cdot\!6H_2O}\]
Diagram
Fig. 3.21-1 — Heating Hydrated Copper Sulphate A hard-glass test tube clamped at a slight downward slope towards its mouth, holding blue crystals at the closed end, heated by a burner. Droplets of water are shown collecting near the cooler mouth of the tube, and a strip of blue cobalt chloride paper held at the mouth is turning pale pink. Label the blue crystals, the white residue forming beneath them, the condensed water and the cobalt chloride paper. Mark the slope of the tube with the note that the mouth points downwards so that water runs out and does not fall back on the hot glass.
Several familiar salts are hydrated, and most of them are better known by their common names than by their chemical ones.
| Common name | Chemical name | Formula |
|---|---|---|
| Washing soda crystals | Sodium carbonate decahydrate | \(\mathrm{Na_2CO_3\!\cdot\!10H_2O}\) |
| Epsom salt | Magnesium sulphate heptahydrate | \(\mathrm{MgSO_4\!\cdot\!7H_2O}\) |
| Potash alum | Hydrated potassium aluminium sulphate | \(\mathrm{K_2SO_4\!\cdot\!Al_2(SO_4)_3\!\cdot\!24H_2O}\) |
| Glauber’s salt | Sodium sulphate decahydrate | \(\mathrm{Na_2SO_4\!\cdot\!10H_2O}\) |
| Blue vitriol | Copper(II) sulphate pentahydrate | \(\mathrm{CuSO_4\!\cdot\!5H_2O}\) |
| Plaster of Paris | Calcium sulphate semihydrate | \(\mathrm{CaSO_4\!\cdot\!\tfrac{1}{2}H_2O}\) |
| White vitriol | Zinc sulphate heptahydrate | \(\mathrm{ZnSO_4\!\cdot\!7H_2O}\) |
| Gypsum | Hydrated calcium sulphate | \(\mathrm{CaSO_4\!\cdot\!2H_2O}\) |
Table 3.14 Some hydrated salts and their common names
Being crystalline is not the same as being hydrated. All crystalline salts do not contain water of crystallisation — common salt, \(\mathrm{NaCl}\), nitre, \(\mathrm{KNO_3}\), sugar, \(\mathrm{C_{12}H_{22}O_{11}}\), potassium permanganate, \(\mathrm{KMnO_4}\), and ammonium chloride, \(\mathrm{NH_4Cl}\), are all crystalline and all anhydrous.
Heat crystals of sodium chloride and you will still hear something. They produce a crackling sound, called decrepitation, which is due to the bigger crystals breaking up into smaller ones — not to any water escaping, since there is none to escape.
3.22 Efflorescence, Hygroscopy And Deliquescence
Leave three different solids in three open dishes for a week. One crumbles to powder, one grows damp, and one turns into a puddle. All three are exchanging water with the air, and the direction and the extent of the exchange give the three phenomena of this section their names.
3.22.1 Efflorescent Substances
Such substances are called efflorescent substances, and they become powdery. Having lost their water of crystallisation, the crystals can no longer hold their shape, so the crystalline form collapses into a powder.
Washing soda, hydrated sodium carbonate, becomes a monohydrate when exposed to dry air.
\[\mathrm{Na_2CO_3\!\cdot\!10H_2O} \xrightarrow{\ \text{dry air}\ } \mathrm{Na_2CO_3\!\cdot\!H_2O} + 9\,\mathrm{H_2O}\]
Glauber’s salt, hydrated sodium sulphate, becomes a powdery anhydrous sodium sulphate when exposed to air.
\[\mathrm{Na_2SO_4\!\cdot\!10H_2O} \xrightarrow{\ \text{dry air}\ } \mathrm{Na_2SO_4} + 10\,\mathrm{H_2O}\]
Epsom salt, magnesium sulphate heptahydrate, becomes a monohydrate when exposed to dry air.
\[\mathrm{MgSO_4\!\cdot\!7H_2O} \xrightarrow{\ \text{dry air}\ } \mathrm{MgSO_4\!\cdot\!H_2O} + 6\,\mathrm{H_2O}\]
The higher the temperature of the air, the greater the efflorescence. Warm air absorbs more water than cold air does, so as the temperature rises and the moisture in the air falls, water is drawn out of the crystals faster.
3.22.2 Hygroscopic Substances
Some salts absorb moisture, that is water vapour, from the atmosphere without dissolving in it. Such substances are called hygroscopic substances, and the phenomenon is called hygroscopy.
Notice the words without dissolving in it. A hygroscopic substance takes up water and stays a solid — damp, perhaps, but still solid. Some examples are these.
- Concentrated sulphuric acid, \(\mathrm{H_2SO_4}\)
- Phosphorus pentoxide, \(\mathrm{P_2O_5}\)
- Quicklime, \(\mathrm{CaO}\)
- Silica gel
3.22.3 Deliquescent Substances
Such a substance is called a deliquescent substance. The examples are caustic soda, \(\mathrm{NaOH}\), caustic potash, \(\mathrm{KOH}\), magnesium chloride, \(\mathrm{MgCl_2}\), zinc chloride, \(\mathrm{ZnCl_2}\), calcium chloride, \(\mathrm{CaCl_2}\), ferric chloride, \(\mathrm{FeCl_3}\), zinc nitrate, \(\mathrm{Zn(NO_3)_2}\), and copper nitrate, \(\mathrm{Cu(NO_3)_2}\).
Deliquescence occurs when the vapour pressure of the salt is much lower compared to the atmospheric vapour pressure. Vapour pressure here is a measure of how strongly water tends to escape from a surface; water moves from where that tendency is greater to where it is smaller, so a salt with a very low vapour pressure pulls water out of damp air.
That comparison also explains when each phenomenon is seen. Deliquescence is minimised during dry conditions, when the air holds little water to give up. Efflorescence, for the same reason, is maximised in dry conditions.
Pure sodium chloride is not deliquescent at all, and this is worth being clear about. It is the commercial version of the salt that contains the impurities, and those impurities are what absorb the water. The impurity can be removed by passing a current of dry hydrogen chloride gas through a saturated solution of the affected salt. Pure sodium chloride is produced as a precipitate, which can be recovered by filtering and washing — first with a little water and finally with alcohol.
One simple test separates the three phenomena. Efflorescent substances lose weight when exposed to the atmosphere, while hygroscopic and deliquescent substances gain weight.
3.23 Drying Agents And Dehydrating Agents
Certain substances remove moisture from other substances and are therefore called desiccants, or desiccating agents, or drying agents. Almost all hygroscopic substances are desiccating agents, which is exactly what you would expect from Section 3.22.2 — a substance that pulls water vapour out of the air will pull it out of a gas passed over it just as readily.
A dehydrating agent goes further. It removes water molecules even from compounds — not water that was merely clinging to the surface, but hydrogen and oxygen that were part of the compound itself. Concentrated sulphuric acid can remove the water molecules from blue vitriol, \(\mathrm{CuSO_4\!\cdot\!5H_2O}\), turning the blue crystals white, so it is a dehydrating agent as well as a drying agent.
That is the distinction to hold on to. Drying takes away water that is already water; dehydrating takes away the elements of water from a substance that contained no water as such.
| Drying agent | Dehydrating agent |
|---|---|
| Removes moisture from other substances. | Removes chemically combined elements of water, in the ratio 2 : 1 of hydrogen to oxygen, from a compound. |
| Used to dry gases such as chlorine, sulphur dioxide and hydrogen chloride; also used in desiccators to keep substances dry. | Used to prepare substances such as carbon monoxide and sugar charcoal. |
| Represents a physical change. | Represents a chemical change. |
| Examples: phosphorus pentoxide, fused calcium chloride, calcium oxide, silica gel, concentrated sulphuric acid. | Example: concentrated sulphuric acid. |
Table 3.16 Differences between a drying agent and a dehydrating agent
Concentrated sulphuric acid appears in both columns, and this is not a mistake. It dries a gas passed through it, which is a physical change and leaves the gas chemically untouched, and it dehydrates a compound such as blue vitriol or sugar, which is a chemical change and leaves a different substance behind.
A drying agent must also be chosen to suit the gas. Quicklime is basic, so it cannot be used to dry an acidic gas such as hydrogen chloride, which it would react with; concentrated sulphuric acid is used instead. In Section 3.19.1 you saw both principles at work — sulphuric acid drying the chlorine, and fused calcium chloride keeping the product dry afterwards.
3.24 Summary
- Compounds are sorted into three classes — acids, bases and salts. Acids are sour, bases are bitter and soapy, and a salt is the product of their reaction. Indicators, not taste, are used to tell them apart.
- Every acid contains hydrogen, every base is a metallic oxide or hydroxide or ammonium hydroxide, and every salt is formed by replacing the hydrogen of an acid with a metallic or ammonium ion.
- An acid is a compound containing one or more hydrogen atoms which, dissolved in water, produces hydronium ions as the only positive ions. A bare \(\mathrm{H^{+}}\) cannot exist in water, so it joins a water molecule: \(\mathrm{H^{+}} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}}\). All the properties of acids are the properties of this ion.
- By source, acids are organic (from plants, contain carbon, weak) or inorganic or mineral (from minerals, generally without carbon, strong). Carbonic acid is the exception — a weak mineral acid containing carbon.
- By molecular composition, acids are oxyacids, containing oxygen, or hydracids, containing hydrogen and a non-metal only.
- Basicity is the number of hydronium ions one molecule of an acid produces in water. It counts ionisable hydrogen, not total hydrogen: acetic acid has four hydrogen atoms and is monobasic. Dibasic and tribasic acids ionise in stages and form acid salts as well as normal salts.
- Concentration is the amount of acid in a definite amount of solution; strength is the concentration of hydronium ions produced. Dilute hydrochloric acid is stronger than concentrated acetic acid. Acid is always poured into water, never water into acid.
- Acids are prepared by synthesis of the elements, by dissolving an acidic oxide in water (oxyacids only), by oxidising a non-metal with concentrated nitric acid, and by displacing a volatile acid from its salt with non-volatile sulphuric acid.
- Acids are sour and corrosive, are electrolytes, turn blue litmus red, methyl orange red or pink, and leave phenolphthalein colourless.
- Dilute hydrochloric and sulphuric acids give hydrogen with active metals; nitric acid does not, as it oxidises the hydrogen formed.
- Dilute acids decompose salts: carbonates and hydrogen carbonates give carbon dioxide, sulphites and hydrogen sulphites give sulphur dioxide, and sulphides give hydrogen sulphide. Chlorides and nitrates react only with concentrated sulphuric acid on warming, unless an insoluble product is formed.
- A base reacts with the hydronium ions of an acid to give salt and water only. The word only excludes lead(IV) oxide, which also releases chlorine.
- An alkali is a base soluble in water, giving \(\mathrm{OH^{-}}\) as the only negative ion. All alkalis are bases, but not all bases are alkalis: \(\mathrm{Fe(OH)_3}\) and \(\mathrm{Cu(OH)_2}\) are insoluble and so are not alkalis.
- Sodium and potassium hydroxides are strong alkalis; calcium and ammonium hydroxides are weak. Acidity is the number of hydroxide ions one molecule of a base supplies, giving monoacidic, diacidic and triacidic bases.
- Bases are prepared from metals and oxygen, from reactive metals and water, from soluble oxides and water, by double decomposition for insoluble hydroxides, by heating sulphides in oxygen, by decomposing carbonates and nitrates, and by dissolving ammonia in water.
- Alkalis turn red litmus blue, methyl orange yellow and phenolphthalein pink; insoluble bases affect no indicator. They absorb carbon dioxide, neutralise acids, precipitate insoluble hydroxides from salts of heavy metals, and give ammonia when warmed with an ammonium salt.
- Zinc, aluminium and lead hydroxides are amphoteric and dissolve in excess sodium or potassium hydroxide, forming soluble zincates and aluminates.
- The pH scale, devised by Sorensen, measures hydrogen ion concentration: \(\mathrm{pH} = -\log_{10}[\mathrm{H^{+}}]\). It runs from 0 to 14, with 7 neutral, below 7 acidic and above 7 alkaline. Each unit is a tenfold change.
- Universal indicator gives green at pH 7, shading through yellow, pink and red as acidity rises, and through blue, indigo and violet as alkalinity rises. pH paper gives a quick reading of the same kind.
- pH governs the body’s narrow range of 7.0 to 7.8, the choice of crop for a soil, acid rain below pH 5.6, medical diagnosis, digestion and antacids, tooth decay below pH 5.5, and the relief of a bee sting with baking soda.
- A salt is formed by the partial or total replacement of the ionisable hydrogen of an acid by a metallic or ammonium ion. In ionic terms, it gives a positive ion other than \(\mathrm{H^{+}}\) and a negative ion other than \(\mathrm{OH^{-}}\).
- There are six kinds of salts: normal, acid, basic, complex, double and mixed. Acid salts retain replaceable hydrogen and show acidic properties; basic salts retain a hydroxyl group from the base.
- Soluble salts are prepared by direct combination, simple displacement, decomposition of a salt by a dilute acid, titration of an alkali, neutralisation of an insoluble base, and the action of alkalis on aluminium and zinc. They are recovered by evaporation followed by crystallisation.
- Solubility rules decide the method: all compounds of ammonium, sodium and potassium, and all nitrates, are soluble; carbonates, sulphides, sulphites and phosphates are insoluble except those three.
- Insoluble salts are prepared by direct combination, by combining an acidic with a basic oxide, and above all by precipitation, in which two solutions exchange radicals and a solid falls out. An insoluble salt is made from another insoluble salt by first converting it into a soluble nitrate.
- Neutralisation is the complete reaction of the hydrogen ions of an acid with the hydroxide ions of a base to give salt and water only. Stripped of spectator ions it is always \(\mathrm{H^{+}}\text{(aq)} + \mathrm{OH^{-}}\text{(aq)} \rightleftharpoons \mathrm{H_2O}\text{(l)}\), and it is exothermic.
- Salts are electrovalent, conduct in the molten state and in solution, and are non-volatile crystalline solids. Their solubility varies with temperature.
- The nature of a salt solution follows from its parents: strong acid with strong base gives a neutral solution, weak acid with strong base an alkaline one, strong acid with weak base an acidic one. This is hydrolysis.
- Water of crystallisation is the definite quantity of water built into a crystal, driven out above \(100^\circ\text{C}\); hydrated copper sulphate turns from blue to white, and the vapour turns cobalt chloride paper pink. Not every crystalline salt is hydrated.
- Efflorescent salts lose water of crystallisation to dry air and lose mass; hygroscopic substances absorb moisture without dissolving and gain mass; deliquescent substances absorb moisture until they dissolve in it.
- Drying agents remove moisture, a physical change; dehydrating agents remove the elements of water from a compound, a chemical change. Concentrated sulphuric acid does both.
Flowchart
FC 3.4 — The Chapter In One Line Acid supplies H⁺ (as H₃O⁺) → base supplies O²⁻ or OH⁻ → they neutralise → salt + water → is the salt soluble? → if soluble: prepare in solution by combination, displacement, decomposition, titration or an insoluble base, then evaporate and crystallise → if insoluble: precipitate it by double decomposition, then filter, wash and dry → test the solution’s nature by hydrolysis and its pH; test the crystals for water of crystallisation
3.25 Key Terms
- Acid — A compound containing one or more hydrogen atoms which, dissolved in water, gives hydronium ions as the only positive ions.
- Acid rain — Rain water with a pH below 5.6, made acidic by pollutants such as sulphur dioxide and oxides of nitrogen.
- Acid salt — A salt formed by the partial replacement of the ionisable hydrogen of a polybasic acid, and still holding replaceable hydrogen.
- Acidic oxide (acidic anhydride) — A non-metallic oxide which dissolves in water to give an acid.
- Acidity of a base — The number of hydroxide ions one molecule of a base can supply in aqueous solution.
- Alkali — A base that dissolves in water, giving hydroxide ions as the only negative ions.
- Amphoteric — Able to react both with acids and with strong alkalis, as the hydroxides of zinc, aluminium and lead do.
- Anhydrous — Containing no water of crystallisation.
- Base — A metallic oxide or hydroxide, or ammonium hydroxide, which reacts with an acid to give salt and water only.
- Basic hydroxide — A metallic hydroxide containing the \(\mathrm{OH^{-}}\) ion, which gives salt and water only with an acid.
- Basic oxide — A metallic oxide containing the \(\mathrm{O^{2-}}\) ion, which gives salt and water only with an acid.
- Basic salt — A salt formed by the partial replacement of the hydroxyl group of a di- or tri-acidic base, and still holding a hydroxyl group.
- Basicity of an acid — The number of hydronium ions produced by one molecule of the acid in aqueous solution.
- Binary acid — An acid whose molecule contains only two elements.
- Complex salt — A salt which on dissociation gives one simple ion and one complex ion.
- Concentrated acid — An acid containing very little water, or none.
- Concentration — The amount of acid present in a definite amount of its aqueous solution.
- Decrepitation — The crackling sound made by crystals, such as sodium chloride, breaking into smaller pieces on heating.
- Degree of ionisation — The percentage of the dissolved molecules of an acid or alkali that have ionised in solution.
- Dehydrating agent — A substance that removes the chemically combined elements of water from a compound.
- Deliquescence — The absorption of moisture from the air by a soluble substance until it dissolves in that water to form a saturated solution.
- Dibasic acid — An acid producing two hydronium ions per molecule on ionisation.
- Dilute acid — An acid containing far more water than its own mass.
- Double decomposition — A change in which two compounds in solution exchange radicals and a solid precipitate is formed.
- Double salt — A salt formed by the union of two simple salts that dissolve in water and crystallise together.
- Drying agent — A substance that removes moisture from another substance without any chemical change in it.
- Efflorescence — The loss of water of crystallisation by a hydrated salt exposed to dry air, leaving it powdery.
- Heat of neutralisation — The heat liberated when one gram equivalent of an acid or a base is completely neutralised.
- Hydracid — An acid containing hydrogen and a non-metal, and no oxygen.
- Hydrated — Containing a definite quantity of water of crystallisation.
- Hydrolysis — The reaction of a salt with water that leaves the solution acidic or alkaline instead of neutral.
- Hydronium ion — The ion \(\mathrm{H_3O^{+}}\), formed when a hydrogen ion joins a water molecule.
- Hygroscopic substance — A substance that absorbs moisture from the air without dissolving in it.
- Indicator — A chemical which shows the nature of a solution by a sharp change in colour.
- Inorganic (mineral) acid — An acid obtained from minerals, generally containing no carbon, and usually strong.
- Ionisation — The splitting of a molecule into ions when it dissolves in water.
- Mixed acid anhydride — An oxide, such as \(\mathrm{NO_2}\), which gives two different acids with water.
- Mixed salt — A salt containing more than one basic or acid radical.
- Monoacidic base — A base supplying one hydroxide ion per molecule.
- Monobasic acid — An acid producing one hydronium ion per molecule on ionisation.
- Neutralisation — The complete reaction of the hydrogen ions of an acid with the hydroxide ions of a base to give salt and water only.
- Non-volatile acid — An acid that does not vaporise easily, such as sulphuric acid.
- Normal salt — A salt formed by the complete replacement of the ionisable hydrogen of an acid, leaving none.
- Olfactory indicator — A substance whose smell, rather than colour, changes with the nature of a solution.
- Organic acid — An acid obtained usually from plants, containing carbon, and weak.
- Oxyacid — An acid containing oxygen along with hydrogen and another element.
- pH paper — Paper used to judge whether a solution is acidic, basic or neutral from the colour a drop of it produces.
- pH scale — A scale from 0 to 14 measuring hydrogen ion concentration, defined by \(\mathrm{pH} = -\log_{10}[\mathrm{H^{+}}]\).
- Salt — A compound formed by the partial or total replacement of the ionisable hydrogen of an acid by a metallic or ammonium ion.
- Strength of an acid — The measure of the concentration of hydronium ions the acid produces in its solution.
- Strong acid — An acid that ionises almost completely in water.
- Strong alkali — An alkali that dissociates almost completely in aqueous solution.
- Titration — The measured addition of one solution to another to find the volumes that exactly neutralise each other.
- Triacidic base — A base supplying three hydroxide ions per molecule.
- Tribasic acid — An acid producing three hydronium ions per molecule on ionisation.
- Universal indicator — A mixture of indicator dyes giving a different colour at each pH value.
- Volatile acid — An acid that vaporises easily at room temperature or on gentle heating.
- Water of crystallisation — The definite quantity of water combined with a salt in its crystals.
- Weak acid — An acid that ionises only partly in water.
- Weak alkali — An alkali that dissociates only partly in aqueous solution.
3.26 Points to Ponder
The properties of every acid are traced in this chapter to one ion, the hydronium ion, and the properties of every alkali to the hydroxide ion. If that is so, what exactly is the negative ion of an acid doing in the solution — and would an acid behave differently if it had no negative ion at all?
A solution of pH 3 holds a hundred times as many hydrogen ions as one of pH 5, yet the two feel equally sour on a tested indicator paper. Why might a scale built on powers of ten be more useful to a chemist than one built on the concentrations themselves?
Zinc hydroxide dissolves in an acid and also in a strong alkali. If a compound can behave as a base towards one reagent and as an acid towards another, how useful is it to speak of a substance as being an acid or a base at all, rather than as behaving like one in a given situation?
Every preparation of a soluble salt in this chapter ends with a way of knowing when to stop — an excess of solid that settles, effervescence that ceases, an indicator that changes colour. What would happen to the final crystals in each method if the chemist stopped too early, or too late?
Water appears on both sides of this chapter: it makes an acid an acid, it is a product of neutralisation, it sits inside crystals, and it moves in and out of salts left in the open air. Is water behaving as the same kind of participant in all four cases, or as something different each time?
Sodium chloride and potassium nitrate both dissolve in water to give neutral solutions, but copper sulphate gives an acidic one although it too was made from an acid and a base. What does that tell you about how much of a compound’s behaviour is decided by the compound itself, and how much by the substances it came from?
3.27 Objective Questions
Multiple Choice
- The only positively charged ion produced when an acid dissolves in water is: [1]
- \(\mathrm{OH^{-}}\)
- \(\mathrm{H_3O^{+}}\)
- \(\mathrm{NH_4^{+}}\)
- \(\mathrm{O^{2-}}\)
- An acid which contains hydrogen and a non-metal but no oxygen is called: [1]
- an oxyacid
- an organic acid
- a hydracid
- a tribasic acid
- The basicity of phosphorous acid, \(\mathrm{H_3PO_3}\), is: [1]
- one
- two
- three
- four
- A white precipitate formed on adding sodium hydroxide solution to a salt solution dissolves when more alkali is added. The metal ion present could be: [1]
- copper
- iron(II)
- zinc
- silver
- A solution has a pH of 11. It will: [1]
- turn methyl orange red
- leave phenolphthalein colourless
- give a green colour with universal indicator
- turn red litmus blue
- The gas given off when dilute hydrochloric acid is added to sodium sulphite is: [1]
- carbon dioxide
- sulphur dioxide
- hydrogen sulphide
- hydrogen
- Which of these salts is prepared by direct combination of its elements? [1]
- iron(II) chloride
- iron(III) chloride
- iron(II) sulphate
- sodium sulphate
- A substance that absorbs moisture from the air and finally dissolves in it is said to be: [1]
- efflorescent
- hygroscopic
- deliquescent
- anhydrous
Assertion And Reason
For each of the following, choose from: [1 each] (1) Both A and R are true and R is the correct explanation of A. (2) Both A and R are true but R is not the correct explanation of A. (3) A is true but R is false. (4) A is false but R is true.
Assertion (A): Dilute hydrochloric acid is a stronger acid than concentrated acetic acid. Reason (R): Strength depends on the number of hydronium ions released in solution, not on the amount of water present.
Assertion (A): Ferric hydroxide is a base but not an alkali. Reason (R): Ferric hydroxide is insoluble in water.
Assertion (A): Nitric acid is not used to prepare hydrogen from a metal. Reason (R): Nitric acid contains no hydrogen.
Assertion (A): Lead sulphate is not prepared by adding dilute sulphuric acid to lead carbonate. Reason (R): An insoluble crust of lead sulphate forms on the carbonate and stops the reaction.
Match The Following
- Match each salt in Column A with the method of preparation in Column B. [5]
| Column A | Column B |
|---|---|
| a. Zinc sulphate | i. Direct combination |
| b. Silver chloride | ii. Neutralisation of an insoluble base |
| c. Sodium nitrate | iii. Simple displacement |
| d. Copper(II) sulphate | iv. Titration |
| e. Lead sulphide | v. Precipitation |
Fill In The Blanks
An acid dissolved in water gives …………… ions as the only positive ions, while an alkali gives …………… ions as the only negative ions. [2]
A salt formed by the …………… replacement of the ionisable hydrogen of a polybasic acid is called an acid salt, while one formed by …………… replacement is called a normal salt. [2]
Blue vitriol turns …………… on heating, because it loses its ……………, and the vapour turns cobalt chloride paper …………… [3]
True Or False
Every base is an alkali. [1]
Efflorescent substances gain mass when exposed to the atmosphere. [1]
All nitrates are soluble in water. [1]
Name The Following
The acid present in the sting of a bee and an ant. [1]
An oxide which gives two different acids on reacting with water. [1]
A salt which on dissociation gives one simple ion and one complex ion. [1]
The reagent used to test a gas that turns lime water milky. [1]
Give One Word Or Phrase For
The number of hydroxide ions supplied by one molecule of a base in aqueous solution. [1]
The crackling sound produced when crystals of common salt are heated. [1]
The reaction of a salt with water that leaves the solution acidic or alkaline. [1]
Odd One Out, And Give A Reason
\(\mathrm{HCl}\), \(\mathrm{HNO_3}\), \(\mathrm{CH_3COOH}\), \(\mathrm{H_2SO_4}\) [1]
\(\mathrm{NaOH}\), \(\mathrm{KOH}\), \(\mathrm{Cu(OH)_2}\), \(\mathrm{NH_4OH}\) [1]
\(\mathrm{Na_2CO_3\!\cdot\!10H_2O}\), \(\mathrm{CuSO_4\!\cdot\!5H_2O}\), \(\mathrm{KNO_3}\), \(\mathrm{MgSO_4\!\cdot\!7H_2O}\) [1]
3.28 Check Your Concepts
Define an alkali, and give one example each of a strong alkali and a weak alkali. [2]
State why a bare hydrogen ion cannot exist in an aqueous solution. [2]
Give a reason why insoluble bases have no effect on litmus. [2]
Name the ions furnished in solution by an acid and by an alkali, and state the effect of each on litmus. [2]
Give a reason why acid is always poured into water and never water into acid. [2]
Explain, with an example, why acetic acid is described as a monobasic acid although its molecule contains four hydrogen atoms. [3]
Differentiate between a drying agent and a dehydrating agent, stating the basis of the difference and giving one example of each. [3]
Why is lead(IV) oxide not classified as a base, although it is a metallic oxide that reacts with hydrochloric acid? [3]
Explain why a solution of copper sulphate is acidic while a solution of potassium nitrate is neutral. [3]
Differentiate between the basicity of an acid and the acidity of a base, stating the basis of difference, and give one example of each with the value two. [4]
What would happen if dilute sulphuric acid were used in place of dilute hydrochloric acid to obtain carbon dioxide from calcium carbonate? Give the reason. [4]
Describe how you would distinguish between the three phenomena of efflorescence, hygroscopy and deliquescence, using a balance and three named substances. [4]
Explain why sodium sulphate must be prepared by titration, while copper(II) sulphate can be prepared without any indicator. Include the reason each method needs, or does not need, a measured volume. [5]
Explain what pH measures, why a change of one pH unit is a large change, and give three situations in everyday life in which the pH of a substance matters. [5]
3.29 Equations, Conversions & Identification
Write the balanced equation for the ionisation of sulphuric acid in water, showing both steps and the formation of hydronium ions. [2]
Write the balanced equation for the action of dilute hydrochloric acid on zinc sulphide, name the gas evolved, and state one test for it. [3]
Complete and balance the following: [4]
- \(\mathrm{P_2O_5} + \mathrm{H_2O} \rightarrow\)
- \(\mathrm{Al(OH)_3} + \mathrm{HNO_3} \rightarrow\)
- \(\mathrm{Zn(OH)_2} + \mathrm{NaOH} \rightarrow\)
- \(\mathrm{NH_4Cl} + \mathrm{Ca(OH)_2} \xrightarrow{\ \Delta\ }\)
State what you would see, and write the balanced equation, when sodium hydroxide solution is added to iron(II) sulphate solution. [3]
How would you obtain each of the following? Give the balanced equation in each case. [4]
- Iron(II) chloride from iron
- Iron(III) chloride from iron
How would you obtain lead(II) chloride starting from lead(II) oxide? Give both steps with balanced equations, and state why the conversion cannot be carried out in a single step. [4]
Identify the substance in each case, giving your reasoning. [4]
- A blue crystalline solid that turns white on heating, and whose vapour turns cobalt chloride paper pink.
- A colourless gas with the smell of rotten eggs, given off when a dilute acid acts on a metal salt.
- A white solid which, left in an open dish, crumbles to a powder and loses mass.
- A solution that turns methyl orange yellow and phenolphthalein pink.
Study Fig. 3.19-1, which shows the laboratory preparation of iron(III) chloride, and answer the following. [4]
- Name the liquid through which the chlorine is passed, and state its purpose.
- State the purpose of the fused calcium chloride in the guard tube.
- Why is the external heating stopped soon after the reaction begins?
- Write the balanced equation for the reaction taking place in the combustion tube.
Write the balanced equation for the preparation of each of the following salts, and name the method used. [4]
- Barium sulphate from barium chloride
- Magnesium sulphate from magnesium carbonate
Give balanced equations for the following conversions, A to C. [3] \[\mathrm{Zn} \xrightarrow{\ \mathrm{A}\ } \mathrm{ZnSO_4} \xrightarrow{\ \mathrm{B}\ } \mathrm{Zn(OH)_2} \xrightarrow{\ \mathrm{C}\ } \mathrm{Na_2ZnO_2}\]