Chapter 4: Analytical Chemistry

In the 1840s a young German student named Carl Remigius Fresenius sat in Justus von Liebig’s laboratory at Giessen with a problem that had nothing to do with discovering new elements. The elements were being found fast enough. What nobody could do reliably was answer a much plainer question: given a bottle of unlabelled powder, what is in it?

Chemists of the time had scattered tests. Add this, look for a colour; add that, look for a solid falling out of the liquid. But the tests were kept in notebooks and passed on by word of mouth, and two chemists testing the same sample could easily disagree. Fresenius did something ordinary and enormously useful. He wrote the tests down in order, so that each step ruled some possibilities out and pointed to the next step. His manual of qualitative analysis appeared while he was still a student, and it turned a collection of tricks into a method.

That method is still the one you will use. A few drops of an alkali fall into a clear salt solution, and a solid appears with a colour of its own — dirty green, reddish brown, pale blue, chalky white. Each colour belongs to one metal ion. Fig. 4.0-1 shows the bench you are about to work at.

Diagram

Fig. 4.0-1 — Reading A Metal Ion From Its Precipitate A wooden test-tube rack holding five test tubes, each half filled with a clear salt solution, with a dropper poised over the first tube releasing a drop of alkali; in each tube show a precipitate settling — label them left to right as dirty green, reddish brown, pale blue, white and gelatinous, and chalky white; label the dropper ‘sodium hydroxide solution, added drop by drop’ and label a reagent bottle standing beside the rack.

Curiosity Corner

Q1. Two clear, colourless liquids are mixed and a brightly coloured solid appears. Where could the colour have been hiding before they were mixed? Q2. Two different salt solutions both give a white precipitate with caustic soda. How could adding still more of the same alkali tell them apart? Q3. Sodium hydroxide and ammonium hydroxide both supply hydroxide ions, yet only one of them will bring down a precipitate from a calcium salt. What must be different about the two solutions? Q4. You already know that a metal reacting with an acid gives hydrogen. Why should zinc give hydrogen with a hot alkali as well, when most metals do not? Q5. A pale blue precipitate of copper hydroxide does not simply dissolve away in excess ammonia — the liquid turns a deep blue instead. What might the copper have become?

4.1 What Analysis Means

A chemist is often handed a substance and asked a question about it. Sometimes the question is what is this? Sometimes it is how much of each thing is in it? Finding out what a given sample is made of is called analysis — the determination of the chemical components present in a sample.

Analysis is not the same as making a new compound. Nothing new is being built here. The substance already exists, and the chemist’s task is to read it — to work out, from the way it behaves with other chemicals, what it contains.

4.1.1 Qualitative And Quantitative Analysis

Chemistry divides the work into two kinds, and the names tell you which question each answers.

Qualitative analysis identifies an unknown substance. It answers which: which metal ion is present in this salt, which gas has been given off. Quantitative analysis determines the composition of a mixture. It answers how much: what mass of each component is present, in what proportion.

A water sample from a river makes the difference clear. Qualitative analysis tells you that the water contains lead. Quantitative analysis tells you that it contains a certain mass of lead in every litre — and that is the figure which decides whether the water is safe. The two are set side by side in Table 4.1.

Compare Table

Table 4.1 — Qualitative Analysis Compared With Quantitative Analysis

Basis	Qualitative analysis	Quantitative analysis
Question answered	Which substances are present?	How much of each is present?
Result obtained	The name of a substance or an ion	A mass, a volume or a percentage
Usual evidence	A colour, a precipitate, a gas, a smell	A weighing, a titration, an instrument reading
Order of working	Done first, on an unknown sample	Done after the components are known

This chapter deals only with qualitative analysis, and with one small corner of it — the identification of metal ions in solution.

Side Note

The word analysis comes from the Greek analusis, meaning a loosening or a taking apart. A chemical analyst takes a substance apart, not with a tool, but with other chemicals.

4.1.2 Reagents And Precipitation

You cannot identify an ion by looking at it. You identify it by adding something that reacts with it in a known, recognisable way. A substance used in this manner is a reagent — a substance that reacts with another substance in a test.

Alkalis As Laboratory Reagents

Alkalis are among the most useful reagents in a school laboratory. When an alkali is added to certain salt solutions, the metal ion in the solution combines with hydroxide ions to form a metal hydroxide. Many of these hydroxides do not dissolve in water, so they appear as a solid suspended in the liquid. Because these solids have characteristic colours, the metal ion is identified by the colour of the solid that appears.

The forming of an insoluble solid when two solutions are mixed is called precipitation, and the solid itself is the precipitate. Notice the condition hidden in that sentence: a precipitate forms only when the new compound is insoluble in water. If it is soluble, it stays dissolved and you see nothing at all.

Definition

Precipitation The process of formation of an insoluble solid when two solutions are mixed. The solid so formed is called a precipitate.

It helps to picture what is happening among the particles, as shown in Fig. 4.1-1. A salt in solution is not a set of intact salt units floating about; it has separated into ions that move freely and independently through the water. When a second solution is added, its ions mix in among them. If any cation from one solution and any anion from the other can form a compound that water cannot hold in solution, those two ions lock together in enormous numbers and drop out of the liquid as a solid.

Diagram

Fig. 4.1-1 — What Happens When A Precipitate Forms Three panels in a row. Panel 1: a beaker of clear solution with scattered free cations labelled M²⁺ and anions labelled X²⁻. Panel 2: a second solution being poured in, showing free OH⁻ ions mixing among the M²⁺ ions. Panel 3: the same beaker with clumps of M(OH)₂ settling at the bottom, labelled ‘precipitate’, and the remaining free ions still in the clear liquid above, labelled ‘salt left in solution’.

Do Not Mix Up

Precipitate And Residue A precipitate appears from within a solution when two solutions are mixed. A residue is what is left behind on the filter paper or in the dish after filtering or evaporating.

Try Yourself

Dissolve a little washing soda in water in one glass and a little Epsom salt in water in another. Both are clear. Now pour one into the other and watch the liquid turn cloudy white. Predict, before you pour, whether the cloudiness will settle or stay suspended.

4.2 Colours Of The Salts And Their Solutions

Before any reagent is added, a salt has already told you something. Some salts are white powders that dissolve to give water-clear solutions. Others are green, blue or orange, and their solutions carry the same colour. That colour is the first piece of evidence in an analysis, and it is free.

4.2.1 Salts Of Representative Elements And Transition Elements

Where an element sits in the periodic table decides, in a general way, whether its salts are coloured. The representative elements are those of Groups 1, 2 and 13 to 17 — the two blocks at the left-hand and right-hand edges of the table. Their salts are generally colourless. The transition elements are those of Groups 3 to 12, the broad block through the middle, and their salts are generally coloured.

So sodium chloride, magnesium sulphate and aluminium nitrate are all colourless, while salts of copper, iron, nickel, chromium and manganese carry colours of their own. The colour belongs to the metal ion, not to the salt as a whole: every common copper(II) salt gives a blue solution, whatever anion it is paired with.

The word generally in both statements is doing real work. Zinc sits in Group 12, among the transition elements, yet the zinc ion is colourless — and so is the solution of every zinc salt you will meet in this chapter. A trend is a guide to what to expect, not a promise.

Do Not Mix Up

Colourless And White A colourless solution lets you read print through it. A white solid is not colourless — it is opaque and pale. Zinc sulphate solution is colourless; the zinc hydroxide precipitated from it is white.

4.2.2 Colourless And Coloured Ions

Since the colour of a salt solution comes from the ions dissolved in it, it is more useful to learn which ions are coloured than which salts are. The common cations are listed in Table 4.2 and the common anions in Table 4.3.

Table 4.2 — Common Colourless And Coloured Cations

Colourless cation	Symbol	Coloured cation	Symbol	Colour
Ammonium ion	\(\mathrm{NH_4^{+}}\)	Cupric ion	\(\mathrm{Cu^{2+}}\)	Blue
Sodium ion	\(\mathrm{Na^{+}}\)	Ferrous ion	\(\mathrm{Fe^{2+}}\)	Light green
Potassium ion	\(\mathrm{K^{+}}\)	Ferric ion	\(\mathrm{Fe^{3+}}\)	Yellowish brown
Calcium ion	\(\mathrm{Ca^{2+}}\)	Nickel ion	\(\mathrm{Ni^{2+}}\)	Green
Magnesium ion	\(\mathrm{Mg^{2+}}\)	Chromium ion	\(\mathrm{Cr^{3+}}\)	Green
Aluminium ion	\(\mathrm{Al^{3+}}\)	Manganese ion	\(\mathrm{Mn^{2+}}\)	Pink
Lead ion	\(\mathrm{Pb^{2+}}\)
Zinc ion	\(\mathrm{Zn^{2+}}\)

Two entries in that table are worth pausing over, because they will return again and again in this chapter. Iron forms two different cations, and they are not the same colour: the ferrous ion, \(\mathrm{Fe^{2+}}\), is light green, while the ferric ion, \(\mathrm{Fe^{3+}}\), is yellowish brown. A single element can therefore give two salts of quite different appearance, and telling them apart is a standard task in analysis.

Table 4.3 — Common Colourless And Coloured Anions

Colourless anion	Symbol	Coloured anion	Symbol	Colour
Chloride ion	\(\mathrm{Cl^{-}}\)	Permanganate ion	\(\mathrm{MnO_4^{-}}\)	Pink or purple
Sulphate ion	\(\mathrm{SO_4^{2-}}\)	Dichromate ion	\(\mathrm{Cr_2O_7^{2-}}\)	Orange
Carbonate ion	\(\mathrm{CO_3^{2-}}\)	Chromate ion	\(\mathrm{CrO_4^{2-}}\)	Yellow
Nitrate ion	\(\mathrm{NO_3^{-}}\)
Hydrogen carbonate ion	\(\mathrm{HCO_3^{-}}\)
Sulphide ion	\(\mathrm{S^{2-}}\)
Bromide ion	\(\mathrm{Br^{-}}\)
Acetate ion	\(\mathrm{CH_3COO^{-}}\)

The coloured anions follow the same pattern as the coloured cations. Each of them contains a transition metal — manganese in the permanganate ion, chromium in the dichromate and chromate ions — locked together with oxygen inside the anion. The metal has moved to the negative side of the compound, but it still brings its colour with it.

Reading the two tables together, you can predict the colour of a solution from its formula. Potassium permanganate gives a purple solution, because potassium contributes nothing to the colour and the permanganate ion contributes all of it. Copper sulphate gives a blue solution for the matching reason.

Curiosity Corner

Q1. Two clear, colourless liquids are mixed and a brightly coloured solid appears. Where could the colour have been hiding before they were mixed?

A. It cannot hide anywhere. If the precipitate is coloured, the metal ion was coloured before mixing, so one of the two solutions was coloured too. Truly colourless solutions give white or colourless products.

Chemistry at Work

Colour That Does A Job A very dilute solution of potassium permanganate is used to disinfect water and to wash wounds, and its deep purple colour is part of its usefulness: as the permanganate reacts, the purple fades to a brown stain, so you can see when the reagent has been spent. The same purple is why permanganate is a favourite in the laboratory — a change you cannot miss.

Side Note

A freshly made ferrous sulphate solution is pale green. Left standing in an open beaker, it slowly turns yellowish brown as air converts ferrous ions to ferric ions. This is why a test calls for a freshly prepared solution.

Extra Practice 4.1

A laboratory holds four unlabelled bottles of salt solution: one orange, one pale green, one blue and one colourless. Name a salt that could be in each bottle, and state the ion responsible for the colour in each coloured one.

4.3 Action Of Sodium Hydroxide Solution On Metallic Salt Solutions

Sodium hydroxide solution, known in the laboratory as caustic soda solution, is the first reagent reached for when a metal ion has to be identified. Added drop by drop to a metallic salt solution, it supplies hydroxide ions, which combine with the metal ion to form the metal hydroxide. Where that hydroxide is insoluble in water, it appears at once as a precipitate, and its colour names the metal ion.

The alkali is always added a drop at a time, and never in a rush. Some precipitates dissolve again when more alkali is added, and that second change is as much a part of the evidence as the first. Pour the alkali in quickly and you may never see the precipitate that formed and vanished.

Reagent Spotlight

Sodium Hydroxide Solution Caustic soda solution, a strong alkali → supplies hydroxide ions freely → precipitates metal hydroxides from salt solutions, and dissolves the few that react further with alkali. Strongly corrosive; keep it off skin and eyes and wear goggles.

Caustic soda solution is corrosive enough to attack skin, so the tube is held away from the face while the reagent is added, and any splash is washed off at once with plenty of water.

4.3.1 Calcium Salts

Take calcium nitrate solution, which is colourless. Adding caustic soda solution gives a white precipitate of calcium hydroxide, together with sodium nitrate, which stays dissolved.

calcium nitrate + sodium hydroxide → calcium hydroxide + sodium nitrate

\[\mathrm{Ca(NO_3)_2}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Ca(OH)_2}\!\downarrow + 2\,\mathrm{NaNO_3}\text{(aq)}\]

The downward arrow marks the substance that leaves the solution as a solid. Calcium hydroxide is only sparingly soluble in water, which is why it appears at all; and because it is sparingly soluble, adding excess caustic soda produces no further change. The white precipitate simply stays.

4.3.2 Iron Salts

Iron gives two cations, and caustic soda tells them apart in a single test. The colours of the two hydroxides are quite different, and neither dissolves in excess alkali.

4.3.2.1 Ferrous Salts

Ferrous sulphate solution is pale green, from the \(\mathrm{Fe^{2+}}\) ion. Caustic soda solution throws down a dirty green, jelly-like precipitate of ferrous hydroxide.

ferrous sulphate + sodium hydroxide → ferrous hydroxide + sodium sulphate

\[\mathrm{FeSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_2}\!\downarrow + \mathrm{Na_2SO_4}\text{(aq)}\]

The precipitate is described as gelatinous because it does not settle as a clean powder but hangs in the liquid as a soft, cloudy mass. Adding excess caustic soda changes nothing: ferrous hydroxide is insoluble in alkali.

Do Not Mix Up

Pale Green And Dirty Green Pale green is the colour of the ferrous solution, before any reagent is added. Dirty green is the colour of the ferrous hydroxide precipitate formed after the alkali goes in.

4.3.2.2 Ferric Salts

Ferric chloride solution is yellow, from the \(\mathrm{Fe^{3+}}\) ion. Here the metal ion carries three positive charges, so three hydroxide ions are needed for each iron atom, and the equation takes three formula units of the alkali.

ferric chloride + sodium hydroxide → ferric hydroxide + sodium chloride

\[\mathrm{FeCl_3}\text{(aq)} + 3\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_3}\!\downarrow + 3\,\mathrm{NaCl}\text{(aq)}\]

The precipitate of ferric hydroxide is reddish brown, and it too is insoluble in excess alkali. A pale green solution giving a dirty green precipitate and a yellow solution giving a reddish brown one are as unlike as any pair of tests in this chapter.

Example 4.1

Identifying An Iron Salt From Its Behaviour

1. A salt dissolves in water to give a pale green solution. Sodium hydroxide solution added drop by drop gives a gelatinous precipitate, whose colour does not change when excess alkali is added. Identify the cation and write the equation for the reaction with the alkali, taking the salt to be the sulphate.

The colour of the solution is the first clue. Of the coloured cations, only the ferrous ion, \(\mathrm{Fe^{2+}}\), gives a pale green solution.
The second clue confirms it. A ferrous salt gives a dirty green, gelatinous precipitate of ferrous hydroxide, and this hydroxide does not dissolve in excess alkali. A ferric salt would have given a reddish brown precipitate instead.
The cation present is the ferrous ion, \(\mathrm{Fe^{2+}}\), and the reaction is

\[\mathrm{FeSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_2}\!\downarrow + \mathrm{Na_2SO_4}\text{(aq)}\]

4.3.3 Copper Salts

Copper sulphate solution is blue. Caustic soda solution precipitates copper(II) hydroxide as a pale blue solid, leaving colourless sodium sulphate in solution.

copper sulphate + sodium hydroxide → copper hydroxide + sodium sulphate

\[\mathrm{CuSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Cu(OH)_2}\!\downarrow + \mathrm{Na_2SO_4}\text{(aq)}\]

Copper hydroxide is insoluble in excess caustic soda. Adding more alkali gives no further change — a point worth holding on to, because copper behaves very differently with the reagent of the next section.

4.3.4 Zinc Salts

Zinc sulphate solution is colourless. With a few drops of caustic soda it gives a white, gelatinous precipitate of zinc hydroxide.

zinc sulphate + sodium hydroxide → zinc hydroxide + sodium sulphate

\[\mathrm{ZnSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Zn(OH)_2}\!\downarrow + \mathrm{Na_2SO_4}\text{(aq)}\]

Now keep adding the alkali. The white precipitate thins and disappears, and the liquid becomes clear again. Zinc hydroxide does not merely dissolve; it reacts with the excess alkali to form sodium zincate, a soluble salt in which zinc has ended up inside the negative ion.

zinc hydroxide + sodium hydroxide (excess) → sodium zincate + water

\[\mathrm{Zn(OH)_2} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]

The two stages are worth watching carefully, and they are drawn in Fig. 4.3-1.

Diagram

Fig. 4.3-1 — A Precipitate That Dissolves In Excess Alkali Three test tubes side by side, each with a dropper above it. Tube 1: clear colourless zinc sulphate solution, labelled ‘before adding alkali’. Tube 2: same tube after a few drops of sodium hydroxide, showing a white cloudy gelatinous precipitate throughout, labelled ‘zinc hydroxide, white precipitate’. Tube 3: same tube after excess sodium hydroxide, clear again, labelled ‘sodium zincate, colourless solution’.

4.3.5 Lead Salts

Lead nitrate solution is colourless. Caustic soda solution gives a chalky white precipitate of lead hydroxide.

lead nitrate + sodium hydroxide → lead hydroxide + sodium nitrate

\[\mathrm{Pb(NO_3)_2}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Pb(OH)_2}\!\downarrow + 2\,\mathrm{NaNO_3}\text{(aq)}\]

Lead behaves like zinc when more alkali is added. The white precipitate dissolves, reacting with the excess caustic soda to give sodium plumbite, again a colourless solution.

lead hydroxide + sodium hydroxide (excess) → sodium plumbite + water

\[\mathrm{Pb(OH)_2} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2PbO_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]

Lead hydroxide has one further property that is useful in the laboratory: the white precipitate is readily soluble in acetic acid.

Curiosity Corner

Q2. Two different salt solutions both give a white precipitate with caustic soda. How could adding still more of the same alkali tell them apart?

A. Because the precipitates behave differently in excess. Zinc and lead hydroxides react with the extra alkali and dissolve to clear solutions; calcium hydroxide does not, and the white precipitate stays.

4.3.6 Ammonium Salts

Ammonium salts contain no metal ion at all, and they give no precipitate. Instead, when sodium hydroxide solution — or any water-soluble base — is heated with an ammonium salt, ammonia gas is driven off.

ammonium chloride + sodium hydroxide → sodium chloride + water + ammonia

\[\mathrm{NH_4Cl} + \mathrm{NaOH} \xrightarrow{\ \Delta\ } \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{NH_3}\!\uparrow\]

ammonium sulphate + sodium hydroxide → sodium sulphate + water + ammonia

\[\mathrm{(NH_4)_2SO_4} + 2\,\mathrm{NaOH} \xrightarrow{\ \Delta\ } \mathrm{Na_2SO_4} + 2\,\mathrm{H_2O} + 2\,\mathrm{NH_3}\!\uparrow\]

The gas is recognised by its sharp, unmistakable smell and by the fact that it turns moist red litmus paper blue. The smell is checked by wafting the vapour towards the nose with the hand, never by putting the face over the mouth of the tube.

Diagram

Fig. 4.3-2 — Testing The Gas Given Off By An Ammonium Salt A test tube held in a holder and heated gently over a small flame, containing a mixture of ammonium salt and sodium hydroxide solution; a strip of moist red litmus paper held on a glass rod at the mouth of the tube, with an arrow showing it turning blue; label the tube contents, the litmus paper and the escaping gas as ammonia.

This is a test for the ammonium ion rather than for a metal ion, and it is the only one in this section that needs heat.

4.3.7 Identifying A Cation With Sodium Hydroxide

All the observations of this section can be gathered into one scheme. Table 4.4 sets them out ion by ion, and FC 4.1 turns the same information into the order in which you would actually work at the bench.

Compare Table

Table 4.4 — Action Of Sodium Hydroxide Solution On Some Salt Solutions

Cation in solution	Colour of solution	Precipitate formed	Colour of precipitate	On adding excess alkali
\(\mathrm{Ca^{2+}}\)	Colourless	Calcium hydroxide	White	No change; sparingly soluble
\(\mathrm{Fe^{2+}}\)	Pale green	Ferrous hydroxide	Dirty green, gelatinous	Insoluble; no change
\(\mathrm{Fe^{3+}}\)	Yellow	Ferric hydroxide	Reddish brown	Insoluble; no change
\(\mathrm{Cu^{2+}}\)	Blue	Copper hydroxide	Pale blue	Insoluble; no change
\(\mathrm{Zn^{2+}}\)	Colourless	Zinc hydroxide	White, gelatinous	Dissolves, giving sodium zincate
\(\mathrm{Pb^{2+}}\)	Colourless	Lead hydroxide	Chalky white	Dissolves, giving sodium plumbite

Flowchart

FC 4.1 — Identifying A Cation With Sodium Hydroxide Solution Metal salt solution + a few drops of NaOH → observe the precipitate; branch A: white ppt → add excess NaOH → no change → Ca²⁺, or ppt dissolves → Pb²⁺ or Zn²⁺; branch B: coloured ppt → pale blue → Cu²⁺, dirty green → Fe²⁺, reddish brown → Fe³⁺ → add excess NaOH → no change, ppt does not dissolve

Two limits of this reagent should be noted before moving on. It stops short of separating zinc from lead, since both white precipitates dissolve in excess; and it treats copper, ferrous and ferric ions alike in excess, since none of the three dissolves. The next reagent settles the first of these.

Potassium hydroxide solution, caustic potash, behaves in exactly the same way throughout. It is also a strong alkali supplying hydroxide ions freely, so every precipitate described in this section forms with it as well, and the soluble products are then potassium zincate and potassium plumbite.

Extra Practice 4.2

A colourless salt solution gives a white precipitate when caustic potash solution is added drop by drop, and the precipitate dissolves when the alkali is added in excess. Name the two cations that could be present, and write the equation for the reaction of the precipitate with the excess alkali in each case.

4.4 Action Of Ammonium Hydroxide On Certain Salt Solutions

Ammonium hydroxide solution — the solution formed when ammonia gas dissolves in water — is the second great reagent of qualitative analysis. Added dropwise to a metallic salt solution, it does much the same work as caustic soda: it supplies hydroxide ions, and the metal hydroxides that are insoluble come down as coloured precipitates.

The interest lies in where the two reagents disagree. Ammonium hydroxide is a weak alkali, which means that only a small fraction of its molecules split into ions in water, so the concentration of hydroxide ions it provides is low. And in excess it does something caustic soda never does: it dissolves the copper precipitate, and turns the liquid deep blue.

Reagent Spotlight

Ammonium Hydroxide Ammonia dissolved in water, a weak alkali → supplies only a small concentration of hydroxide ions → precipitates most metal hydroxides but not calcium; in excess it dissolves the hydroxides of copper and zinc by forming soluble ammonia compounds.

4.4.1 Calcium Salts

With a calcium salt solution, ammonium hydroxide gives no precipitate at all, however much of it is added.

The reason is the weakness of the alkali. Calcium hydroxide is only sparingly soluble, so it needs a fairly high concentration of hydroxide ions before it will begin to come out of solution. Ammonium hydroxide ionises so little that the hydroxide-ion concentration never reaches that level, and the calcium ions stay in solution.

Curiosity Corner

Q3. Sodium hydroxide and ammonium hydroxide both supply hydroxide ions, yet only one of them will bring down a precipitate from a calcium salt. What must be different about the two solutions?

A. Their degree of ionisation. Ammonium hydroxide is a weak alkali and releases too few hydroxide ions to precipitate calcium hydroxide; caustic soda, a strong alkali, releases plenty.

Do Not Mix Up

Weak And Dilute Weak describes how far a substance ionises, and cannot be changed by adding water. Dilute describes how much water it is mixed with. A concentrated ammonium hydroxide solution is still a weak alkali.

4.4.2 Iron Salts

Iron behaves with ammonium hydroxide exactly as it did with caustic soda. Both hydroxides are precipitated, both keep their characteristic colours, and neither dissolves in excess of the reagent.

4.4.2.1 Iron(II) Salts

Green ferrous sulphate solution gives a dirty green precipitate of ferrous hydroxide, with ammonium sulphate left colourless in solution.

ferrous sulphate + ammonium hydroxide → ferrous hydroxide + ammonium sulphate

\[\mathrm{FeSO_4}\text{(aq)} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_2}\!\downarrow + \mathrm{(NH_4)_2SO_4}\text{(aq)}\]

The precipitate is insoluble in excess ammonium hydroxide.

4.4.2.2 Iron(III) Salts

Yellow ferric chloride solution gives the reddish brown precipitate of ferric hydroxide, with ammonium chloride in solution.

ferric chloride + ammonium hydroxide → ferric hydroxide + ammonium chloride

\[\mathrm{FeCl_3}\text{(aq)} + 3\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_3}\!\downarrow + 3\,\mathrm{NH_4Cl}\text{(aq)}\]

Ferric sulphate solution, also yellow, gives the same reddish brown hydroxide. Because each ferric sulphate unit carries two iron atoms, the equation needs six formula units of the alkali and yields two of the hydroxide.

ferric sulphate + ammonium hydroxide → ferric hydroxide + ammonium sulphate

\[\mathrm{Fe_2(SO_4)_3}\text{(aq)} + 6\,\mathrm{NH_4OH}\text{(aq)} \rightarrow 2\,\mathrm{Fe(OH)_3}\!\downarrow + 3\,\mathrm{(NH_4)_2SO_4}\text{(aq)}\]

Neither iron precipitate dissolves when the alkali is added in excess.

4.4.3 Copper(II) Salts And The Deep Blue Solution

Blue copper sulphate solution gives, at first, the same pale blue precipitate of copper hydroxide that caustic soda gave.

copper sulphate + ammonium hydroxide → copper hydroxide + ammonium sulphate

\[\mathrm{CuSO_4}\text{(aq)} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{Cu(OH)_2}\!\downarrow + \mathrm{(NH_4)_2SO_4}\text{(aq)}\]

Keep adding the reagent and the pale blue solid dissolves, and the liquid takes on a deep, intense blue quite unlike the blue of the original solution, as shown in Fig. 4.4-1. The copper has not simply gone back into solution unchanged. Four ammonia molecules have attached themselves to each copper ion, building a new, soluble ion which carries the deep blue colour.

ammonium sulphate present:

\[\mathrm{Cu(OH)_2} + \mathrm{(NH_4)_2SO_4} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{[Cu(NH_3)_4]SO_4}\text{(aq)} + 4\,\mathrm{H_2O}\text{(l)}\]

with excess ammonium hydroxide alone:

\[\mathrm{Cu(OH)_2} + 4\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{[Cu(NH_3)_4](OH)_2}\text{(aq)} + 4\,\mathrm{H_2O}\text{(l)}\]

The first product is tetraammine copper(II) sulphate and the second is tetraammine copper hydroxide. The square brackets hold together the group in which the four ammonia molecules are bound to the copper ion. This deep blue colour is a characteristic property of the \(\mathrm{Cu^{2+}}\) ion, and it is used to detect copper in qualitative analysis.

Diagram

Fig. 4.4-1 — Copper Sulphate Solution With Ammonium Hydroxide Three test tubes in a rack. Tube 1: clear blue copper sulphate solution, labelled ‘blue solution’. Tube 2: after a few drops of ammonium hydroxide, a pale blue precipitate suspended in the liquid, labelled ‘copper hydroxide, pale blue precipitate’. Tube 3: after excess ammonium hydroxide, no solid and a very dark blue clear solution, labelled ‘tetraammine copper(II) ion, deep blue solution’.

Curiosity Corner

Q5. A pale blue precipitate of copper hydroxide does not simply dissolve away in excess ammonia — the liquid turns a deep blue instead. What might the copper have become?

A. Part of a new ion. Four ammonia molecules bind to each copper ion to form the tetraammine copper ion, which is soluble and far more deeply blue than the copper ion in water.

Chemistry at Work

A Blue Liquid That Dissolves Cotton The deep blue solution formed from copper hydroxide and ammonia is known as Schweizer’s reagent, after the Swiss chemist who described it in the nineteenth century. It has an unusual property: it dissolves cellulose, the material of cotton and wood fibre. Forcing the resulting syrup through fine holes into an acid bath regenerates the cellulose as a long thread, and this was one of the earliest routes to a manufactured fibre.

4.4.4 Zinc Salts

Colourless zinc sulphate solution gives a white, gelatinous precipitate of zinc hydroxide.

zinc sulphate + ammonium hydroxide → zinc hydroxide + ammonium sulphate

\[\mathrm{ZnSO_4}\text{(aq)} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{Zn(OH)_2}\!\downarrow + \mathrm{(NH_4)_2SO_4}\text{(aq)}\]

In excess of the reagent this precipitate dissolves, and for the same reason as copper’s: ammonia molecules bind to the zinc ion and carry it into solution. The solution formed is colourless, since the zinc ion has no colour of its own.

\[\mathrm{Zn(OH)_2} + \mathrm{(NH_4)_2SO_4} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{[Zn(NH_3)_4]SO_4}\text{(aq)} + 4\,\mathrm{H_2O}\text{(l)}\]

or, with excess ammonium hydroxide alone,

\[\mathrm{Zn(OH)_2} + 4\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{[Zn(NH_3)_4](OH)_2}\text{(aq)} + 4\,\mathrm{H_2O}\text{(l)}\]

The products are tetraammine zinc(II) sulphate and tetraammine zinc hydroxide.

4.4.5 Lead Salts

Colourless lead nitrate solution gives a chalky white precipitate of lead hydroxide.

lead nitrate + ammonium hydroxide → lead hydroxide + ammonium nitrate

\[\mathrm{Pb(NO_3)_2}\text{(aq)} + 2\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{Pb(OH)_2}\!\downarrow + 2\,\mathrm{NH_4NO_3}\text{(aq)}\]

Here lead parts company with zinc. Lead hydroxide is insoluble in excess ammonium hydroxide, and the white precipitate remains however much reagent is added. This single difference is what makes the two reagents together so powerful.

4.4.6 Identifying A Cation With Ammonium Hydroxide

The whole scheme for this reagent is set out in FC 4.2, and the observations are collected in Table 4.5.

Flowchart

FC 4.2 — Identifying A Cation With Ammonium Hydroxide Solution Metal salt solution + a few drops of NH₄OH → observe; branch A: no precipitate at all → Ca²⁺; branch B: white ppt → add excess NH₄OH → no change → Pb²⁺, or ppt dissolves to a colourless solution → Zn²⁺; branch C: coloured ppt → pale blue → Cu²⁺ → add excess NH₄OH → ppt dissolves to a deep blue solution; dirty green → Fe²⁺ → ppt does not dissolve; reddish brown → Fe³⁺ → ppt does not dissolve

Compare Table

Table 4.5 — Action Of Ammonium Hydroxide On Some Salt Solutions

Cation in solution	Precipitate formed	Colour of precipitate	On adding excess ammonium hydroxide
\(\mathrm{Ca^{2+}}\)	None	—	No change; still no precipitate
\(\mathrm{Fe^{2+}}\)	Ferrous hydroxide	Dirty green	Insoluble; no change
\(\mathrm{Fe^{3+}}\)	Ferric hydroxide	Reddish brown	Insoluble; no change
\(\mathrm{Cu^{2+}}\)	Copper hydroxide	Pale blue	Dissolves, giving a deep blue solution
\(\mathrm{Zn^{2+}}\)	Zinc hydroxide	White, gelatinous	Dissolves, giving a colourless solution
\(\mathrm{Pb^{2+}}\)	Lead hydroxide	Chalky white	Insoluble; no change

Using The Two Reagents Together

Neither reagent settles every case alone, but between them they leave no doubt. Caustic soda separates calcium from zinc and lead, because calcium’s precipitate stays while the other two dissolve. Ammonium hydroxide then separates zinc from lead, because zinc’s precipitate dissolves in excess while lead’s does not. An analyst therefore takes two portions of the unknown solution and tests one with each reagent.

Example 4.2

Naming A Salt From Two Tests

1. A salt gives a colourless solution in water. With a little caustic soda solution a white gelatinous precipitate forms, which dissolves when more caustic soda is added. A fresh portion of the solution behaves in the same way with ammonium hydroxide. Identify the cation and write the equations for both reactions in excess, taking the salt to be the sulphate.

The colourless solution rules out the coloured cations, leaving calcium, zinc and lead among those studied.
The precipitate dissolves in excess caustic soda, which rules out calcium. It also dissolves in excess ammonium hydroxide, which rules out lead. Only zinc behaves in both ways.
The cation is the zinc ion, \(\mathrm{Zn^{2+}}\), and the two reactions in excess are

\[\mathrm{Zn(OH)_2} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]

\[\mathrm{Zn(OH)_2} + 4\,\mathrm{NH_4OH}\text{(aq)} \rightarrow \mathrm{[Zn(NH_3)_4](OH)_2}\text{(aq)} + 4\,\mathrm{H_2O}\text{(l)}\]

Extra Practice 4.3

You are given two bottles, one holding copper nitrate solution and the other ferric nitrate solution, and only ammonium hydroxide as a reagent. Describe what you would see in each case, first with a few drops and then with excess, and write the equation for the first precipitate formed in each.

4.5 Action Of Alkalis On Certain Metals

You already expect a metal to react with an acid and give off hydrogen. Alkalis are the opposite of acids, so it seems reasonable that a metal should be safe in them — and for most metals it is. Iron, copper and magnesium are not attacked by caustic soda solution.

Three metals break the pattern. Zinc, aluminium and lead react with hot concentrated caustic alkalis, sodium hydroxide or potassium hydroxide, giving the corresponding soluble salt and liberating hydrogen gas.

The reason is the one that runs through the rest of this chapter. The hydroxides of these three metals are the very ones that dissolved in excess alkali in Sections 4.3 and 4.4, forming zincate, aluminate and plumbite salts. Because those salts are stable and soluble in strong alkali, the alkali is able to attack the metal itself, taking the metal into the negative ion and releasing the hydrogen.

Do Not Mix Up

Metal With Acid And Metal With Alkali With an acid, the metal ends up in the positive ion — zinc chloride, zinc sulphate. With an alkali, it ends up inside the negative ion — sodium zincate. Hydrogen is given off in both cases.

4.5.1 Zinc

Zinc granules warmed with hot concentrated caustic soda solution give a colourless solution of sodium zincate and a steady stream of hydrogen. The set-up and the test for the gas are shown in Fig. 4.5-1.

zinc + sodium hydroxide → sodium zincate + hydrogen

\[\mathrm{Zn} + 2\,\mathrm{NaOH}\ \text{(hot and conc.)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + \mathrm{H_2}\!\uparrow\]

Caustic potash behaves in the same way, and the product is potassium zincate.

\[\mathrm{Zn} + 2\,\mathrm{KOH}\ \text{(hot and conc.)} \rightarrow \mathrm{K_2ZnO_2}\text{(aq)} + \mathrm{H_2}\!\uparrow\]

The gas is collected and a burning splint is brought to its mouth. It burns with a squeaky pop — the standard test that identifies hydrogen. Hot concentrated alkali is far more dangerous than the dilute solution used for precipitation, so the tube is heated gently, pointed away from everyone, and handled with tongs.

Diagram

Fig. 4.5-1 — Hydrogen From Zinc And Hot Concentrated Caustic Soda A boiling tube held in a clamp at a slight angle, containing zinc granules under hot concentrated sodium hydroxide solution, heated by a small burner flame; bubbles of gas rising through the liquid; a burning splint being brought to the mouth of the tube with a small burst of flame drawn beside it. Label the zinc granules, the hot concentrated caustic soda solution, the hydrogen bubbles and the burning splint.

Curiosity Corner

Q4. You already know that a metal reacting with an acid gives hydrogen. Why should zinc give hydrogen with a hot alkali as well, when most metals do not?

A. Because zinc’s hydroxide reacts with alkalis. The alkali can carry zinc away into a soluble zincate, in which zinc sits inside the negative ion, and the hydrogen displaced is set free.

4.5.2 Aluminium

Aluminium reacts with boiling caustic alkali solution. Here water is used up in the reaction as well as the alkali, and three molecules of hydrogen are released for every two atoms of aluminium.

aluminium + sodium hydroxide + water → sodium meta aluminate + hydrogen

\[2\,\mathrm{Al} + 2\,\mathrm{NaOH}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{NaAlO_2}\text{(aq)} + 3\,\mathrm{H_2}\!\uparrow\]

\[2\,\mathrm{Al} + 2\,\mathrm{KOH}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{KAlO_2}\text{(aq)} + 3\,\mathrm{H_2}\!\uparrow\]

The products are sodium meta aluminate and potassium meta aluminate. In ordinary use the shorter names sodium aluminate and potassium aluminate are the ones you will hear.

With fused alkali — solid caustic soda melted rather than dissolved — aluminium gives a different aluminate, again with hydrogen.

\[2\,\mathrm{Al} + 6\,\mathrm{NaOH}\ \text{(fused)} \rightarrow 2\,\mathrm{Na_3AlO_3} + 3\,\mathrm{H_2}\!\uparrow\]

Chemistry at Work

Why Aluminium Vessels Dislike Cleaning Powders The thin oxide film that protects an aluminium pan resists water, air and most foods, but it does not resist alkali. Strongly alkaline oven cleaners and washing powders eat through the film and then attack the metal, leaving the surface dull, pitted and grey. This is also why alkaline foods are not stored for long in bare aluminium, and why such vessels are often given a coating.

4.5.3 Lead

Lead reacts with caustic alkalis to give a plumbite and hydrogen.

lead + sodium hydroxide → sodium plumbite + hydrogen

\[\mathrm{Pb} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2PbO_2}\text{(aq)} + \mathrm{H_2}\!\uparrow\]

\[\mathrm{Pb} + 2\,\mathrm{KOH}\text{(aq)} \rightarrow \mathrm{K_2PbO_2}\text{(aq)} + \mathrm{H_2}\!\uparrow\]

Both plumbite solutions are colourless. Table 4.6 gathers the three metals together.

Compare Table

Table 4.6 — Metals That Liberate Hydrogen With Caustic Alkalis

Metal	Conditions	Salt formed with NaOH	Salt formed with KOH	Gas evolved
Zinc	Hot and concentrated alkali	Sodium zincate	Potassium zincate	Hydrogen
Aluminium	Boiling alkali solution, water also used	Sodium meta aluminate	Potassium meta aluminate	Hydrogen
Lead	Caustic alkali solution	Sodium plumbite	Potassium plumbite	Hydrogen

Example 4.3

A Metal Attacked By Both An Acid And An Alkali

1. A light metal dissolves in dilute hydrochloric acid, giving a gas that burns with a pop. A fresh piece of the same metal dissolves in boiling caustic soda solution, giving the same gas. Name the metal and write the balanced equation for each reaction.

A gas that burns with a pop is hydrogen, so the metal reacts with acid and with alkali alike. Of the metals studied, zinc, aluminium and lead do this, and of those aluminium is the light one.
With the acid, aluminium forms its chloride, in which aluminium is the positive ion:

\[2\,\mathrm{Al} + 6\,\mathrm{HCl}\text{(aq)} \rightarrow 2\,\mathrm{AlCl_3}\text{(aq)} + 3\,\mathrm{H_2}\!\uparrow\]

With the alkali, aluminium forms an aluminate, in which aluminium sits inside the negative ion, and water is used up as well:

\[2\,\mathrm{Al} + 2\,\mathrm{NaOH}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{NaAlO_2}\text{(aq)} + 3\,\mathrm{H_2}\!\uparrow\]

The metal is aluminium.

Extra Practice 4.4

Four metal strips — copper, magnesium, lead and iron — are each warmed with hot concentrated caustic potash solution. State which of them would give off a gas, name the gas and the test for it, and write the balanced equation for that reaction.

4.6 Action Of Alkalis On Metal Oxides And Hydroxides

The last section showed three metals being attacked by an alkali. Their oxides and hydroxides behave in the same unusual way, and this section sets that behaviour against the ordinary behaviour of metal oxides.

4.6.1 Basic Oxides And Hydroxides

Oxides of most metals are basic in nature. Many of them dissolve in water to form hydroxides, and where that hydroxide is soluble the solution obtained is an alkali. Sodium oxide is the plainest example.

sodium oxide + water → sodium hydroxide

\[\mathrm{Na_2O}\text{(s)} + \mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{NaOH}\text{(aq)}\]

A basic oxide, and the hydroxide formed from it, neutralise acids. Each reaction gives a salt and water, and nothing else.

\[\mathrm{Na_2O} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow 2\,\mathrm{NaCl}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]

\[\mathrm{NaOH}\text{(aq)} + \mathrm{HCl}\text{(aq)} \rightarrow \mathrm{NaCl}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]

What such oxides and hydroxides do not do is react with bases. A base put with another base has nothing to react with, since neutralisation needs an acid on one side and a base on the other. That is the general rule, and it holds for the great majority of metal oxides.

Do Not Mix Up

Base And Alkali Every alkali is a base, but not every base is an alkali. A base that dissolves in water to give hydroxide ions is called an alkali. Copper hydroxide is a base; it is not an alkali, because it is insoluble.

4.6.2 Amphoteric Oxides And Hydroxides

A few metallic oxides and hydroxides show a dual character. They react with acids, as a base does, and they also react with alkalis, as an acidic oxide does. Compounds of this kind are said to be amphoteric in nature.

Zinc oxide shows the two faces clearly. Put it with hydrochloric acid and it behaves as a base, giving zinc chloride and water.

\[\mathrm{ZnO} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]

\[\mathrm{Zn(OH)_2} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]

Put the same oxide with concentrated caustic soda and it behaves as an acidic oxide would, giving sodium zincate and water. The zinc has moved into the negative ion, exactly as it did in Sections 4.3 and 4.5.

\[\mathrm{ZnO} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]

\[\mathrm{Zn(OH)_2} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]

In every one of the four reactions the products are a salt and water. That is what the definition rests on.

Definition

Amphoteric Oxides And Hydroxides Amphoteric oxides and hydroxides are those compounds which react with both acids and alkalis to form salt and water.

The contrast with an ordinary basic oxide is set out in Table 4.7.

Compare Table

Table 4.7 — Basic Oxides Compared With Amphoteric Oxides

Basis	Basic oxide or hydroxide	Amphoteric oxide or hydroxide
Example	Sodium oxide, sodium hydroxide	Zinc oxide, aluminium hydroxide
Reaction with an acid	Reacts, giving salt and water	Reacts, giving salt and water
Reaction with an alkali	Does not react	Reacts, giving salt and water
Metal in the product	Always in the positive ion	In the positive ion with an acid, in the negative ion with an alkali

Diagram

Fig. 4.6-1 — Aluminium Hydroxide Reacting Both Ways Two test tubes side by side, each starting with the same white gelatinous solid labelled ‘freshly precipitated aluminium hydroxide’. Left tube: dilute hydrochloric acid being added, solid dissolving, product labelled ‘aluminium chloride solution’. Right tube: caustic soda solution being added, solid dissolving, product labelled ‘sodium meta aluminate solution’. Caption the pair ‘the same solid, two opposite reagents’.

Example 4.4

Showing That A Hydroxide Is Amphoteric

1. Freshly precipitated aluminium hydroxide dissolves both in dilute hydrochloric acid and in caustic soda solution. Write the balanced equation for each reaction, and state what these two results show.

With the acid, aluminium hydroxide acts as a base. Aluminium has a valency of three, so its chloride is \(\mathrm{AlCl_3}\) and three molecules of acid are needed:

\[\mathrm{Al(OH)_3} + 3\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{AlCl_3}\text{(aq)} + 3\,\mathrm{H_2O}\text{(l)}\]

With the alkali, it acts as an acidic oxide would, forming an aluminate:

\[\mathrm{Al(OH)_3} + \mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{NaAlO_2}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]

Both reactions give a salt and water. Aluminium hydroxide therefore reacts with an acid and with an alkali, which is exactly what makes a compound amphoteric.

4.6.3 Reactions Of Amphoteric Oxides And Hydroxides With Alkalis

Three metals of this chapter have amphoteric oxides and hydroxides: zinc, aluminium and lead. Their reactions with the two caustic alkalis are collected in Table 4.8. Notice the colours of the solids named there — zinc oxide and aluminium oxide are white, while lead oxide is yellow — and notice that every product solution is colourless.

Table 4.8 — Amphoteric Oxides And Hydroxides With Caustic Alkalis

Oxide or hydroxide	With sodium hydroxide	With potassium hydroxide
Zinc oxide (white)	\(\mathrm{ZnO} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + \mathrm{H_2O}\)	\(\mathrm{ZnO} + 2\,\mathrm{KOH} \rightarrow \mathrm{K_2ZnO_2} + \mathrm{H_2O}\)
Zinc hydroxide	\(\mathrm{Zn(OH)_2} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + 2\,\mathrm{H_2O}\)	\(\mathrm{Zn(OH)_2} + 2\,\mathrm{KOH} \rightarrow \mathrm{K_2ZnO_2} + 2\,\mathrm{H_2O}\)
Aluminium oxide (white)	\(\mathrm{Al_2O_3} + 2\,\mathrm{NaOH} \rightarrow 2\,\mathrm{NaAlO_2} + \mathrm{H_2O}\)	\(\mathrm{Al_2O_3} + 2\,\mathrm{KOH} \rightarrow 2\,\mathrm{KAlO_2} + \mathrm{H_2O}\)
Aluminium hydroxide	\(\mathrm{Al(OH)_3} + \mathrm{NaOH} \rightarrow \mathrm{NaAlO_2} + 2\,\mathrm{H_2O}\)	\(\mathrm{Al(OH)_3} + \mathrm{KOH} \rightarrow \mathrm{KAlO_2} + 2\,\mathrm{H_2O}\)
Lead oxide (yellow)	\(\mathrm{PbO} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2PbO_2} + \mathrm{H_2O}\)	\(\mathrm{PbO} + 2\,\mathrm{KOH} \rightarrow \mathrm{K_2PbO_2} + \mathrm{H_2O}\)
Lead hydroxide	\(\mathrm{Pb(OH)_2} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2PbO_2} + 2\,\mathrm{H_2O}\)	\(\mathrm{Pb(OH)_2} + 2\,\mathrm{KOH} \rightarrow \mathrm{K_2PbO_2} + 2\,\mathrm{H_2O}\)

Read down the table and one pattern stands out. The oxide and the hydroxide of a metal give the same salt; they differ only in how much water is produced, because the hydroxide brings extra hydrogen and oxygen of its own into the reaction.

Side Note

Yellow lead(II) oxide has an old trade name, litharge. It was long used to make the dense, sparkling glass of cut crystal, and to prepare the plates of lead–acid batteries.

Extra Practice 4.5

Sodium oxide and zinc oxide are both white solids. Describe a single test, using one reagent, that would tell them apart, state what you would observe in each case, and write the balanced equation for the reaction that occurs.

4.7 Summary

Analysis is the determination of the chemical components present in a given sample.
Qualitative analysis identifies the substances present; quantitative analysis determines how much of each is present.
A reagent is a substance used in a test because it reacts with another substance in a known way. Alkalis are important laboratory reagents for identifying metal ions.
Precipitation is the formation of an insoluble solid when two solutions are mixed; the solid formed is the precipitate. Only compounds insoluble in water form precipitates.
Salts of the representative elements, those of Groups 1, 2 and 13 to 17, are generally colourless; salts of the transition elements, Groups 3 to 12, are generally coloured.
Colourless cations include \(\mathrm{NH_4^{+}}\), \(\mathrm{Na^{+}}\), \(\mathrm{K^{+}}\), \(\mathrm{Ca^{2+}}\), \(\mathrm{Mg^{2+}}\), \(\mathrm{Al^{3+}}\), \(\mathrm{Pb^{2+}}\) and \(\mathrm{Zn^{2+}}\). Coloured cations include \(\mathrm{Cu^{2+}}\) blue, \(\mathrm{Fe^{2+}}\) light green, \(\mathrm{Fe^{3+}}\) yellowish brown, \(\mathrm{Ni^{2+}}\) green, \(\mathrm{Cr^{3+}}\) green and \(\mathrm{Mn^{2+}}\) pink.
Coloured anions include the permanganate ion \(\mathrm{MnO_4^{-}}\), pink or purple, the dichromate ion \(\mathrm{Cr_2O_7^{2-}}\), orange, and the chromate ion \(\mathrm{CrO_4^{2-}}\), yellow.
Sodium hydroxide solution is added drop by drop to a salt solution, because a precipitate that dissolves in excess would otherwise be missed.
With sodium hydroxide, calcium salts give a white precipitate, sparingly soluble and unchanged in excess.
Ferrous salts give a dirty green gelatinous precipitate and ferric salts a reddish brown one; both are insoluble in excess alkali.
Copper salts give a pale blue precipitate of copper hydroxide, insoluble in excess sodium hydroxide.
Zinc salts give a white gelatinous precipitate and lead salts a chalky white one; both dissolve in excess sodium hydroxide, forming sodium zincate and sodium plumbite. Lead hydroxide is also readily soluble in acetic acid.
\[\mathrm{Zn(OH)_2} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + 2\,\mathrm{H_2O}\]
An ammonium salt heated with sodium hydroxide, or any water-soluble base, gives off ammonia gas, which turns moist red litmus blue.
\[\mathrm{NH_4Cl} + \mathrm{NaOH} \xrightarrow{\ \Delta\ } \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{NH_3}\]
Potassium hydroxide solution behaves throughout like sodium hydroxide, giving potassium zincate and potassium plumbite.
Ammonium hydroxide is a weak alkali. It ionises so little that it gives no precipitate at all with a calcium salt.
With ammonium hydroxide, ferrous and ferric salts give the same dirty green and reddish brown precipitates, both insoluble in excess; lead salts give a chalky white precipitate, also insoluble in excess.
With excess ammonium hydroxide the pale blue copper hydroxide dissolves to a deep blue solution of a tetraammine copper compound. This is a characteristic test for the \(\mathrm{Cu^{2+}}\) ion.
Zinc hydroxide also dissolves in excess ammonium hydroxide, giving a colourless tetraammine zinc compound. Zinc and lead are told apart by this difference.
Zinc, aluminium and lead react with hot concentrated caustic alkalis to give a soluble salt and hydrogen, which burns with a pop.
\[\mathrm{Zn} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + \mathrm{H_2}\]
\[2\,\mathrm{Al} + 2\,\mathrm{NaOH} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaAlO_2} + 3\,\mathrm{H_2}\]
Oxides of most metals are basic. They neutralise acids to give salt and water, but do not react with bases.
Amphoteric oxides and hydroxides react with both acids and alkalis to form salt and water. The oxides and hydroxides of zinc, aluminium and lead are amphoteric.
\[\mathrm{ZnO} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2ZnO_2} + \mathrm{H_2O}\]
With an acid an amphoteric compound puts its metal into the positive ion; with an alkali it puts the same metal into the negative ion, as a zincate, an aluminate or a plumbite.

Flowchart

FC 4.3 — The Whole Identification Scheme Unknown salt solution → note its colour → divide into two portions; portion 1 + NaOH drop by drop, then in excess → no ppt, colourless throughout → NH₄⁺ (confirm by heating with NaOH and testing for ammonia); white ppt unchanged in excess → Ca²⁺; white ppt dissolving in excess → Zn²⁺ or Pb²⁺; pale blue ppt → Cu²⁺; dirty green ppt → Fe²⁺; reddish brown ppt → Fe³⁺ → portion 2 + NH₄OH in excess to settle the doubtful cases → white ppt dissolves → Zn²⁺; white ppt stays → Pb²⁺; pale blue ppt dissolves to deep blue solution → Cu²⁺ confirmed; no ppt at all → Ca²⁺ confirmed

4.8 Key Terms

Amphoteric — describing an oxide or hydroxide that reacts with both acids and alkalis to form a salt and water.
Analysis — the determination of the chemical components present in a given sample.
Precipitate — the insoluble solid that appears when two solutions are mixed.
Precipitation — the process by which an insoluble solid is formed when two solutions are mixed.
Qualitative analysis — the branch of analysis that identifies which substances or ions are present in a sample.
Quantitative analysis — the branch of analysis that determines how much of each component is present in a mixture.
Reagent — a substance used in a test because it reacts with another substance in a known, recognisable way.
Representative elements — the elements of Groups 1, 2 and 13 to 17, whose salts are generally colourless.
Sodium meta aluminate — the soluble salt \(\mathrm{NaAlO_2}\), formed when aluminium or its oxide or hydroxide reacts with caustic soda; commonly called sodium aluminate.
Sodium plumbite — the soluble salt \(\mathrm{Na_2PbO_2}\), formed when lead or its oxide or hydroxide reacts with caustic soda.
Sodium zincate — the soluble salt \(\mathrm{Na_2ZnO_2}\), formed when zinc or its oxide or hydroxide reacts with caustic soda.
Transition elements — the elements of Groups 3 to 12, whose salts are generally coloured.
Weak alkali — an alkali that ionises only slightly in water and therefore supplies a low concentration of hydroxide ions.

4.9 Points To Ponder

Every test in this chapter depends on something being insoluble. If all metal hydroxides were freely soluble in water, what would be left of qualitative analysis, and what other property of the ions might a chemist have to use instead?
Zinc, aluminium and lead behave like metals with acids and rather like non-metals with alkalis. Where in the periodic table would you expect to find elements that sit on such a boundary, and why should position and dual behaviour be connected?
Ammonium hydroxide fails to precipitate calcium hydroxide only because the concentration of hydroxide ions is low. Could a very concentrated ammonium hydroxide solution ever succeed, or is the failure built into the alkali itself?
Sodium hydroxide and ammonium hydroxide give identical results for iron but different results for copper and zinc. What does that suggest about which part of the reagent — the hydroxide ion or the ammonia — is doing the work in each case?
The tests of this chapter tell you which ion is present but say nothing about how much. What difficulty would a chemist face in trying to turn a precipitate into a measurement of quantity, and what would have to be controlled?
An ammonium salt gives no precipitate with an alkali, yet the test for it is still counted as part of the same scheme. Is a change you can smell weaker evidence than a change you can see, or simply a different kind of evidence?

4.10 Objective Questions

Multiple Choice

The salt solution that gives no precipitate with ammonium hydroxide, even in excess, is: [1]
1. zinc sulphate
2. calcium nitrate
3. lead nitrate
4. ferric chloride
A reddish brown precipitate, insoluble in excess sodium hydroxide, indicates the presence of: [1]
1. \(\mathrm{Fe^{2+}}\)
2. \(\mathrm{Cu^{2+}}\)
3. \(\mathrm{Fe^{3+}}\)
4. \(\mathrm{Pb^{2+}}\)
The compound formed when zinc hydroxide dissolves in excess caustic potash solution is: [1]
1. potassium zincate
2. sodium zincate
3. zinc chloride
4. tetraammine zinc hydroxide
Which of the following oxides will react with a caustic alkali as well as with a dilute acid? [1]
1. Sodium oxide
2. Copper oxide
3. Lead oxide
4. Magnesium oxide
Ammonium hydroxide added in excess to copper sulphate solution finally gives: [1]
1. a pale blue precipitate
2. a colourless solution
3. a deep blue solution
4. a dirty green precipitate
A colourless solution gives a white precipitate with sodium hydroxide which dissolves in excess, and a white precipitate with ammonium hydroxide which does not dissolve in excess. The cation present is: [1]
1. \(\mathrm{Zn^{2+}}\)
2. \(\mathrm{Ca^{2+}}\)
3. \(\mathrm{Pb^{2+}}\)
4. \(\mathrm{Al^{3+}}\)

Assertion–Reason

Choose your answer from the codes given below. (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): Ammonium hydroxide gives no precipitate with a calcium salt solution. Reason (R): Ammonium hydroxide is a weak alkali and supplies a low concentration of hydroxide ions. [1]
Assertion (A): Aluminium liberates hydrogen when boiled with caustic soda solution. Reason (R): Aluminium oxide and aluminium hydroxide are amphoteric. [1]
Assertion (A): Salts of the transition elements are generally coloured. Reason (R): Every salt of a Group 12 element gives a coloured solution. [1]

Match The Following

Match each observation in Column 1 with the correct cation in Column 2. [5]

Column 1	Column 2
a. White precipitate with NaOH, dissolving in excess; white precipitate with NH₄OH, not dissolving in excess	i. \(\mathrm{Fe^{2+}}\)
b. Pale blue precipitate dissolving in excess NH₄OH to a deep blue solution	ii. \(\mathrm{Zn^{2+}}\)
c. Dirty green gelatinous precipitate, insoluble in excess of either alkali	iii. \(\mathrm{Ca^{2+}}\)
d. White precipitate with NaOH, no change in excess; no precipitate at all with NH₄OH	iv. \(\mathrm{Cu^{2+}}\)
e. White precipitate dissolving in excess of both NaOH and NH₄OH	v. \(\mathrm{Pb^{2+}}\)

Fill In The Blanks

The precipitate of ……………. hydroxide is dirty green, while that of ……………. hydroxide is reddish brown. [2]
Aluminium oxide is ……………. (basic/amphoteric) in nature, whereas sodium oxide is ……………. (basic/amphoteric). [2]
Heating an ammonium salt with caustic soda solution liberates ……………. gas, which turns moist red litmus ……………. . [2]

True Or False

Copper hydroxide dissolves in excess sodium hydroxide solution. [1]
Lead hydroxide is readily soluble in acetic acid. [1]

Name The Following

The gas liberated when lead is warmed with caustic potash solution. [1]
The salt formed when aluminium hydroxide reacts with caustic soda solution. [1]
A coloured cation that gives a precipitate insoluble in excess of both sodium hydroxide and ammonium hydroxide. [1]

Give One Word Or Phrase For

The branch of analysis that answers the question ‘how much’. [1]
A compound that reacts with both an acid and an alkali to give a salt and water. [1]

Odd One Out

Zinc hydroxide, lead hydroxide, ferric hydroxide, aluminium hydroxide. Give a reason for your choice. [2]
Permanganate ion, sulphate ion, dichromate ion, chromate ion. Give a reason for your choice. [2]

4.11 Check Your Concepts

Define precipitation, and state the condition a compound must satisfy before it can be obtained as a precipitate. [2]
Give a reason why an alkali is added to a salt solution drop by drop rather than all at once. [2]
Name two coloured anions, and state the colour of each. [2]
Why does a ferrous salt give a pale green solution while a ferric salt of the same metal gives a yellow one? Answer in terms of the ion present. [2]
Differentiate between qualitative analysis and quantitative analysis, stating the basis of difference. [3]
Explain, with one example each, why copper hydroxide is called a base but not an alkali. [3]
What would you observe when caustic soda solution is added, first a little and then in excess, to lead nitrate solution? [3]
Name two bases that are not alkalis but that dissolve in a strong alkali, and state what makes them behave in this way. [3]
Ammonium hydroxide and sodium hydroxide behave alike towards iron salts but differently towards copper and zinc salts. Explain both parts of this statement. [4]
Describe how you would use only ammonium hydroxide solution to distinguish between zinc nitrate solution and lead nitrate solution. [4]
Most metals are unaffected by caustic alkalis, yet zinc, aluminium and lead are attacked by them. Explain why these three behave differently, and state what is formed in each case. [5]

4.12 Equations, Conversions & Identification

Write balanced equations for the reaction of sodium hydroxide solution with: [4]
1. ferric chloride solution
2. zinc sulphate solution, using excess alkali
3. copper sulphate solution
4. calcium nitrate solution
Complete and balance the following equations: [3]
1. \(\mathrm{Pb(OH)_2} + \mathrm{KOH} \rightarrow\)
2. \(\mathrm{Al_2O_3} + \mathrm{NaOH} \rightarrow\)
3. \(\mathrm{(NH_4)_2SO_4} + \mathrm{NaOH} \xrightarrow{\ \Delta\ }\)
Ammonium chloride is warmed with caustic soda solution. Name the gas evolved, state one test by which you would identify it, and write the balanced equation for the reaction. [3]
Zinc granules are warmed with hot concentrated caustic soda solution, as shown in Fig. 4.5-1. Name the labelled reagent in the tube and the gas collected, state its test, and write the balanced equation. [3]
State what you would see in each case when ammonium hydroxide solution is added, first dropwise and then in excess, to: [3]
1. ferrous sulphate solution
2. zinc sulphate solution
3. lead nitrate solution
How would you obtain, in one or more steps: [4]
1. sodium zincate from zinc sulphate solution
2. tetraammine copper(II) hydroxide from copper sulphate solution
A colourless salt solution gives a chalky white precipitate with sodium hydroxide solution, which dissolves when more alkali is added. The same precipitate, obtained separately, dissolves readily in acetic acid. Identify the cation, and write the equation for its reaction with excess caustic soda. [3]
Study the three test tubes drawn in Fig. 4.4-1. Name the reagent being added, name the precipitate seen in the second tube, and explain the change that produces the colour of the third tube. [3]
Give balanced equations to show that lead oxide is amphoteric, using dilute nitric acid and caustic soda solution, and name the products formed. [4]