Chapter 13: Practical Work
In a laboratory at Heidelberg, in the late 1850s, Robert Bunsen had a problem he could not solve with his eyes alone. He knew that a salt held in a hot, nearly colourless flame gives that flame a colour of its own, and that the colour tells you which metal the salt contains. It was a beautiful test, and an unreliable one. Traces of sodium are almost everywhere, and their strong yellow drowned every fainter colour.
His colleague Gustav Kirchhoff suggested passing the light of the flame through a prism before looking at it. Instead of one blended colour, the two men saw a set of sharp bright lines — and every element laid down its own pattern of lines. The yellow of sodium was now just two narrow lines, easy to recognise and easy to ignore while the rest of the pattern was read.
In 1860 they boiled down a large quantity of mineral water and found, in its flame, two blue lines that belonged to no element then known. They named the new metal caesium, from the Latin word for sky blue. A year later a pair of deep red lines gave them rubidium.
Two elements had been found without anyone first seeing, weighing or holding them. They were identified by what they did.
Diagram
Fig. 13.0-1 — Identifying A Metal By The Colour It Gives A Flame Two panels side by side. Left panel: a Bunsen burner with a non-luminous blue flame, a thin wire held in the edge of the flame carrying a small paste of a salt, and the flame coloured along the wire; label the burner, the air hole, the non-luminous flame and the wire with salt paste. Right panel: the same coloured light entering a prism and spreading on a screen into a dark band crossed by a few sharp bright lines; label the coloured light, the prism and the bright lines, with a caption line reading ‘each metal gives its own set of lines’.
13.1 Recognition And Identification Of Gases
A great deal of practical chemistry begins with a simple situation. Something is heated, or an acid is poured on a solid, and a gas comes off. You cannot see what it is, and you may not be able to smell it safely. What you can do is put the gas through a short series of tests and let it tell you its own name.
Every test in this section works the same way. You do something to the gas, you record exactly what you see, and only then do you say what the gas is. The observation always comes first; the conclusion follows from it.
Three preliminary observations cost nothing and are made before any reagent is used. First, the colour of the gas: almost all of the common gases are colourless, so a coloured gas has already narrowed the field. Second, the odour. Never put your nose over the mouth of a test tube. Hold it a little away, and use your hand to fan a small amount of the gas towards you — this is called wafting, and it is shown in Fig. 13.1-2.
Third, the effect of the gas on moist litmus paper. The paper must be moist, because a dry gas has no effect on dry litmus; the gas has to dissolve in the film of water before it can act as an acid or a base. An acidic gas turns moist blue litmus red, a basic gas turns moist red litmus blue, and a neutral gas leaves both unchanged.
Diagram
Fig. 13.1-1 — Testing A Gas As It Is Given Off A test tube held in a slanted clamp containing a small quantity of solid with a little dilute acid poured on it, bubbles rising in the liquid; a bent delivery tube leads from a stopper in the mouth of the test tube down into a second test tube containing lime water, with the end of the delivery tube dipping below the surface and a faint white cloudiness in the lime water. Label the solid, the dilute acid, the delivery tube, the lime water and the milkiness. Add a small inset showing a strip of moist blue litmus paper held just inside the mouth of a test tube of gas.
Diagram
Fig. 13.1-2 — The Safe Way To Smell A Gas A student standing to one side of a test tube of gas held at chest height, the free hand cupped above the mouth of the tube and fanning gently towards the face, the face turned slightly away and not directly over the tube. Label ‘never place the nose over the mouth of the tube’ and ‘fan a little gas towards you’.
After these three observations comes a confirmatory test — one reaction that only that gas gives, or a small set of reactions that together leave only one possible answer. The ten gases below are the ones you are expected to recognise.
13.1.1 Hydrogen
Hydrogen is prepared by adding dilute hydrochloric acid or dilute sulphuric acid to a reactive metal — that is, to a metal placed above hydrogen in the activity series, such as magnesium, zinc or iron. The metal takes the place of hydrogen in the acid, and the hydrogen is set free as a gas.
active metal + dilute acid → salt + hydrogen
\[\mathrm{Mg} + \mathrm{H_2SO_4} \rightarrow \mathrm{MgSO_4} + \mathrm{H_2}\]
\[\mathrm{Zn} + 2\,\mathrm{HCl} \rightarrow \mathrm{ZnCl_2} + \mathrm{H_2}\]
The steady stream of bubbles you see rising through the acid is called effervescence, and it is worth recording as an observation in its own right.
Recognising Hydrogen
- The gas is colourless, odourless and neutral to litmus.
- Pure hydrogen burns with a pale blue flame when a burning splint is brought to the mouth of the tube.
- A mixture of hydrogen and air burns with a pop sound when a burning taper is brought near it.
Hydrogen is flammable, so the burning-splint test is done on a small test tube of gas held away from the face, and never on a large jar.
13.1.2 Oxygen
Oxygen is obtained by heating either a higher metallic oxide or a metal nitrate. A higher oxide is one in which the metal holds more oxygen than it can hold at a high temperature, so on heating it gives up the surplus.
heavy metal oxide → metal oxide or metal + oxygen
\[2\,\mathrm{Pb_3O_4} \xrightarrow{\ \Delta\ } 6\,\mathrm{PbO} + \mathrm{O_2}\]
\[2\,\mathrm{HgO} \xrightarrow{\ \Delta\ } 2\,\mathrm{Hg} + \mathrm{O_2}\]
metal nitrate → metal oxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Cu(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{CuO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
\[2\,\mathrm{Zn(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Recognising Oxygen
- The gas is colourless, odourless and neutral to litmus.
- It rekindles a glowing splinter. A splint whose flame has just been blown out, so that its tip is still red hot, bursts back into flame in oxygen.
- It is absorbed by a colourless alkaline solution of pyrogallol, which turns dark brown.
13.1.3 Carbon Dioxide
There are two ordinary ways of getting carbon dioxide in the laboratory. The first is to heat a metallic carbonate — every carbonate except sodium carbonate and potassium carbonate, which are too stable to decompose at the temperature of a burner.
metallic carbonate → metal oxide + carbon dioxide
\[\mathrm{ZnCO_3} \xrightarrow{\ \Delta\ } \mathrm{ZnO} + \mathrm{CO_2}\]
\[\mathrm{CuCO_3} \xrightarrow{\ \Delta\ } \mathrm{CuO} + \mathrm{CO_2}\]
The second way is to add a dilute acid to any carbonate or hydrogen carbonate, which needs no heating at all.
metal carbonate or hydrogen carbonate + dilute acid → salt + water + carbon dioxide
\[\mathrm{CaCO_3} + 2\,\mathrm{HCl} \rightarrow \mathrm{CaCl_2} + \mathrm{H_2O} + \mathrm{CO_2}\]
\[\mathrm{NaHCO_3} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{CO_2}\]
Recognising Carbon Dioxide
- The gas is colourless and odourless.
- It turns moist blue litmus faint red — a weak effect, because carbon dioxide is only feebly acidic in water.
- Passed through lime water, it turns the lime water milky. The milkiness is a white precipitate of calcium carbonate, which is insoluble.
\[\mathrm{Ca(OH)_2} + \mathrm{CO_2} \rightarrow \mathrm{CaCO_3}\downarrow + \mathrm{H_2O}\]
If you keep passing the gas, the milkiness disappears again. The insoluble carbonate reacts with more carbon dioxide and water to give calcium hydrogen carbonate, which is soluble, so the solution turns clear once more.
\[\mathrm{CaCO_3} + \mathrm{CO_2} + \mathrm{H_2O} \rightarrow \mathrm{Ca(HCO_3)_2}\]
- The gas has no effect on filter paper dipped in acidified potassium dichromate or in acidified potassium permanganate. This negative result matters as much as the positive ones, as the next few pages will show.
13.1.4 Chlorine
Chlorine is prepared by adding concentrated hydrochloric acid to an oxidising agent such as red lead (\(\mathrm{Pb_3O_4}\)), lead dioxide (\(\mathrm{PbO_2}\)) or manganese dioxide (\(\mathrm{MnO_2}\)). The oxidising agent takes hydrogen from the acid as water and releases the chlorine.
oxidising agent + concentrated hydrochloric acid → metal chloride + water + chlorine
\[\mathrm{Pb_3O_4} + 8\,\mathrm{HCl} \rightarrow 3\,\mathrm{PbCl_2} + 4\,\mathrm{H_2O} + \mathrm{Cl_2}\]
\[\mathrm{MnO_2} + 4\,\mathrm{HCl} \rightarrow \mathrm{MnCl_2} + 2\,\mathrm{H_2O} + \mathrm{Cl_2}\]
Chlorine is poisonous and this preparation is carried out in a fume cupboard, never on an open bench.
Recognising Chlorine
- The gas is greenish yellow with a sharp, pungent, choking odour. Its colour alone separates it from every other gas in this list except nitrogen dioxide.
- It turns moist blue litmus paper red and then bleaches it, so that the paper ends up white. Bleaching means the destruction of a colour, and here it happens because chlorine and water give an acid together with nascent oxygen, written \([\mathrm{O}]\), which attacks the dye.
\[\mathrm{Cl_2} + \mathrm{H_2O} \rightarrow \mathrm{HCl} + \mathrm{HClO}\]
\[\mathrm{HClO} \rightarrow \mathrm{HCl} + [\mathrm{O}]\]
colouring matter + \([\mathrm{O}]\) → colourless or bleached product
- It turns moist starch iodide paper blue black. Chlorine displaces iodine from the iodide, and iodine with starch gives the deep blue-black colour.
\[\mathrm{Cl_2} + 2\,\mathrm{KI} \rightarrow 2\,\mathrm{KCl} + \mathrm{I_2}\]
starch + iodine → blue-black colour
- Passed through silver nitrate solution, it gives a white precipitate.
13.1.5 Hydrogen Chloride
Hydrogen chloride gas is made by warming a metal chloride with concentrated sulphuric acid. The acid is used concentrated because water would simply dissolve the gas as it formed.
metal chloride + concentrated sulphuric acid → metal hydrogen sulphate + hydrogen chloride
\[\mathrm{NaCl} + \mathrm{H_2SO_4} \xrightarrow{\ \Delta\ } \mathrm{NaHSO_4} + \mathrm{HCl}\]
\[\mathrm{KCl} + \mathrm{H_2SO_4} \xrightarrow{\ \Delta\ } \mathrm{KHSO_4} + \mathrm{HCl}\]
Recognising Hydrogen Chloride
- The gas is colourless with a pungent, choking odour.
- It turns moist blue litmus paper red, and does not bleach it.
- A glass rod dipped in ammonia solution and held near the gas produces dense white fumes of ammonium chloride. The fumes are solid particles of the salt, formed in the air between the rod and the gas.
\[\mathrm{NH_3} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl}\]
- Passed through silver nitrate solution, it gives a white precipitate of silver chloride.
\[\mathrm{AgNO_3} + \mathrm{HCl} \rightarrow \mathrm{AgCl}\downarrow + \mathrm{HNO_3}\]
The precipitate dissolves when excess ammonium hydroxide is added, because a soluble compound is formed.
\[\mathrm{AgCl} + 2\,\mathrm{NH_4OH} \rightarrow \mathrm{Ag(NH_3)_2Cl} + 2\,\mathrm{H_2O}\]
13.1.6 Sulphur Dioxide
Sulphur dioxide is obtained by adding dilute hydrochloric acid or dilute sulphuric acid to a metallic sulphite or hydrogen sulphite.
metal sulphite or hydrogen sulphite + dilute acid → salt + water + sulphur dioxide
\[\mathrm{Na_2SO_3} + 2\,\mathrm{HCl} \rightarrow 2\,\mathrm{NaCl} + \mathrm{H_2O} + \mathrm{SO_2}\]
\[\mathrm{Na_2SO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{Na_2SO_4} + \mathrm{H_2O} + \mathrm{SO_2}\]
\[2\,\mathrm{NaHSO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{Na_2SO_4} + 2\,\mathrm{H_2O} + 2\,\mathrm{SO_2}\]
Recognising Sulphur Dioxide
- The gas is colourless with the suffocating smell of burning sulphur.
- It turns moist blue litmus red and finally bleaches it.
- It turns lime water turbid, exactly as carbon dioxide does, because insoluble calcium sulphite is formed.
\[\mathrm{Ca(OH)_2} + \mathrm{SO_2} \rightarrow \mathrm{CaSO_3}\downarrow + \mathrm{H_2O}\]
- It decolourises pink potassium permanganate solution. This is the point at which sulphur dioxide parts company with carbon dioxide.
\[2\,\mathrm{KMnO_4} + 2\,\mathrm{H_2O} + 5\,\mathrm{SO_2} \rightarrow \mathrm{K_2SO_4} + 2\,\mathrm{MnSO_4} + 2\,\mathrm{H_2SO_4}\]
- It changes the orange or yellow colour of acidified potassium dichromate to green.
\[\mathrm{K_2Cr_2O_7} + \mathrm{H_2SO_4} + 3\,\mathrm{SO_2} \rightarrow \mathrm{Cr_2(SO_4)_3} + \mathrm{K_2SO_4} + \mathrm{H_2O}\]
- It has no effect on lead acetate paper, which is what separates it from hydrogen sulphide.
In both the permanganate and the dichromate tests, sulphur dioxide is acting as a reducing agent — it takes oxygen from the reagent, and the reagent’s colour changes as its metal is reduced. Carbon dioxide cannot do this, because the carbon in it is already fully oxidised.
13.1.7 Hydrogen Sulphide
Hydrogen sulphide is prepared by adding dilute hydrochloric acid or dilute sulphuric acid to a metallic sulphide such as iron(II) sulphide or zinc sulphide.
metal sulphide + dilute acid → salt + hydrogen sulphide
\[\mathrm{FeS} + \mathrm{H_2SO_4} \rightarrow \mathrm{FeSO_4} + \mathrm{H_2S}\]
\[\mathrm{ZnS} + 2\,\mathrm{HCl} \rightarrow \mathrm{ZnCl_2} + \mathrm{H_2S}\]
Recognising Hydrogen Sulphide
- The gas is colourless with the foul smell of rotten eggs.
- It turns moist blue litmus red.
- It turns lead acetate paper black, because black lead sulphide is deposited on the paper.
\[\mathrm{(CH_3COO)_2Pb} + \mathrm{H_2S} \rightarrow \mathrm{PbS}\downarrow + 2\,\mathrm{CH_3COOH}\]
- It turns lead nitrate solution black, for the same reason.
\[\mathrm{Pb(NO_3)_2} + \mathrm{H_2S} \rightarrow \mathrm{PbS}\downarrow + 2\,\mathrm{HNO_3}\]
13.1.8 Ammonia
Ammonia is obtained by warming an ammonium salt, such as ammonium chloride or ammonium sulphate, with an alkali. The alkali takes the hydrogen ion away from the ammonium ion and leaves ammonia behind.
ammonium salt + alkali → salt + water + ammonia
\[2\,\mathrm{NH_4Cl} + \mathrm{Ca(OH)_2} \xrightarrow{\ \Delta\ } \mathrm{CaCl_2} + 2\,\mathrm{H_2O} + 2\,\mathrm{NH_3}\]
\[\mathrm{NH_4Cl} + \mathrm{NaOH} \xrightarrow{\ \Delta\ } \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{NH_3}\]
Recognising Ammonia
- The gas is colourless with a sharp, pungent, characteristic smell.
- It turns moist red litmus blue. It is the only basic gas in this list, so this single observation is a strong clue.
- A glass rod dipped in concentrated hydrochloric acid and held near the gas gives dense white fumes of ammonium chloride.
- It turns colourless Nessler’s reagent — a solution of potassium tetraiodomercurate(II), \(\mathrm{K_2HgI_4}\) — brown. Nessler’s reagent is the standard reagent for ammonia and for the ammonium ion.
13.1.9 Water Vapour
Many crystalline salts hold water inside the crystal itself. This water is called the water of crystallisation, and a salt that contains it is a hydrated salt. On heating, the water is driven off as vapour and an anhydrous salt — a salt with no water of crystallisation — is left behind.
hydrated salt → anhydrous salt + water vapour
\[\mathrm{Na_2CO_3\!\cdot\!10H_2O} \xrightarrow{\ \Delta\ } \mathrm{Na_2CO_3} + 10\,\mathrm{H_2O}\]
\[\mathrm{CuSO_4\!\cdot\!5H_2O} \xrightarrow{\ \Delta\ } \mathrm{CuSO_4} + 5\,\mathrm{H_2O}\]
Blue hydrated copper sulphate turns white as it loses its water. The change of colour is itself an observation you should record.
Recognising Water Vapour
- The vapour is colourless and odourless, and it condenses as a clear liquid on the cooler upper parts of the test tube.
- The liquid is neutral to litmus.
- It turns anhydrous copper sulphate from white back to blue.
\[\mathrm{CuSO_4} + 5\,\mathrm{H_2O} \rightarrow \mathrm{CuSO_4\!\cdot\!5H_2O}\]
- It turns blue cobalt chloride paper pink.
\[\mathrm{CoCl_2} + 6\,\mathrm{H_2O} \rightarrow \mathrm{CoCl_2\!\cdot\!6H_2O}\]
Both of these tests show that water is present. Neither shows that the liquid is pure water, which is why the litmus test is done as well.
Diagram
Fig. 13.1-3 — The Test Papers Used In Gas Identification Four labelled strips of paper shown side by side, each drawn twice, before and after. (i) Blue litmus paper, before blue, after red. (ii) Red litmus paper, before red, after blue. (iii) Lead acetate paper, before white, after black. (iv) Cobalt chloride paper, before blue, after pink. Beneath each pair, print the gas that causes the change: an acidic gas, ammonia, hydrogen sulphide, water vapour. Add a fifth pair, starch iodide paper, before white, after blue-black, labelled chlorine. Note on the drawing that every paper except cobalt chloride paper is used moist.
13.1.10 Nitrogen Dioxide
Nitrogen dioxide is given off when a heavy metal nitrate, such as lead nitrate, is heated strongly.
metal nitrate → metal oxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Pb(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Copper nitrate and zinc nitrate behave in the same way, as their equations in section 13.1.2 show. Sodium nitrate and potassium nitrate do not produce nitrogen dioxide on heating.
Recognising Nitrogen Dioxide
- The gas is brown, with an irritating, pungent odour, and it does not burn.
- It turns moist blue litmus paper red.
- It turns starch iodide paper from colourless to blue black, and turns potassium iodide paper brown, because iodine is set free.
\[2\,\mathrm{KI} + 2\,\mathrm{NO_2} \rightarrow 2\,\mathrm{KNO_2} + \mathrm{I_2}\]
- It turns green acidified ferrous sulphate solution brown.
13.1.11 Physical Properties Of The Common Gases
Before any reagent comes out of the cupboard, colour, odour and action on litmus have already told you a great deal. Table 13.2 collects those three properties for all ten gases, and it is worth reading down each column rather than across each row.
Table 13.2 Colour, odour and nature of the common gases
| S. No. | Gas | Colour | Odour | Nature |
|---|---|---|---|---|
| 1. | Hydrogen (\(\mathrm{H_2}\)) | Colourless | Odourless | Neutral |
| 2. | Oxygen (\(\mathrm{O_2}\)) | Colourless | Odourless | Neutral |
| 3. | Carbon dioxide (\(\mathrm{CO_2}\)) | Colourless | Odourless | Acidic |
| 4. | Chlorine (\(\mathrm{Cl_2}\)) | Greenish yellow | Pungent, choking | Acidic |
| 5. | Hydrogen chloride (\(\mathrm{HCl}\)) | Colourless | Pungent, choking | Acidic |
| 6. | Sulphur dioxide (\(\mathrm{SO_2}\)) | Colourless | Burning sulphur | Acidic |
| 7. | Hydrogen sulphide (\(\mathrm{H_2S}\)) | Colourless | Rotten eggs | Acidic |
| 8. | Ammonia (\(\mathrm{NH_3}\)) | Colourless | Sharp pungent | Basic |
| 9. | Water vapour (\(\mathrm{H_2O}\)) | Colourless | Odourless | Neutral |
| 10. | Nitrogen dioxide (\(\mathrm{NO_2}\)) | Brown | Irritating | Acidic |
Only two of the ten are coloured, and only one is basic. That is why the order of testing matters: colour first, then litmus, then the confirmatory reagent. FC 13.1 sets out that order as a single scheme.
Flowchart
FC 13.1 — A Scheme For Identifying An Unknown Gas Note the colour of the gas → brown: nitrogen dioxide, confirm by turning starch iodide paper blue black; greenish yellow: chlorine, confirm by bleaching moist litmus; colourless: test with moist litmus → moist red litmus turns blue: ammonia, confirm with a glass rod dipped in concentrated hydrochloric acid; litmus unchanged: hydrogen, oxygen or water vapour → burning splint pops: hydrogen; glowing splinter rekindles: oxygen; blue cobalt chloride paper turns pink: water vapour; moist blue litmus turns red: carbon dioxide, sulphur dioxide, hydrogen sulphide or hydrogen chloride → lead acetate paper turns black: hydrogen sulphide; acidified dichromate turns green: sulphur dioxide; lime water milky but dichromate unaffected: carbon dioxide; dense white fumes with a rod dipped in ammonia solution: hydrogen chloride
13.2 Action Of Heat On Solids
Heating a solid in a dry test tube is the oldest test in analytical chemistry, and it is still one of the most informative. A small quantity of the substance is taken in a clean, dry test tube, held in a holder with the mouth pointing away from everyone, and heated gently at first and then strongly. Fig. 13.2-1 shows how it is done.
Four things are worth watching, and it helps to write them down in this order. What happens to the solid itself — does it melt, crackle, change colour? What condenses on the cooler upper part of the tube? What gas comes off, and what does it do to a test paper or to lime water? And what is left behind at the end. That last thing is called the residue, and its colour, both while it is hot and after it has cooled, is often the deciding clue.
Diagram
Fig. 13.2-1 — Heating A Solid In A Dry Test Tube A dry test tube held in a wooden test-tube holder at about forty-five degrees, mouth pointing away from the student and away from the neighbouring bench, containing a small quantity of solid at the closed end; a burner flame with an open air hole playing on the closed end. Show droplets of liquid condensed on the cooler upper wall of the tube and a wisp of gas leaving the mouth. Label the solid, the holder, the non-luminous flame, the condensed droplets, the escaping gas and the residue.
Heating is a dry test, and it destroys the portion of the sample used, so only a little is taken. The rest is kept for the tests that follow in later sections.
13.2.1 Action Of Heat On Carbonates
A carbonate that decomposes on heating gives a metal oxide and carbon dioxide. Both halves of that statement are checked separately — the gas by lime water, the residue by its colour.
Copper Carbonate
Copper carbonate is a light green amorphous powder. On strong heating it turns black.
- A colourless, odourless gas comes off, which puts out a burning wooden splinter.
- The gas turns lime water milky.
- The gas has no effect on acidified potassium dichromate or acidified potassium permanganate.
The gas is therefore carbon dioxide and not sulphur dioxide, and the black residue is copper oxide.
\[\mathrm{CuCO_3} \xrightarrow{\ \Delta\ } \mathrm{CuO} + \mathrm{CO_2}\]
Deduction: the light green powder is copper carbonate.
Zinc Carbonate
Zinc carbonate is a light amorphous white solid. On strong heating it changes to pale yellow.
- A colourless, odourless gas comes off which turns lime water milky, and the milkiness disappears when excess gas is passed.
- The gas has no effect on acidified potassium dichromate or acidified potassium permanganate.
- On cooling, the residue changes to a white colour.
\[\mathrm{ZnCO_3} \xrightarrow{\ \Delta\ } \mathrm{ZnO} + \mathrm{CO_2}\]
Deduction: the white powder is zinc carbonate, and the residue is zinc oxide, which is yellow when hot and white when cold.
13.2.2 Action Of Heat On Nitrates
A nitrate is a more complicated case, because a hydrated nitrate gives up its water first and its nitrate afterwards. You therefore see two stages, and as many as three gases in all.
Zinc Nitrate Hexahydrate
The solid is white, crystalline and deliquescent — that is, it takes up so much moisture from the air that it slowly dissolves in the water it has absorbed.
- On heating, it melts to a white sticky mass and gives off steamy vapours. These are water vapour, driven out of the crystal.
- On strong heating, the white residue gives off reddish brown fumes which turn moist blue litmus paper red.
- A glowing wooden splint held in the reddish brown gas bursts into flame, because oxygen is coming off mixed with the nitrogen dioxide.
- The residue is pale yellow while it is hot, and turns white on cooling.
\[\mathrm{Zn(NO_3)_2\!\cdot\!6H_2O} \xrightarrow{\ \Delta\ } \mathrm{Zn(NO_3)_2} + 6\,\mathrm{H_2O}\]
\[2\,\mathrm{Zn(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
The gases evolved are water vapour, nitrogen dioxide and oxygen; the residue is zinc oxide; and the substance is zinc nitrate hexahydrate.
Copper Nitrate Hexahydrate
The solid is a bluish green crystalline substance.
- On heating, it melts to a bluish green mass and gives off steamy vapours which condense as droplets of water on the cooler parts of the test tube.
- On further heating, the bluish green mass changes to a black residue, copper(II) oxide.
- A reddish brown gas is given off.
- A second gas is given off which rekindles a glowing splinter, so oxygen is present.
\[\mathrm{Cu(NO_3)_2\!\cdot\!6H_2O} \xrightarrow{\ \Delta\ } \mathrm{Cu(NO_3)_2} + 6\,\mathrm{H_2O}\]
The anhydrous salt then decomposes exactly as in the equation given in section 13.1.2, leaving copper oxide and giving off nitrogen dioxide and oxygen. The substance is copper(II) nitrate hexahydrate.
Lead Nitrate
Lead nitrate is a heavy white crystalline solid, and it is not hydrated, so no steamy vapours appear.
- On strong heating it crumbles with a crackling noise. This crackling of crystals as they are heated is called decrepitation, and it happens because the crystals split apart as they expand unevenly.
- A reddish brown gas is given off, which turns moist blue litmus paper red.
- A glowing wooden splinter held in the reddish brown gas relights, showing the presence of oxygen.
- The residue is reddish brown when hot; on cooling it changes to yellow, partly fuses into the glass and stains it yellow.
The equation is the one already written in section 13.1.10. The gases evolved are nitrogen dioxide and oxygen, the residue is lead(II) oxide, and the substance is lead(II) nitrate.
Diagram
Fig. 13.2-2 — Residues That Change Colour On Cooling Three pairs of small drawings, each pair showing the closed end of a test tube with the residue inside, first while hot and then after cooling. Pair one, zinc oxide: yellow when hot, white when cold. Pair two, lead(II) oxide: reddish brown when hot, yellow when cold, with the residue drawn sticking to and staining the glass. Pair three, copper(II) oxide: black when hot, black when cold. Label each pair with the name of the oxide and the words ‘hot’ and ‘cold’.
Table 13.3 sets the two families of solids side by side, since a great many identification questions come down to telling a carbonate from a nitrate.
13.2.3 General Rules For The Action Of Heat
The individual results above settle into a small number of rules. Each one is worth knowing as a rule and as a reason.
- An ammonium salt heated with an alkali, other than ammonium hydroxide, produces ammonia gas. Ammonium chloride heated by itself leaves no residue behind, because it passes straight into vapour and re-forms as a solid on the cool glass higher up the tube; ammonium nitrate also leaves no residue, and because it decomposes violently it is not heated in the school laboratory. Ammonium dichromate on heating leaves a greenish grey mass of chromium(III) oxide.
\[\mathrm{(NH_4)_2Cr_2O_7} \xrightarrow{\ \Delta\ } \mathrm{Cr_2O_3} + \mathrm{N_2} + 4\,\mathrm{H_2O}\]
Lead dioxide, red lead, mercury(II) oxide, potassium nitrate and sodium nitrate all produce oxygen gas on heating. Potassium nitrate and sodium nitrate give oxygen alone and no brown gas, because they lose only part of their oxygen and are left behind as nitrites.
Carbonates and hydrogen carbonates evolve carbon dioxide on heating, except potassium carbonate and sodium carbonate, which are stable at the temperature of a burner.
Hydrated salts produce water vapour on heating.
Sulphites, and some sulphates, produce sulphur dioxide when heated.
Lead compounds decompose to give lead monoxide, \(\mathrm{PbO}\), also called litharge. It is brown when hot and yellow when cold, and it sticks to the glass of the test tube.
Zinc compounds decompose on heating to give zinc oxide, which is yellow when hot and white when cold.
Copper compounds decompose to give black copper oxide.
Rules 6, 7 and 8 are the reason a residue is examined twice — once while it is still hot in the tube, and again after it has cooled. A single glance at a hot residue can send you to the wrong metal altogether.
13.3 Recognition Of Substances By Colour, Odour And Physical State
Before a single reagent is added, a salt has already told you something about itself. Its colour, its smell and the look and feel of the solid narrow the list of possibilities, and they cost nothing to observe.
These observations are called preliminary tests, and it is important to be clear about what they can and cannot do. A colour points to an ion; it does not prove it is there. That is why the tables in this section are headed ‘suspected ion’. The confirmation comes later, from the reactions in section 13.4.
13.3.1 Colour Of The Substance
Colour in a salt usually comes from the metal ion, and the same ion carries its colour into solution. Copper salts are blue or bluish green, iron(II) salts are light green, iron(III) salts are yellow or yellowish brown. Most of the other metal ions you meet at this level are colourless, so their salts are white solids that give colourless solutions.
Table 13.4 Colour of the substance and the cation suspected
| Colour observed | Suspected ion |
|---|---|
| Blue or bluish green | \(\mathrm{Cu^{2+}}\) |
| Light green | \(\mathrm{Fe^{2+}}\) |
| Yellow or yellowish brown | \(\mathrm{Fe^{3+}}\) |
| White (or colourless in solution) | \(\mathrm{Pb^{2+}}\), \(\mathrm{Zn^{2+}}\), \(\mathrm{Ca^{2+}}\), \(\mathrm{Na^{+}}\), \(\mathrm{K^{+}}\) or \(\mathrm{NH_4^{+}}\) |
The last row is the least helpful and the most common. A white salt has told you only which metals it is not — and that is still worth knowing, because it sends you straight to the alkali tests rather than to a colour test.
Diagram
Fig. 13.3-1 — Colours That Point To A Cation Four watch glasses in a row on a white tile, each holding a small heap of a crystalline salt, with a labelled test tube of its solution standing behind it. From left to right: a blue solid with a blue solution, labelled copper(II) ion; a light green solid with a pale green solution, labelled iron(II) ion; a yellowish brown solid with a yellow-brown solution, labelled iron(III) ion; a white solid with a colourless solution, labelled ‘colour gives no clue — go to the alkali tests’.
13.3.2 Odour Of The Substance
A few salts give away their anion by smell, either at once or after a drop of water is added. The smell comes from a gas escaping from the salt, so what you are really recognising is the gas of section 13.1. Smell is tested by wafting, never by holding the tube under the nose.
Table 13.5 Odour of the substance and the ion suspected
| Odour observed | Suspected ion |
|---|---|
| Smell of ammonia gas | \(\mathrm{NH_4^{+}}\) |
| Smell of hydrogen sulphide gas | \(\mathrm{S^{2-}}\) |
| Smell of sulphur dioxide gas | \(\mathrm{SO_3^{2-}}\) |
Notice that the ammonium ion is the one cation in this whole chapter that can be suspected by smell. Every other ion in the table is an anion.
13.3.3 Physical State And Nature Of The Substance
How the solid looks and behaves is the third free observation.
- An amorphous salt — a powder with no regular crystal shape — suggests the carbonate ion, \(\mathrm{CO_3^{2-}}\).
- A salt that is hygroscopic, meaning it absorbs moisture from the air and turns damp without dissolving, or one that is deliquescent, suggests the chloride ion, \(\mathrm{Cl^{-}}\), or the nitrate ion, \(\mathrm{NO_3^{-}}\).
Both of these are hints and nothing more. A damp bottle of a white powder is a reason to test for chloride and nitrate first, not a reason to write them down as the answer.
13.3.4 Solubility Of Salts And Bases In Water
Almost every wet test in the next section works by forming a precipitate. A precipitate appears when two solutions are mixed and one of the possible new compounds is insoluble in water. To predict that, or to explain it afterwards, you need to know which salts dissolve and which do not.
Table 13.6 gives this information for the ions you meet at this level. The Roman numeral after each ion is the valency of that radical or metal. Read down a column to learn a metal’s habits, and across a row to learn an anion’s.
Table 13.6 Solubility of salts and bases in water (Roman numerals indicate the valency of the radical or metal)
| Anion ↓ / Cation → | K (I) | Na (I) | Ba (II) | Ca (II) | Mg (II) | Al (III) | Zn (II) | Fe (II) | Fe (III) | Mn (II) | Pb (II) | Cu (II) | Ag (I) | Hg (II) |
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| \(\mathrm{NO_3}\) (I) | S | S | S | S | S | S | S | S | S | S | S | S | S | S |
| \(\mathrm{CH_3COO}\) (I) | S | S | S | S | S | Sp | S | S | S | S | S | S | S | S |
| \(\mathrm{Cl}\) (I) | S | S | S | S | S | S | S | S | S | S | N | S | N | N |
| \(\mathrm{SO_4}\) (II) | S | S | N | Sp | S | S | S | S | S | S | N | S | N | S |
| \(\mathrm{OH}\) (I) | S | S | S | Sp | Sp | N | N | N | N | N | N | N | — | — |
| \(\mathrm{S}\) (II) | S | S | N | N | N | — | N | N | — | N | N | N | N | N |
| \(\mathrm{SO_3}\) (II) | S | S | N | N | N | — | N | N | — | N | N | N | N | N |
| \(\mathrm{PO_4}\) (III) | S | S | N | N | N | N | N | N | N | N | N | N | N | N |
| \(\mathrm{CO_3}\) (II) | S | S | N | N | N | — | N | N | N | N | N | N | N | N |
Note: S — soluble; N — insoluble; — — does not exist; Sp — sparingly soluble.
Three patterns in this table do most of the work. Every nitrate is soluble, without exception. Every potassium and sodium salt is soluble, which is why these two metals never appear as a precipitate. And nearly every hydroxide, sulphide, sulphite, phosphate and carbonate is insoluble except those of potassium and sodium — which is exactly why adding an alkali to an unknown solution produces a precipitate you can look at.
13.4 Identification Of Ions
A salt is made of two parts, a cation and an anion, and each is identified by its own set of tests. Once you know both, you know the salt. This section is the heart of practical analysis, and it is worth learning as a habit rather than as a list.
13.4.1 Dry Tests And Wet Tests
A dry test is one carried out on the solid itself. It covers the colour of the substance, its density, its physical state, the action of dry heat on it, and the flame test of section 13.7. You have already met most of these.
A wet test is one in which a reagent is added to a solution of the substance, and the substance is identified from what you observe — most often a colour change or a precipitate.
Identification of ions is usually done in the solution state. There is a plain reason for this. Ions can only move, meet and react when the salt is dissolved and its ions are free; in a dry solid they are locked in place.
So the salt is first dissolved in water. If it does not dissolve in water, nitric acid is added instead. The acid converts the salt into its nitrate, and every nitrate is soluble in water, as Table 13.6 shows.
13.4.2 Identification Of Cations By The Action Of Alkalies
Cations are tested by the action of alkalies, because most metal hydroxides are insoluble and come down as coloured precipitates. The colour of the precipitate points to the metal, and what happens when more alkali is added often settles the matter. The reactions themselves were met in chapter 4.
How the alkali is added matters as much as which alkali it is. Add it slowly at first, one drop at a time, and watch the tube against a white background. If it is added too quickly it is easy to miss a precipitate that forms and then redissolves in the excess. Only after you have seen the first precipitate do you add the alkali in excess and record the change.
Diagram
Fig. 13.4-1 — Adding Alkali Drop By Drop And Then In Excess Three test tubes in a row against a white tile, each with a dropper above it. Tube 1: a clear solution with the first drop of alkali entering and a faint cloud forming where it lands, labelled ‘add one drop at a time’. Tube 2: a solid coloured precipitate settled through the liquid, labelled ‘note the colour of the precipitate’. Tube 3: the same tube after excess alkali has been added, the precipitate gone and the liquid clear, labelled ‘note whether it dissolves in excess’. Show a white tile behind all three.
Table 13.7 gives the results for the five metal ions you are expected to recognise. Read it in pairs: the colour first, then the behaviour in excess.
Table 13.7 Action of sodium hydroxide solution and of ammonium hydroxide solution on solutions of metal salts
| Metal ion | Colour of precipitate with sodium hydroxide | With excess sodium hydroxide | Colour of precipitate with ammonium hydroxide | With excess ammonium hydroxide |
|---|---|---|---|---|
| Copper | Pale blue | Insoluble | Pale blue | Soluble — deep blue solution |
| Iron(II) | Pale green, turning brown | Insoluble | Pale green, turning brown | Insoluble |
| Iron(III) | Reddish brown | Insoluble | Rust (brown) | Insoluble |
| Lead | White chalky | Soluble | White chalky | Insoluble |
| Zinc | White gelatinous | Soluble | White gelatinous | Soluble |
13.4.2.1 With Sodium Hydroxide Solution
Sodium hydroxide separates the coloured ions at once. A pale blue precipitate that stays put means copper. A reddish brown precipitate means iron(III). A pale green precipitate that slowly turns brown at the surface means iron(II) — the iron(II) hydroxide is being oxidised by the air to iron(III) hydroxide, which is why the brown colour starts at the top and works downwards.
The two white precipitates, lead and zinc, both dissolve when excess sodium hydroxide is added, so sodium hydroxide alone cannot tell them apart. Their appearance is a small clue: the lead precipitate is chalky, the zinc precipitate is gelatinous, like a jelly. A clue is not a confirmation.
13.4.2.2 With Ammonium Hydroxide Solution
Ammonium hydroxide is a weaker alkali, and this is exactly why a second test with it is worth doing. It separates the two cases that sodium hydroxide leaves undecided.
- With a copper salt, the pale blue precipitate dissolves in excess to give a deep inky blue solution. This is a confirmation of copper, and sodium hydroxide gives nothing like it.
- With a zinc salt, the white gelatinous precipitate dissolves in excess.
- With a lead salt, the white chalky precipitate does not dissolve in excess.
So a white precipitate that dissolves in excess of both alkalies is zinc, while one that dissolves in excess sodium hydroxide but not in excess ammonium hydroxide is lead.
13.4.2.3 Test For The Ammonium Ion
The ammonium ion is not a metal ion and gives no hydroxide precipitate at all, so it has its own two tests.
- When a caustic alkali, sodium hydroxide or potassium hydroxide, is added to any ammonium salt, ammonia gas is evolved on heating. The equation for this was written in section 13.1.8, and the gas is then identified by any of the tests given there.
- If Nessler’s reagent is added to a solution of an ammonium salt, it turns brown.
13.4.3 Identification Of Anions
Anions are tested by the reaction of acids on the salt, or by specific reagents added to its solution. The acid tests work by driving out a gas, which is then identified exactly as in section 13.1; the reagent tests work by forming a precipitate.
13.4.3.1 By Adding Dilute Sulphuric Acid
Take a small quantity of the salt in a test tube and add dilute sulphuric acid, warming it if there is no action in the cold. Three anions answer to this treatment.
Carbonate Ion
There is brisk effervescence, and a colourless, odourless gas is given off. It does not support combustion, so a burning splinter held in it is put out, and it turns moist blue litmus faint red. Passed through lime water, it turns the lime water milky; passed in excess, or on adding excess dilute nitric acid, the white precipitate dissolves to give a colourless solution. The gas is carbon dioxide, and the salt contains the carbonate ion, \(\mathrm{CO_3^{2-}}\).
Sulphide Ion
A gas smelling of rotten eggs is given off, which turns moist blue litmus paper red. A moist lead acetate paper brought near the gas turns black. The gas is hydrogen sulphide, and the salt contains the sulphide ion, \(\mathrm{S^{2-}}\).
Sulphite Ion
A gas with the suffocating odour of burning sulphur is given off, warming the tube if necessary. A filter paper moistened with acidified potassium dichromate and held near the gas changes from golden yellow or orange to green. The gas is sulphur dioxide, and the salt contains the sulphite ion, \(\mathrm{SO_3^{2-}}\).
Notice how the identification runs. The acid only releases the gas; it is the gas that names the anion, and the gas is named by the tests you already know.
13.4.3.2 By Adding Concentrated Sulphuric Acid
Two more anions need the concentrated acid, because a dilute acid would not displace their gases. Concentrated sulphuric acid is corrosive and gives out a great deal of heat with water, so it is added a few drops at a time, the tube is warmed gently, and it is never poured into a wet tube.
Chloride Ion
On warming the salt gently with concentrated sulphuric acid, a colourless gas with a pungent odour is evolved. The chloride ion may be present. Three confirmations are available.
- A glass rod carrying a drop of ammonia solution is brought near the gas: dense white fumes are produced. Chloride confirmed.
- A pinch of manganese dioxide is added to the salt, followed by concentrated sulphuric acid, and the mixture is heated. A greenish yellow gas with a pungent odour is evolved, which turns moist starch iodide paper blue black. The manganese dioxide has oxidised the hydrogen chloride to chlorine. Chloride confirmed.
- Silver nitrate solution is added to the salt solution. A white precipitate of silver chloride forms, which dissolves in excess ammonium hydroxide, as the equations in section 13.1.5 show. Chloride confirmed.
Nitrate Ion
On warming the salt gently with concentrated sulphuric acid, reddish brown fumes are evolved, and the fumes become thick when copper turnings are added. The gas is nitrogen dioxide, and the salt contains the nitrate ion, \(\mathrm{NO_3^{-}}\).
The confirmation is the brown ring test. To the salt solution, add freshly prepared ferrous sulphate solution. Then pour a few drops of concentrated sulphuric acid cautiously down the inside wall of the sloping test tube, so that the dense acid sinks and forms a separate layer underneath instead of mixing. A brown ring appears at the junction of the two liquid layers, and it disappears when the tube is shaken. Nitrate confirmed.
The ring forms only at the junction because that is the one place where all three substances meet. Shaking mixes the layers, the junction is destroyed, and the ring goes with it.
Diagram
Fig. 13.4-2 — The Brown Ring Test A test tube held slanting in a clamp. The upper layer is the salt solution mixed with freshly prepared ferrous sulphate solution; the lower, denser layer is concentrated sulphuric acid, drawn slightly darker. A dropper is shown touching the inside wall of the tube, with arrows showing the acid running down the wall rather than falling through the liquid. A distinct brown ring is drawn at the junction of the two layers. Label the salt solution with ferrous sulphate, the concentrated sulphuric acid layer, the wall of the tube, and the brown ring.
13.4.3.3 Test For The Sulphate Ion
The sulphate ion gives off no gas with either acid, so it is found by precipitation instead. There are two tests, and in both the added acid is there to dissolve away any carbonate or sulphite precipitate that would otherwise be mistaken for a sulphate.
- To the salt solution, add a little nitric acid and then barium chloride solution. A white precipitate is obtained which is insoluble in mineral acid. Sulphate confirmed.
- To the solution, add acetic acid and then lead acetate solution. A white precipitate is obtained which is soluble in excess of ammonium acetate solution. Sulphate confirmed.
13.4.3.4 General Identification Of Anions
Table 13.8 collects the anion tests into the order in which they are actually carried out at the bench. FC 13.2 then shows the same work as a single decision scheme.
Table 13.8 General identification of anions
| Add to the salt or salt solution | Observation | Anion present |
|---|---|---|
| 1. Dilute hydrochloric acid or dilute sulphuric acid, and heat | (i) Carbon dioxide gas is evolved | \(\mathrm{CO_3^{2-}}\) |
| (ii) Sulphur dioxide gas is evolved | \(\mathrm{SO_3^{2-}}\) | |
| (iii) Hydrogen sulphide gas is evolved | \(\mathrm{S^{2-}}\) | |
| 2. Concentrated sulphuric acid, and heat | (i) Hydrogen chloride gas is evolved | \(\mathrm{Cl^{-}}\) |
| (ii) Nitrogen dioxide gas is evolved | \(\mathrm{NO_3^{-}}\) | |
| 3. Silver nitrate solution | Curdy white precipitate, insoluble in dilute nitric acid but dissolving in ammonium hydroxide | \(\mathrm{Cl^{-}}\) |
| 4. Lead acetate solution | White precipitate insoluble in dilute hydrochloric acid and in dilute nitric acid | \(\mathrm{SO_4^{2-}}\) |
| 5. Barium nitrate solution or barium chloride solution | (i) White precipitate soluble in dilute hydrochloric acid or dilute nitric acid | \(\mathrm{CO_3^{2-}}\) or \(\mathrm{SO_3^{2-}}\) |
| (ii) White precipitate insoluble in dilute hydrochloric acid or dilute nitric acid | \(\mathrm{SO_4^{2-}}\) |
Flowchart
FC 13.2 — Deciding Which Anion Is Present Add dilute sulphuric acid to the salt, warming if there is no action in the cold → a gas is evolved → colourless and odourless, lime water milky, no effect on acidified dichromate: carbonate → rotten egg smell, lead acetate paper black: sulphide → burning sulphur smell, acidified dichromate paper turns green: sulphite; no gas with the dilute acid → warm the salt with concentrated sulphuric acid → colourless pungent gas, dense white fumes with a rod dipped in ammonia solution: chloride → reddish brown fumes thickening with copper turnings, brown ring test positive: nitrate; still no gas → test the solution with barium chloride after adding dilute nitric acid → white precipitate insoluble in the acid: sulphate
13.5 Identification Of Acidic And Basic Solutions
Two bottles stand on the bench, both holding a clear colourless liquid. One is a dilute acid, the other an alkali. Nothing about their appearance separates them, and neither may be tasted or touched. They are told apart by what they do to an indicator, and then by what they do to a chosen salt.
An indicator is a substance that shows one colour in an acidic solution and a different colour in an alkaline one. A small quantity of the liquid is taken on a spotting tile or in a test tube, and a drop of the indicator is added to it — the indicator is never added to the stock bottle.
Table 13.9 gives three indicators and two chemical tests. Two tests for each solution are enough to be sure, and the two should not both be indicators.
Diagram
Fig. 13.5-1 — What Each Indicator Shows Three pairs of test tubes standing in a rack against a white card, each pair labelled with the name of the indicator and with the left tube marked ‘acid’ and the right tube marked ‘alkali’. Pair one, litmus: left tube red, right tube blue. Pair two, methyl orange: left tube pink, right tube yellow. Pair three, phenolphthalein: left tube colourless, right tube pink. Draw a dropper above the first tube to show that the indicator is added to a small portion, not to the stock bottle.
Why The Two Chemical Tests Work
The indicator tests tell you what the solution is. The two chemical tests show it doing what an acid or an alkali does, and that is why they count as independent evidence.
Sodium carbonate is the salt of a weak acid, so any stronger acid drives carbon dioxide out of it. The effervescence, and the milkiness the gas gives with lime water, together identify the liquid as an acid.
\[\mathrm{Na_2CO_3} + 2\,\mathrm{HCl} \rightarrow 2\,\mathrm{NaCl} + \mathrm{H_2O} + \mathrm{CO_2}\]
Ammonium carbonate works the other way round. It is the salt of a weak base, ammonia, so an alkali displaces that ammonia and the gas comes off. An acid can do nothing of the kind.
\[\mathrm{(NH_4)_2CO_3} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2CO_3} + 2\,\mathrm{H_2O} + 2\,\mathrm{NH_3}\]
The ammonia is recognised in the ordinary way, by turning moist red litmus blue and by the dense white fumes it gives with a glass rod dipped in hydrochloric acid.
13.6 Action Of Concentrated Hydrochloric Acid On Copper Oxide And Manganese Dioxide
Copper oxide and manganese dioxide are both black powders. Nothing in their colour, and nothing in the heating test, separates them. What separates them is the way each one behaves towards concentrated hydrochloric acid, because the two oxides are chemically different kinds of substance.
Copper oxide is a basic oxide. It reacts with an acid in the ordinary way, giving a salt and water, and no gas at all.
\[\mathrm{CuO} + 2\,\mathrm{HCl} \rightarrow \mathrm{CuCl_2} + \mathrm{H_2O}\]
Manganese dioxide is an oxidising agent. It takes hydrogen from the acid to form water and sets the chlorine free, as the equation in section 13.1.4 shows. That chlorine is what you see and smell.
The experiment is carried out in three steps, and both powders are put through the same three steps side by side. Chlorine is poisonous and concentrated hydrochloric acid gives off choking fumes, so the work is done in a fume cupboard and the tubes are warmed gently, never boiled.
Diagram
Fig. 13.6-1 — Comparing The Two Black Powders Three panels in a row inside an outlined fume cupboard. Panel (a): two test tubes side by side in a rack, each holding a black powder with concentrated hydrochloric acid being warmed by a small flame; over the left tube draw a greenish yellow gas with a strip of moist starch iodide paper turning blue black, over the right tube draw no gas. Label the tubes ‘manganese dioxide’ and ‘copper oxide’. Panel (b): each mixture being poured through a filter funnel and filter paper into a beaker; label the left filtrate ‘brownish’ and the right filtrate ‘bluish’. Panel (c): a dropper adding ammonium hydroxide to each filtrate; left beaker unchanged and labelled ‘no precipitate’, right beaker showing a pale blue precipitate that becomes a deep blue solution when more is added.
What Is Observed At Each Step
- Concentrated hydrochloric acid is added to the black powder and the mixture is heated. With manganese dioxide, a greenish yellow gas is evolved, which is chlorine. With copper oxide, no chlorine is evolved.
- The mixture is filtered. The filtrate from manganese dioxide is brownish in colour; the filtrate from copper oxide is bluish.
- Ammonium hydroxide is added to the filtrate. With the manganese dioxide filtrate, no precipitate is formed. With the copper oxide filtrate, a pale blue precipitate is formed, which dissolves in excess ammonium hydroxide to give an azure, deep blue solution.
The third step is the one that names the metal. A pale blue precipitate turning to a deep blue solution in excess ammonium hydroxide is the test for copper you met in section 13.4.2.2, and here it is doing exactly the same work on a filtrate instead of on a salt solution.
13.7 The Flame Test
Some metals give a burner flame a colour of their own, and that colour can be used to name the metal. This is the flame test, and it is a dry test — it is done on the solid, not on a solution.
The test is carried out on the chloride of the metal, not on whatever salt you happen to have. Concentrated hydrochloric acid is added to a little of the salt to make a paste, which converts the salt into its chloride. Chlorides of metals generally vaporise and ionise readily when heated, and it is this that lets the metal reach the flame and colour it.
How The Test Is Carried Out
A thin platinum wire is used, because platinum does not react with the acid, withstands the temperature of the flame, and gives no colour of its own.
- The wire is first cleaned thoroughly by dipping it into concentrated hydrochloric acid.
- It is then heated in the non-luminous flame of the burner. When the wire imparts no colour at all to the flame, it is clean and ready for use.
- A little of the paste is now taken on the wire.
- The wire is introduced into the non-luminous part of the flame, and the colour imparted to the flame is observed.
The flame must be the non-luminous one, obtained with the air hole of the burner open. A luminous flame is yellow to begin with, and a yellow flame cannot show you a yellow result.
Diagram
Fig. 13.7-1 — Carrying Out A Flame Test Three stages drawn left to right. Stage 1: a thin platinum wire fixed in a glass holder being dipped into a small watch glass of concentrated hydrochloric acid; label ‘clean the wire’. Stage 2: the same wire held in the edge of a non-luminous blue flame of a Bunsen burner with the air hole shown open, the flame uncoloured; label ‘heat until the wire gives no colour’. Stage 3: the wire carrying a small paste of the salt, held in the non-luminous part of the flame, which is now coloured; label ‘introduce the paste and watch the colour’. Show a second watch glass of the salt paste beside stage 3.
The Colours And What They Mean
Table 13.11 gives the four colours you are expected to recognise. The middle column is the same flame viewed through a piece of blue cobalt glass, and that column is more useful than it looks.
Table 13.11 Colour imparted to the flame by salts of different metals
| Colour imparted to the flame (observed with the naked eye) | Colour seen through the blue glass | Name of metal |
|---|---|---|
| Golden yellow (persistent) | Yellow colour vanishes | Sodium ion (\(\mathrm{Na^{+}}\)) |
| Violet (lilac) | Violet or pink | Potassium ion (\(\mathrm{K^{+}}\)) |
| Brick red | Light green | Calcium ion (\(\mathrm{Ca^{2+}}\)) |
| Peacock bluish green | Bluish green | Copper ion (\(\mathrm{Cu^{2+}}\)) |
Sodium is the difficulty in this test, and the word ‘persistent’ in the first row is the warning. Traces of sodium are present almost everywhere — on the fingers, in the dust, in ordinary glass — and their strong golden yellow lasts and lasts, hiding fainter colours underneath it. Blue cobalt glass absorbs that yellow light almost completely. Looking through it, the sodium colour vanishes and the lilac of potassium, or the red of calcium, can be seen for what it is.
Diagram
Fig. 13.7-2 — The Four Flame Colours Four burner flames drawn side by side against a dark background, each with the platinum wire and its paste at the base of the flame. From left to right: a golden yellow flame labelled sodium salt; a lilac or violet flame labelled potassium salt; a brick red flame labelled calcium salt; a bluish green flame labelled copper salt. Beneath each flame print the name of the metal ion. Add a small drawing of a rectangle of blue cobalt glass held in front of the first two flames, with the sodium flame shown greyed out through it and the potassium flame shown clearly.
13.8 The Hydrogen Ion Scale
The tests of section 13.5 tell you whether a solution is acidic or basic. They do not tell you how acidic or how basic it is, and there is a great difference between the acid in a lemon and the acid in a car battery.
Acidic and basic strengths are compared on the hydrogen ion scale, also called the pH scale, which gives a solution a number called its pH value. The scale is named after the hydrogen ion because it is the hydrogen ion released in solution that makes a solution acidic; the more of it there is, the lower the pH value.
A solution of pH 7 is neutral, and pure water is the standard example. A value of less than 7 means the solution is acidic, and a value of more than 7 means it is basic. Fig. 13.8-1 shows the whole scale.
Diagram
Fig. 13.8-1 — The pH Scale A horizontal line marked with equal divisions and numbered 0 to 14, with a tick and label at every whole number. A downward arrow at 7 labelled NEUTRAL. To the left of 7, a long arrow pointing left labelled INCREASING ACIDIC NATURE. To the right of 7, a long arrow pointing right labelled INCREASING BASIC NATURE.
Two things about the scale are worth fixing in mind. It runs in the direction opposite to what you might expect — the smaller the number, the stronger the acid. And 7 is not the bottom of the scale but the middle of it, the point at which a solution is neither acidic nor basic.
13.9 Indicators
An indicator answers a question about a solution by changing colour in it. The tests in section 13.5 used three of them; this section sets out their colours in full and then introduces an indicator that does more than say acid or alkali.
13.9.1 Common Indicators And Their Colour Changes
Table 13.12 gives the colour of each indicator in a neutral, an acidic and an alkaline solution. It is worth reading the neutral column first, because that is the colour you start from.
Table 13.12 Colours of common indicators
| Indicator | Colour in neutral solution | Colour in acidic solution | Colour in alkaline solution |
|---|---|---|---|
| Litmus | Purple | Red | Blue |
| Methyl orange | Orange | Pink/Red | Yellow |
| Phenolphthalein | Colourless | Colourless | Pink |
| Alkaline phenolphthalein | Pink | Colourless | Pink |
Phenolphthalein has an awkward habit: it is colourless in a neutral solution and colourless in an acid too, so it cannot tell those two apart. Alkaline phenolphthalein is the same substance already turned pink by an alkali, and it is used for the opposite job — it loses its pink colour when an acid is added, and that disappearance is easy to see.
Litmus is the most generous of the three, since it gives a definite and different colour in all three kinds of solution.
13.9.2 The Universal Indicator
Each of the indicators above gives one colour change at one point on the pH scale, so each can only sort solutions into two or three groups. A universal indicator is a mixture, and it gives a whole range of colours with solutions of different pH values.
- In a neutral solution, at pH 7, a universal indicator produces a green colour.
- In a basic solution the colour changes from blue to indigo and then to violet as the pH increases from 7 to 14.
- In an acidic solution the colour changes from yellow to pink and then to red as the pH decreases from 7 to 1.
Fig. 13.9-1 shows the full range of colours against the scale, with familiar substances placed at their pH values.
Diagram
Fig. 13.9-1 — Colours Of Universal Indicator At Different pH Values A vertical colour strip divided into fifteen bands, labelled pH = 0 at the top down to pH = 14 at the bottom. The bands run red at pH 0–1, through orange and yellow to pH 5–6, green at pH 7, then blue at pH 8–10, indigo at pH 11–12 and violet at pH 13–14. To the right of the strip, tie labels to their bands: hydrochloric acid near pH 0–1, vinegar near pH 3, coffee near pH 5, pure water at pH 7, hand soap near pH 10, ammonia solution near pH 12, sodium hydroxide near pH 14. To the left of the strip, label gastric acid near pH 1, tomato juice near pH 4, milk near pH 6 and ocean water near pH 8.
Finding The pH Of A Substance
Make a solution of the substance to be tested and put a drop of it on a strip of pH paper — paper soaked in universal indicator and dried. The paper takes up a colour, and that colour is matched against the chart of colours supplied by the manufacturer, which gives the pH value.
Try this with hydrochloric acid, distilled water, sodium hydroxide solution and ammonium hydroxide on four different strips. You will see four different shades, and the order in which they fall along the chart is itself the answer to which solution is the most acidic and which the most basic.
13.10 Summary
In practical work the observation always comes before the conclusion. A gas is recorded as colourless, odourless and neutral to litmus first; only then is it named.
A gas is smelt by wafting a little of it towards the face, never by holding the tube under the nose, and litmus paper is used moist because a dry gas has no effect on dry litmus.
Hydrogen is prepared by adding dilute hydrochloric acid or dilute sulphuric acid to a metal above hydrogen in the activity series. It is colourless, odourless and neutral, burns with a pale blue flame, and burns with a pop sound when mixed with air.
Oxygen is obtained by heating a higher metallic oxide or a metal nitrate. It is neutral, rekindles a glowing splinter, and is absorbed by alkaline pyrogallol solution, which turns dark brown.
Carbon dioxide is obtained by heating a metallic carbonate other than sodium or potassium carbonate, or by adding a dilute acid to any carbonate or hydrogen carbonate. It turns moist blue litmus faint red and turns lime water milky.
\[\mathrm{Ca(OH)_2} + \mathrm{CO_2} \rightarrow \mathrm{CaCO_3}\downarrow + \mathrm{H_2O}\]
The milkiness disappears when excess carbon dioxide is passed, because soluble calcium hydrogen carbonate is formed.
Chlorine is prepared by adding concentrated hydrochloric acid to an oxidising agent such as red lead, lead dioxide or manganese dioxide. It is greenish yellow, pungent, reddens and then bleaches moist litmus, and turns moist starch iodide paper blue black.
Hydrogen chloride is prepared by warming a metal chloride with concentrated sulphuric acid. It is colourless and pungent, reddens moist litmus without bleaching it, and gives dense white fumes with a rod dipped in ammonia solution.
Sulphur dioxide is obtained by adding a dilute acid to a sulphite or hydrogen sulphite. It smells of burning sulphur, turns lime water milky like carbon dioxide, and — unlike carbon dioxide — decolourises pink potassium permanganate and turns acidified potassium dichromate from orange to green.
Hydrogen sulphide is obtained by adding a dilute acid to a metallic sulphide. It smells of rotten eggs and turns lead acetate paper and lead nitrate solution black.
Ammonia is obtained by warming an ammonium salt with an alkali. It is the only basic gas in the list: it turns moist red litmus blue, gives dense white fumes with a rod dipped in concentrated hydrochloric acid, and turns Nessler’s reagent brown.
Water vapour is obtained by heating a hydrated salt. It condenses as a neutral liquid, turns anhydrous copper sulphate from white to blue, and turns blue cobalt chloride paper pink.
Nitrogen dioxide is obtained by heating a heavy metal nitrate. It is brown and irritating, reddens moist litmus, turns starch iodide paper blue black, and turns green acidified ferrous sulphate solution brown.
Of the ten gases, only chlorine and nitrogen dioxide are coloured and only ammonia is basic, so colour and litmus are always the first two observations.
On heating, a carbonate gives one colourless gas, carbon dioxide, and leaves the metal oxide.
\[\mathrm{ZnCO_3} \xrightarrow{\ \Delta\ } \mathrm{ZnO} + \mathrm{CO_2}\]
- On heating, a hydrated nitrate gives water vapour first, then nitrogen dioxide and oxygen, and leaves the metal oxide.
\[2\,\mathrm{Zn(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Lead nitrate decrepitates on heating; zinc oxide is yellow when hot and white when cold; lead monoxide is brown when hot and yellow when cold; copper oxide is black throughout. A residue is therefore examined both hot and cold.
Ammonium salts heated with an alkali give ammonia; ammonium chloride and ammonium nitrate leave no residue; carbonates and hydrogen carbonates give carbon dioxide except those of sodium and potassium; hydrated salts give water vapour; sulphites and some sulphates give sulphur dioxide.
Colour suggests the cation: blue or bluish green suggests copper, light green iron(II), yellowish brown iron(III), and white or colourless leaves lead, zinc, calcium, sodium, potassium or ammonium possible.
Odour suggests the anion or the ammonium ion, and an amorphous powder suggests a carbonate, while a hygroscopic or deliquescent solid suggests a chloride or a nitrate. These are hints, not proofs.
All nitrates, and all sodium and potassium salts, are soluble in water. Most hydroxides, sulphides, sulphites, phosphates and carbonates are insoluble, which is why precipitation tests work.
A dry test uses the solid itself — colour, physical state, dry heating, flame test. A wet test uses reagents added to a solution. Ions are usually identified in the solution state.
Cations are identified by adding an alkali drop by drop and then in excess. Copper gives a pale blue precipitate, iron(II) a pale green one turning brown, iron(III) a reddish brown one, and lead and zinc white precipitates.
Sodium hydroxide dissolves both the lead and the zinc precipitates in excess; ammonium hydroxide dissolves the zinc one and the copper one but not the lead one. Using both alkalies separates all five ions.
The ammonium ion gives ammonia when warmed with sodium or potassium hydroxide, and turns Nessler’s reagent brown.
Dilute sulphuric acid identifies three anions by the gas evolved: carbon dioxide for a carbonate, hydrogen sulphide for a sulphide, sulphur dioxide for a sulphite.
Concentrated sulphuric acid identifies two more: hydrogen chloride for a chloride, confirmed by silver nitrate solution, and nitrogen dioxide for a nitrate, confirmed by the brown ring test with freshly prepared ferrous sulphate solution.
A sulphate gives no gas. It is confirmed by a white precipitate with barium chloride after acidifying with nitric acid, insoluble in mineral acid, or by a white precipitate with lead acetate that dissolves in excess ammonium acetate solution.
\[\mathrm{Na_2SO_4} + \mathrm{BaCl_2} \rightarrow \mathrm{BaSO_4}\downarrow + 2\,\mathrm{NaCl}\]
An acidic solution reddens blue litmus, turns methyl orange pink and leaves phenolphthalein colourless, and gives carbon dioxide with sodium carbonate. An alkali does the reverse and gives ammonia with ammonium carbonate.
Concentrated hydrochloric acid separates the two black powders: manganese dioxide gives chlorine and a brownish filtrate, copper oxide gives no chlorine and a bluish filtrate in which ammonium hydroxide produces a pale blue precipitate soluble in excess.
\[\mathrm{CuO} + 2\,\mathrm{HCl} \rightarrow \mathrm{CuCl_2} + \mathrm{H_2O}\]
In a flame test the salt is made into a paste with concentrated hydrochloric acid and held on a clean platinum wire in the non-luminous flame. Sodium gives golden yellow, potassium violet, calcium brick red and copper peacock bluish green.
Blue glass absorbs the yellow of sodium, so potassium’s lilac can be seen through it even when sodium is present.
On the hydrogen ion scale, pH 7 is neutral, below 7 is acidic and above 7 is basic.
Litmus is purple, red and blue in neutral, acidic and alkaline solutions; methyl orange is orange, pink or red, and yellow; phenolphthalein is colourless, colourless and pink.
A universal indicator gives a different colour at each part of the scale — green at pH 7, yellow through red as the solution becomes more acidic, and blue through indigo to violet as it becomes more basic — so it gives a value rather than only a verdict.
Flowchart
FC 13.3 — Identifying An Unknown Salt From Start To Finish Note colour, odour and physical state of the solid → heat a little of it in a dry test tube and identify any gas and the residue → make a solution in water, or in nitric acid if it will not dissolve → test one portion for the cation by adding sodium hydroxide drop by drop and then in excess, and a second portion with ammonium hydroxide → confirm a sodium, potassium, calcium or copper cation by a flame test on a clean platinum wire → test a fresh portion for the anion with dilute sulphuric acid, then with concentrated sulphuric acid if no gas is evolved, then with barium chloride after acidifying → combine the cation and the anion to name the salt
13.11 Key Terms
- Anhydrous — describing a salt that contains no water of crystallisation.
- Bleaching — the destruction of a colour, as when chlorine whitens moist litmus paper.
- Brown ring test — the confirmatory test for a nitrate, in which a brown ring appears where concentrated sulphuric acid meets a salt solution containing freshly prepared ferrous sulphate.
- Decrepitation — the crackling and crumbling of crystals when they are heated, seen with lead nitrate.
- Deliquescent — describing a solid that absorbs so much moisture from the air that it dissolves in it.
- Dry test — a test carried out on the solid substance itself, using colour, physical state, dry heating or the flame test.
- Effervescence — the rapid bubbling seen when a gas is set free in a liquid.
- Flame test — a dry test in which a salt held on a clean platinum wire colours a non-luminous flame according to the metal it contains.
- Hydrated salt — a salt whose crystals contain water of crystallisation.
- Hygroscopic — describing a solid that absorbs moisture from the air and turns damp without dissolving in it.
- Indicator — a substance that shows one colour in an acidic solution and another in an alkaline one.
- Nessler’s reagent — an alkaline solution of potassium tetraiodomercurate(II), which turns brown with ammonia or an ammonium salt.
- pH paper — paper soaked in universal indicator and dried, used to find the pH of a solution by matching its colour against a chart.
- pH scale — the hydrogen ion scale, running from 0 to 14, on which 7 is neutral, below 7 acidic and above 7 basic.
- Residue — the solid left behind at the end of a reaction, especially after heating or filtering.
- Universal indicator — a mixture of indicators that gives a different colour at each part of the pH scale.
- Wafting — fanning a little of a gas towards the face with the hand, the only safe way to smell a gas.
- Wet test — a test in which reagents are added to a solution of the substance and the substance is identified from what is observed.
13.12 Points To Ponder
Every test in this chapter identifies a substance by what it does rather than by what it is. Is there anything about a salt that could be observed directly, without making it react at all — and how far would such observations take you?
Several conclusions in this chapter rest on a negative result: dichromate paper staying orange, a precipitate refusing to dissolve in excess alkali. Why is ‘nothing happened’ usually weaker evidence than ‘something happened’, and what has to be true before a negative result can be trusted?
Sodium’s golden yellow is an obstacle in one test and a clear answer in another. What decides whether a strong, easily seen effect helps an experiment or spoils it?
Lime water cannot distinguish carbon dioxide from sulphur dioxide, and yet no one has replaced it. What makes a test worth keeping even when it cannot give a unique answer on its own?
You are handed an unknown white solid and allowed exactly three tests. Which three would you choose, in what order, and what would you knowingly be giving up by not choosing others?
Suppose a salt gave no flame colour, no precipitate with either alkali, and no gas with dilute or concentrated acid. Would that mean it could not be identified, or only that this particular set of tests is exhausted? What kind of evidence would you look for next?
13.13 Objective Questions
Multiple Choice
- A colourless gas turns lime water milky and also decolourises pink potassium permanganate solution. The gas is: [1]
- carbon dioxide
- sulphur dioxide
- hydrogen sulphide
- nitrogen dioxide
- A solid is heated strongly and the residue is yellow while hot and white after cooling. The residue is: [1]
- lead monoxide
- copper oxide
- zinc oxide
- mercury(II) oxide
- A salt solution gives a white precipitate with sodium hydroxide solution which dissolves in excess, but gives a white precipitate with ammonium hydroxide solution which does not dissolve in excess. The cation present is: [1]
- \(\mathrm{Zn^{2+}}\)
- \(\mathrm{Pb^{2+}}\)
- \(\mathrm{Cu^{2+}}\)
- \(\mathrm{Fe^{3+}}\)
- The reagent used along with concentrated sulphuric acid to confirm the nitrate ion is: [1]
- barium chloride solution
- lead acetate solution
- freshly prepared ferrous sulphate solution
- ammonium acetate solution
- A gas is passed over blue cobalt chloride paper and the paper turns pink. The gas is: [1]
- ammonia
- water vapour
- hydrogen chloride
- oxygen
- A flame test on a salt gives a persistent golden yellow colour which vanishes when viewed through blue glass. The cation present is: [1]
- \(\mathrm{K^{+}}\)
- \(\mathrm{Ca^{2+}}\)
- \(\mathrm{Na^{+}}\)
- \(\mathrm{Cu^{2+}}\)
Assertion–Reason
The following questions are assertion–reason based. Choose the answer based on 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): In the flame test the wire used is made of platinum. Reason (R): Platinum has a high melting point and imparts no colour of its own to the flame. [1]
- Assertion (A): Nitric acid is added before barium chloride solution in the test for a sulphate. Reason (R): Barium sulphate dissolves in dilute nitric acid. [1]
- Assertion (A): Manganese dioxide gives chlorine when warmed with concentrated hydrochloric acid, but copper oxide does not. Reason (R): Manganese dioxide is an oxidising agent while copper oxide is a basic oxide. [1]
- Assertion (A): An alkali is added to a salt solution one drop at a time. Reason (R): A precipitate that redissolves in excess alkali may otherwise be missed. [1]
Match The Following
- Match each reagent in Column I with the ion it confirms in Column II. [5]
| Column I | Column II |
|---|---|
| a. Lead acetate paper | i. Chloride ion |
| b. Nessler’s reagent | ii. Sulphate ion |
| c. Silver nitrate solution, precipitate soluble in ammonium hydroxide | iii. Sulphide ion |
| d. Barium chloride solution after acidifying with nitric acid | iv. Sulphite ion |
| e. Acidified potassium dichromate paper turning green | v. Ammonium ion |
Fill In The Blanks
A gas that turns moist red litmus paper blue is ……… in nature. [1]
On heating, sodium nitrate gives off ……… but no brown gas. [1]
Anhydrous copper sulphate is ……… in colour and turns ……… on adding water. [1]
A solution of pH ……… is neutral, and a universal indicator gives it a ……… colour. [1]
True Or False
Litmus paper must be moist before a gas is tested with it. [1]
Every nitrate is soluble in water. [1]
Phenolphthalein can be used to distinguish an acidic solution from a neutral one. [1]
Name The Following
The gas that rekindles a glowing splinter. [1]
The gas that turns lead acetate paper black. [1]
The alkali that dissolves the white precipitate of zinc hydroxide but not that of lead hydroxide. [1]
The metal ion that gives a peacock bluish green colour in a flame test. [1]
Give One Word Or Phrase For
The crackling of crystals of a salt when they are heated. [1]
A solid that absorbs moisture from the air and dissolves in the water it has absorbed. [1]
The liquid that passes through the filter paper during filtration. [1]
Odd One Out, And Give A Reason
Zinc carbonate, copper carbonate, calcium carbonate, sodium carbonate. [1]
Chlorine, ammonia, hydrogen chloride, nitrogen dioxide. [1]
13.14 Check Your Concepts
Define a dry test and a wet test, and give one example of each from this chapter. [2]
State why litmus paper is moistened before a gas is tested with it. [2]
Name the two gases in this chapter that are coloured, and give the colour of each. [2]
Give a reason why a hydrated nitrate gives three gases on strong heating while a carbonate gives only one. [3]
Explain why the milkiness produced when carbon dioxide is passed through lime water disappears when the gas is passed for a longer time. [3]
Give a reason why nitric acid is added to the salt solution before barium chloride solution in the test for a sulphate. [3]
Differentiate between a hygroscopic solid and a deliquescent solid, stating the basis of difference. [3]
Why is the flame test carried out in the non-luminous flame of the burner and not in the luminous flame? Give two reasons. [3]
What would happen if a student added sodium hydroxide solution quickly, in one large volume, to a solution containing the zinc ion? Explain what information would be lost. [4]
Explain, with the reason in each case, why sodium hydroxide solution alone cannot distinguish a lead salt from a zinc salt, and how the distinction is made. [4]
A blue glass is used while performing a flame test on a mixture of two salts. Explain what the blue glass does and why it is needed. [4]
Describe how you would show, using two tests for each, which of two unlabelled solutions is a dilute acid and which is an alkali. [5]
Sulphur dioxide and carbon dioxide behave alike towards lime water but differently towards acidified potassium permanganate. Explain the similarity and the difference, and state which property of sulphur dioxide is responsible. [5]
13.15 Equations, Conversions & Identification
Write the balanced equation for the action of dilute hydrochloric acid on magnesium. [2]
Write the balanced equation for the action of heat on lead nitrate. [2]
Write the balanced equation for the reaction between magnesium sulphate solution and barium chloride solution. [2]
Complete and balance the following equations: [3]
- \(\mathrm{CuCO_3} \xrightarrow{\ \Delta\ }\)
- \(\mathrm{MnO_2} + \mathrm{HCl} \rightarrow\)
- \(\mathrm{NH_4Cl} + \mathrm{NaOH} \rightarrow\)
Name the gas evolved and give one test for it in each case: [3]
- dilute sulphuric acid is added to zinc sulphide;
- ammonium sulphate is warmed with sodium hydroxide solution;
- concentrated sulphuric acid is warmed with sodium chloride.
State what you would see when: [4]
- copper nitrate hexahydrate is heated strongly in a dry test tube;
- ammonium hydroxide solution is added in excess to copper sulphate solution;
- moist starch iodide paper is held in a jar of chlorine;
- a glowing splinter is introduced into a jar of oxygen.
How would you obtain each of the following, starting from the substance named? [4]
- chlorine from manganese dioxide;
- carbon dioxide from zinc carbonate, without heating it;
- ammonia from ammonium chloride;
- water vapour from hydrated sodium carbonate.
Identify the substance in each case: [4]
- a bluish green crystalline solid which, on heating, gives steamy vapours, a reddish brown gas and a gas that rekindles a glowing splint, leaving a black residue;
- a colourless gas with a foul smell which turns lead nitrate solution black;
- a white salt whose solution gives a pale green precipitate with sodium hydroxide solution that slowly turns brown, and a white precipitate with barium chloride solution that is insoluble in mineral acid;
- a black powder which gives a greenish yellow gas when warmed with concentrated hydrochloric acid.
Fig. 13.4-2 shows the brown ring test. Name the two liquid layers marked in the figure, state where the brown ring appears, and give one reason why the tube must not be shaken. [3]
Fig. 13.7-1 shows the stages of a flame test. Name the acid used to make the paste, state the purpose of the second stage shown, and state the part of the flame in which the observation is made. [3]