Chapter 9: Practical Work
In 1860, in a laboratory at Heidelberg in Germany, two men spent weeks boiling down mineral water. Robert Bunsen, a chemist, and Gustav Kirchhoff, a physicist, were working with tens of thousands of litres brought from the springs at Dürkheim. They were not searching for anything in particular. They were simply looking at what the water had left behind.
Their method was disarmingly simple. A trace of the residue was carried into a burner flame on a clean wire, and the light of that flame was passed through a prism. Bunsen already knew that a sodium salt paints a flame golden yellow and a potassium salt a pale lilac. The prism went further: it spread the flame’s light into a row of separate coloured lines, and each metal produced its own set.
In the light from the Dürkheim residue there were two brilliant blue lines that belonged to no known metal. Bunsen named the new element caesium, from a Latin word for sky blue. The following year, working the same way on the mineral lepidolite, the pair found deep red lines and named a second new element rubidium.
Neither element was ever seen as a lump of metal during that work. Both were identified purely by what they did to a flame — by observation, carefully made and carefully read. That is what the whole of this chapter is about.
Diagram
Fig. 9.0-1 — Reading A Flame With A Prism Laboratory bench scene, left to right: a gas burner with a small, non-luminous flame coloured by a salt held on a clean wire loop; a narrow slit; a glass prism on a stand; and a screen or eyepiece on which the flame’s light appears as a row of separate bright vertical lines against a dark band. Label the burner, the wire loop carrying the salt, the slit, the prism and the line spectrum. Show two closely spaced bright lines on the screen picked out with a pointer.
9.1 Recognition And Identification Of Gases
In a laboratory you are often handed a jar of gas and asked what it is. You cannot read a label, because there is none. You have to work it out from what the gas looks like, what it smells like, and how it behaves towards a few simple substances.
Ten gases are met again and again at this level. Each one has a small set of properties that, taken together, belongs to it alone. Learn to make the observations carefully, and the identification follows almost by itself.
Four checks are worth making on almost any gas before anything else: its colour, its smell, its effect on moist litmus paper, and what a burning splinter does in it. These first four checks are shown in Fig. 9.1-1.
Diagram
Fig. 9.1-1 — Four First Checks On An Unknown Gas Four small panels in a row, each showing a gas jar with a cover slip partly slid aside. Panel 1: student looking at the jar from the side, labelled ‘colour’. Panel 2: student wafting the gas towards the nose with one hand, well away from the face, labelled ‘odour — waft, never sniff directly’. Panel 3: a strip of moist blue litmus paper and a strip of moist red litmus paper being lowered into the jar, labelled ‘action on moist litmus’. Panel 4: a burning wooden splinter being lowered into the jar, labelled ‘burning splinter’. Label the gas jar, cover slip and splinter in the first panel only.
A caution runs through all of this work. Several of these gases are poisonous or choking, so they are prepared and tested in a fume cupboard, and a gas is never sniffed directly — a little is wafted towards the nose with the hand.
9.1.1 Hydrogen
Hydrogen is given off when a reactive metal — one placed above hydrogen in the activity series, such as magnesium, zinc or iron — is added to a dilute acid. The metal takes the place of hydrogen in the acid, and the hydrogen escapes as a gas.
active metal + dilute acid → salt + hydrogen
\[\\mathrm{Mg}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(aq)} \\rightarrow \\mathrm{MgSO\_4}\\text{(aq)} + \\mathrm{H\_2}\\text{(g)}\]
\[\\mathrm{Zn}\\text{(s)} + 2,\\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{ZnCl\_2}\\text{(aq)} + \\mathrm{H\_2}\\text{(g)}\]
You recognise hydrogen by the following.
The gas is colourless and odourless.
It is neutral to litmus: neither moist blue nor moist red litmus paper changes colour.
Pure hydrogen burns with a pale blue flame when a burning splinter is brought to the mouth of the jar. The gas itself burns; it does not support the burning of the splinter, which goes out.
\[2,\\mathrm{H\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)} \\rightarrow 2,\\mathrm{H\_2O}\\text{(l)}\]
- Hydrogen mixed with air burns with a sharp pop sound when a burning paper is brought near it.
9.1.2 Oxygen
Some heavy metal oxides are unstable to heat. On strong heating they give up part of their oxygen, leaving behind a lower oxide or even the metal itself.
heavy metal oxide → metal oxide or metal + oxygen
\[2,\\mathrm{Pb\_3O\_4}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 6,\\mathrm{PbO}\\text{(s)} + \\mathrm{O\_2}\\text{(g)}\]
\[2,\\mathrm{PbO\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{PbO}\\text{(s)} + \\mathrm{O\_2}\\text{(g)}\]
\[2,\\mathrm{HgO}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{Hg}\\text{(l)} + \\mathrm{O\_2}\\text{(g)}\]
You recognise oxygen by the following.
The gas is colourless, odourless and neutral to litmus.
It rekindles a glowing wooden splinter. A splinter that has just been blown out, still glowing red at the tip, bursts back into flame in the jar. Oxygen does not burn itself; it supports burning.
The gas is absorbed by a colourless alkaline solution of pyrogallol, and the solution turns dark brown.
9.1.3 Carbon Dioxide
Carbon dioxide has two common sources. Most metallic carbonates give it off when heated, the exceptions being sodium carbonate and potassium carbonate, which are stable to heat. Any carbonate or hydrogen carbonate also gives it off at once when a dilute acid is added.
metallic carbonate → metal oxide + carbon dioxide
\[\\mathrm{ZnCO\_3}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{ZnO}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)}\]
\[\\mathrm{CuCO\_3}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CuO}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)}\]
metal carbonate or hydrogen carbonate + dilute acid → salt + water + carbon dioxide
\[\\mathrm{CaCO\_3}\\text{(s)} + 2,\\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{CaCl\_2}\\text{(aq)} + \\mathrm{H\_2O}\\text{(l)} + \\mathrm{CO\_2}\\text{(g)}\]
\[\\mathrm{NaHCO\_3}\\text{(s)} + \\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{NaCl}\\text{(aq)} + \\mathrm{H\_2O}\\text{(l)} + \\mathrm{CO\_2}\\text{(g)}\]
You recognise carbon dioxide by the following.
The gas is colourless and odourless.
It turns moist blue litmus a faint red. The colour is faint because carbon dioxide dissolves in the moisture to give a very weak acid.
Passed through lime water — a clear, saturated solution of calcium hydroxide — it turns the solution milky. The milkiness is a fine white precipitate, the solid that appears when two solutions react to give an insoluble product. Here it is calcium carbonate. The way the gas is passed in is shown in Fig. 9.1-2.
\[\\mathrm{Ca(OH)\_2}\\text{(aq)} + \\mathrm{CO\_2}\\text{(g)} \\rightarrow \\mathrm{CaCO\_3}\\text{(s)} + \\mathrm{H\_2O}\\text{(l)}\]
If you go on passing the gas, the milkiness disappears again. The insoluble carbonate is converted into calcium hydrogen carbonate, which is soluble, so the liquid becomes clear once more.
\[\\mathrm{CaCO\_3}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)} + \\mathrm{H\_2O}\\text{(l)} \\rightarrow \\mathrm{Ca(HCO\_3)\_2}\\text{(aq)}\]
- The gas has no effect on filter paper dipped in acidified potassium dichromate or in acidified potassium permanganate. This last observation looks dull, but it is the one that separates carbon dioxide from sulphur dioxide.
Diagram
Fig. 9.1-2 — Passing A Gas Through Lime Water Test tube held in a stand, one-third filled with clear lime water; a delivery tube from a second tube dips below the surface of the liquid; bubbles rising through the lime water; the liquid drawn cloudy white in the upper part. Label the delivery tube, lime water, and the white precipitate of calcium carbonate settling at the bottom.
9.1.4 Chlorine
Chlorine is obtained by warming an oxidising agent, such as red lead, lead dioxide or manganese dioxide, with concentrated hydrochloric acid. The oxidising agent takes hydrogen from the acid and sets the chlorine free.
oxidising agent + concentrated hydrochloric acid → metal chloride + water + chlorine
\[\\mathrm{Pb\_3O\_4}\\text{(s)} + 8,\\mathrm{HCl}\\text{(aq)} \\xrightarrow{\\ \\Delta\\ } 3,\\mathrm{PbCl\_2}\\text{(s)} + 4,\\mathrm{H\_2O}\\text{(l)} + \\mathrm{Cl\_2}\\text{(g)}\]
\[\\mathrm{MnO\_2}\\text{(s)} + 4,\\mathrm{HCl}\\text{(aq)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{MnCl\_2}\\text{(aq)} + 2,\\mathrm{H\_2O}\\text{(l)} + \\mathrm{Cl\_2}\\text{(g)}\]
Chlorine is poisonous and is always handled in a fume cupboard. You recognise it by the following.
The gas is greenish yellow, with a sharp, pungent, choking odour.
It turns moist blue litmus paper red and then bleaches it, that is, decolourises it. Bleaching is the destruction of a colour, and here it is permanent. Chlorine first reacts with the moisture to give hydrochloric acid and hypochlorous acid; the second of these breaks up and releases oxygen, which destroys the colour.
\[\\mathrm{Cl\_2}\\text{(g)} + \\mathrm{H\_2O}\\text{(l)} \\rightarrow \\mathrm{HCl}\\text{(aq)} + \\mathrm{HClO}\\text{(aq)}\]
\[\\mathrm{HClO}\\text{(aq)} \\rightarrow \\mathrm{HCl}\\text{(aq)} + \[\\mathrm{O}]\]
colouring matter + [O] → colourless or bleached product
- It turns moist starch iodide paper — filter paper soaked in potassium iodide and starch solution — blue black. Chlorine displaces iodine from the potassium iodide, and the iodine gives the familiar blue black colour with starch.
\[\\mathrm{Cl\_2}\\text{(g)} + 2,\\mathrm{KI}\\text{(aq)} \\rightarrow 2,\\mathrm{KCl}\\text{(aq)} + \\mathrm{I\_2}\\text{(aq)}\]
starch + iodine → blue black colour
- Passed through silver nitrate solution, it gives a white precipitate of silver chloride. The chlorine first forms hydrochloric acid with the water, and it is the chloride from that acid which is precipitated.
9.1.5 Hydrogen Chloride
Hydrogen chloride gas is obtained by warming a metal chloride, such as sodium chloride or potassium chloride, with concentrated sulphuric acid. The sulphuric acid, being far less volatile, holds on to the metal and drives the hydrogen chloride out as a gas.
metal chloride + concentrated sulphuric acid → metal hydrogensulphate + hydrogen chloride
\[\\mathrm{NaCl}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(l)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{NaHSO\_4}\\text{(s)} + \\mathrm{HCl}\\text{(g)}\]
\[\\mathrm{KCl}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(l)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{KHSO\_4}\\text{(s)} + \\mathrm{HCl}\\text{(g)}\]
You recognise hydrogen chloride by the following.
The gas is colourless and has a pungent, choking odour.
It turns moist blue litmus paper red, and does so strongly.
If a glass rod dipped in ammonia solution is held near the gas, dense white fumes of solid ammonium chloride form in the air between them. This is shown in Fig. 9.1-3.
\[\\mathrm{NH\_3}\\text{(aq)} + \\mathrm{HCl}\\text{(g)} \\rightarrow \\mathrm{NH\_4Cl}\\text{(s)}\]
- Passed through silver nitrate solution, it gives a white precipitate of silver chloride, and the precipitate dissolves when excess ammonium hydroxide is added.
\[\\mathrm{AgNO\_3}\\text{(aq)} + \\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{AgCl}\\text{(s)} + \\mathrm{HNO\_3}\\text{(aq)}\]
Diagram
Fig. 9.1-3 — Dense White Fumes Of Ammonium Chloride A gas jar of hydrogen chloride with the cover slip slid aside; a glass rod held just above the mouth, its tip wet with ammonia solution; a thick white cloud drawn in the space between the rod tip and the jar mouth. Label the glass rod, the drop of ammonia solution on its tip, and the white fumes of ammonium chloride.
9.1.6 Sulphur Dioxide
Sulphur dioxide is given off when a dilute acid is added to a metallic sulphite or a hydrogen sulphite. The acid displaces the weaker sulphurous acid, which immediately breaks up into water and sulphur dioxide.
metal sulphite or hydrogen sulphite + dilute acid → salt + water + sulphur dioxide
\[\\mathrm{Na\_2SO\_3}\\text{(s)} + 2,\\mathrm{HCl}\\text{(aq)} \\rightarrow 2,\\mathrm{NaCl}\\text{(aq)} + \\mathrm{H\_2O}\\text{(l)} + \\mathrm{SO\_2}\\text{(g)}\]
\[\\mathrm{Na\_2SO\_3}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(aq)} \\rightarrow \\mathrm{Na\_2SO\_4}\\text{(aq)} + \\mathrm{H\_2O}\\text{(l)} + \\mathrm{SO\_2}\\text{(g)}\]
\[2,\\mathrm{NaHSO\_3}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(aq)} \\rightarrow \\mathrm{Na\_2SO\_4}\\text{(aq)} + 2,\\mathrm{H\_2O}\\text{(l)} + 2,\\mathrm{SO\_2}\\text{(g)}\]
You recognise sulphur dioxide by the following.
The gas is colourless, with the suffocating smell of burning sulphur.
A burning wooden splinter is extinguished in it. Sulphur dioxide neither burns nor supports burning.
It turns moist blue litmus red and then bleaches it. Here the bleaching is temporary, and the colour slowly returns on standing in air.
It turns lime water milky, just as carbon dioxide does, because insoluble calcium sulphite is formed.
\[\\mathrm{Ca(OH)\_2}\\text{(aq)} + \\mathrm{SO\_2}\\text{(g)} \\rightarrow \\mathrm{CaSO\_3}\\text{(s)} + \\mathrm{H\_2O}\\text{(l)}\]
On passing excess gas the milkiness disappears, exactly as before, because the soluble hydrogen sulphite is formed.
\[\\mathrm{CaSO\_3}\\text{(s)} + \\mathrm{H\_2O}\\text{(l)} + \\mathrm{SO\_2}\\text{(g)} \\rightarrow \\mathrm{Ca(HSO\_3)\_2}\\text{(aq)}\]
- It decolourises purplish pink acidified potassium permanganate solution. This distinguishes it from carbon dioxide, which leaves the solution unchanged.
\[2,\\mathrm{KMnO\_4} + 2,\\mathrm{H\_2O} + 5,\\mathrm{SO\_2} \\rightarrow \\mathrm{K\_2SO\_4} + 2,\\mathrm{MnSO\_4} + 2,\\mathrm{H\_2SO\_4}\]
- It changes the orange solution of acidified potassium dichromate to green. This also distinguishes it from carbon dioxide.
\[\\mathrm{K\_2Cr\_2O\_7} + \\mathrm{H\_2SO\_4} + 3,\\mathrm{SO\_2} \\rightarrow \\mathrm{Cr\_2(SO\_4)\_3} + \\mathrm{K\_2SO\_4} + \\mathrm{H\_2O}\]
- It has no effect on lead acetate paper. This distinguishes it from hydrogen sulphide.
Notice how the last three observations work. Sulphur dioxide shares its two most obvious properties — colourless, turns lime water milky — with carbon dioxide, so those two observations alone can never decide between them. The dichromate and permanganate tests are brought in precisely because the two gases behave differently towards them.
9.1.7 Hydrogen Sulphide
Hydrogen sulphide is given off when a dilute acid is added to a metallic sulphide such as iron(II) sulphide or zinc sulphide.
metal sulphide + dilute acid → salt + hydrogen sulphide
\[\\mathrm{FeS}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(aq)} \\rightarrow \\mathrm{FeSO\_4}\\text{(aq)} + \\mathrm{H\_2S}\\text{(g)}\]
\[\\mathrm{ZnS}\\text{(s)} + 2,\\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{ZnCl\_2}\\text{(aq)} + \\mathrm{H\_2S}\\text{(g)}\]
You recognise hydrogen sulphide by the following.
The gas is colourless, with a foul smell like that of rotten eggs.
It turns moist blue litmus red.
It turns lead acetate paper black, because insoluble black lead sulphide is formed on the paper.
\[\\mathrm{(CH\_3COO)\_2Pb}\\text{(aq)} + \\mathrm{H\_2S}\\text{(g)} \\rightarrow \\mathrm{PbS}\\text{(s)} + 2,\\mathrm{CH\_3COOH}\\text{(aq)}\]
- It turns colourless lead nitrate solution black, for the same reason.
\[\\mathrm{Pb(NO\_3)\_2}\\text{(aq)} + \\mathrm{H\_2S}\\text{(g)} \\rightarrow \\mathrm{PbS}\\text{(s)} + 2,\\mathrm{HNO\_3}\\text{(aq)}\]
Hydrogen sulphide also turns acidified potassium permanganate from pink to colourless and acidified potassium dichromate from orange to green, as sulphur dioxide does. There is one difference in what you see: yellow sulphur is set free at the same time, so the solution is left cloudy rather than clear.
9.1.8 Ammonia
Ammonia is obtained by heating an ammonium salt, such as ammonium chloride or ammonium sulphate, with an alkali.
ammonium salt + alkali → salt + water + ammonia
\[2,\\mathrm{NH\_4Cl}\\text{(s)} + \\mathrm{Ca(OH)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CaCl\_2}\\text{(s)} + 2,\\mathrm{H\_2O}\\text{(g)} + 2,\\mathrm{NH\_3}\\text{(g)}\]
\[\\mathrm{NH\_4Cl}\\text{(s)} + \\mathrm{NaOH}\\text{(aq)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{NaCl}\\text{(aq)} + \\mathrm{H\_2O}\\text{(l)} + \\mathrm{NH\_3}\\text{(g)}\]
You recognise ammonia by the following.
The gas is colourless, with a characteristic sharp, pungent smell.
It turns moist red litmus blue. Of the ten gases in this chapter, ammonia is the only one that does this — it is the only basic gas among them.
If a glass rod dipped in concentrated hydrochloric acid is brought near the gas, dense white fumes of ammonium chloride are formed. This is the same reaction met in 9.1.5, approached from the other side.
It turns colourless Nessler’s reagent, which is potassium mercuric iodide, \(\\mathrm{K\_2HgI\_4}\), brown.
Passed through copper sulphate solution, it first gives a pale blue precipitate. On passing excess ammonia, this precipitate dissolves to form a deep blue solution.
9.1.9 Water Vapour
Many crystals hold water inside the crystal itself. Heating such a crystalline substance — hydrated sodium carbonate or hydrated copper sulphate, for example — drives that water out as vapour and leaves the dry solid behind.
hydrated salt → anhydrous salt + water vapour
\[\\mathrm{Na\_2CO\_3!\\cdot!10H\_2O}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Na\_2CO\_3}\\text{(s)} + 10,\\mathrm{H\_2O}\\text{(g)}\]
\[\\mathrm{CuSO\_4!\\cdot!5H\_2O}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CuSO\_4}\\text{(s)} + 5,\\mathrm{H\_2O}\\text{(g)}\]
The blue crystals of hydrated copper sulphate leave a white solid behind. You recognise the vapour by the following.
The vapours are colourless and odourless, and condense to a clear liquid on the cooler part of the test tube.
The liquid is neutral to litmus.
The liquid turns white anhydrous copper sulphate blue.
\[\\mathrm{CuSO\_4}\\text{(s)} + 5,\\mathrm{H\_2O}\\text{(l)} \\rightarrow \\mathrm{CuSO\_4!\\cdot!5H\_2O}\\text{(s)}\]
- It turns blue cobalt chloride paper pink.
\[\\mathrm{CoCl\_2}\\text{(s)} + 2,\\mathrm{H\_2O}\\text{(l)} \\rightarrow \\mathrm{CoCl\_2!\\cdot!2H\_2O}\\text{(s)}\]
These last two tests show only that the liquid is water. They do not show that it is pure water, since a solution of something in water would give the same result.
9.1.10 Nitrogen Dioxide
Heating a heavy metal nitrate, such as copper nitrate or lead nitrate, breaks it down into the metal oxide, oxygen and the brown gas nitrogen dioxide.
metal nitrate → metal oxide + oxygen + nitrogen dioxide
\[2,\\mathrm{Cu(NO\_3)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{CuO}\\text{(s)} + \\mathrm{O\_2}\\text{(g)} + 4,\\mathrm{NO\_2}\\text{(g)}\]
\[2,\\mathrm{Pb(NO\_3)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{PbO}\\text{(s)} + \\mathrm{O\_2}\\text{(g)} + 4,\\mathrm{NO\_2}\\text{(g)}\]
Sodium nitrate and potassium nitrate behave differently. They do not produce nitrogen dioxide on heating.
You recognise nitrogen dioxide by the following.
The gas is brown, has an irritating, pungent odour, and is non-combustible.
It turns moist blue litmus paper red.
It liberates iodine, seen as violet vapours, from potassium iodide solution, and turns potassium iodide paper brown.
\[2,\\mathrm{KI}\\text{(aq)} + 2,\\mathrm{NO\_2}\\text{(g)} \\rightarrow 2,\\mathrm{KNO\_2}\\text{(aq)} + \\mathrm{I\_2}\\text{(s)}\]
- It turns green acidified ferrous sulphate solution brown.
Of the ten gases, only nitrogen dioxide has a colour you can see at a glance apart from chlorine. Brown points to nitrogen dioxide; greenish yellow points to chlorine.
9.1.11 The Ten Gases At A Glance
The colour, the odour and the acidic, basic or neutral nature of all ten gases are collected in Table 9.1. Read down the columns rather than across the rows: the colour column alone narrows the field to two gases, and the odour column narrows it much further.
Table 9.1 — Colour, Odour And Nature Of The Ten 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 |
Three of the ten are neutral to litmus, six are acidic, and exactly one — ammonia — is basic. That single fact makes moist red litmus paper one of the most useful strips in the laboratory.
9.2 Action Of Heat On A Given Substance
Heat alone tells you a surprising amount about a solid. Some solids give off the water locked inside their crystals. Some decompose and release a gas. Some change colour, and some pass straight into vapour without melting at all.
None of this needs a single reagent. You take a small quantity of the unknown solid in a clean, dry test tube, heat it gently at first and then more strongly, and watch. The correct way to hold and heat the tube is shown in Fig. 9.2-1.
Three things are worth watching at once: any change in the colour of the solid, anything given off as a gas or vapour, and the colour of whatever residue is left when the tube has cooled. Put together, these observations often name the substance.
Diagram
Fig. 9.2-1 — Heating A Solid In A Dry Test Tube Test tube held in a wooden holder, sloped at about 45 degrees with the mouth pointing away from the student and away from anyone nearby; a small quantity of solid at the bottom; a burner flame under the tube; the cooler upper part of the tube shown clear. Label the test-tube holder, the solid, the flame, the cooler upper part of the tube, and mark the direction of the mouth with the note ‘mouth pointed away from everyone’.
9.2.1 Carbonates On Heating
Most metallic carbonates break down on strong heating into the metal oxide and carbon dioxide. You met this decomposition in 9.1.3 as a way of obtaining the gas. Here you use it in the opposite direction — the gas is the clue, and the solid is what you are trying to name.
Take copper carbonate first. It is a light green amorphous powder, which means it is a powder with no regular crystal shape. On strong heating it turns black.
\[\\mathrm{CuCO\_3}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CuO}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)}\]
The gas that comes off is colourless and odourless, it extinguishes a burning wooden splinter, and it turns lime water milky. It has no effect on acidified potassium dichromate or acidified potassium permanganate, which rules out sulphur dioxide. The gas is therefore carbon dioxide, the black residue is copper oxide, and the original light green powder was copper carbonate.
Zinc carbonate tells a similar story with one extra observation. It is a light amorphous white solid, and on strong heating it changes to pale yellow.
\[\\mathrm{ZnCO\_3}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{ZnO}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)}\]
Again the gas turns lime water milky, and the milkiness disappears when excess gas is passed. The important detail is what happens as the tube cools: the yellow residue turns white. Zinc oxide is yellow when hot and white when cold, and that reversible colour change is a signature you should watch for.
9.2.2 Hydrated Salts On Heating
Many crystals contain water built into the crystal structure itself. This is called the water of crystallisation, and it is the reason a dry-looking crystal can give off steam when heated. A salt from which this water has been driven out is described as anhydrous.
Washing soda is hydrated sodium carbonate. Its white crystalline solid behaves dramatically on strong heating: it swells, crumbles and then melts, giving off steamy vapours.
\[\\mathrm{Na\_2CO\_3!\\cdot!10H\_2O}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Na\_2CO\_3}\\text{(s)} + 10,\\mathrm{H\_2O}\\text{(g)}\]
The vapours condense on the cooler parts of the test tube as tiny droplets of a colourless liquid. That liquid turns white anhydrous copper sulphate blue and turns blue cobalt chloride paper pink, so it is water. On cooling, a white amorphous residue of sodium carbonate is left.
Copper sulphate crystals go further, because on stronger heating the anhydrous salt itself decomposes. The blue crystalline solid first crumbles to a white amorphous powder, giving off steamy vapours that condense to water on the cooler part of the tube.
\[\\mathrm{CuSO\_4!\\cdot!5H\_2O}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CuSO\_4}\\text{(s)} + 5,\\mathrm{H\_2O}\\text{(g)}\]
Heat it harder and a black solid is left, and a fresh gas comes off.
\[2,\\mathrm{CuSO\_4}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{CuO}\\text{(s)} + 2,\\mathrm{SO\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)}\]
This second gas turns moist blue litmus red and changes acidified potassium dichromate solution from orange to green, so it is sulphur dioxide. The black residue is copper oxide. The blue crystalline solid you started with was copper sulphate crystals, long known as blue vitriol.
Notice that the same substance gave two different results at two different temperatures. Gentle heating removed only the water; strong heating destroyed the salt itself. When you report the action of heat, you must say how strongly the substance was heated.
9.2.3 Nitrates On Heating
Heavy metal nitrates decompose on heating to give the metal oxide, the brown gas nitrogen dioxide, and oxygen. You met that decomposition in 9.1.10. Where the nitrate is also hydrated, water comes off first, and the two stages are seen one after the other.
Zinc nitrate hexahydrate is a white crystalline deliquescent solid — deliquescent meaning that it absorbs so much moisture from the air that it eventually dissolves in the water it has taken up. On heating it melts to a white sticky mass and gives off steamy vapours.
\[\\mathrm{Zn(NO\_3)\_2!\\cdot!6H\_2O}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Zn(NO\_3)\_2}\\text{(s)} + 6,\\mathrm{H\_2O}\\text{(g)}\]
On heating strongly, the white residue gives off reddish brown fumes that turn moist blue litmus paper red.
\[2,\\mathrm{Zn(NO\_3)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{ZnO}\\text{(s)} + 4,\\mathrm{NO\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)}\]
A glowing wooden splint held in the reddish brown gas bursts into flame. That is not the nitrogen dioxide doing it — nitrogen dioxide does not support burning. It is the oxygen mixed with it. The residue is pale yellow when hot and white on cooling, so it is zinc oxide, and the original solid was zinc nitrate hexahydrate.
Copper nitrate hexahydrate is a bluish green crystalline solid. It behaves in exactly the same two stages, melting to a bluish green mass and giving off steamy vapours that condense as water.
\[\\mathrm{Cu(NO\_3)\_2!\\cdot!6H\_2O}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Cu(NO\_3)\_2}\\text{(s)} + 6,\\mathrm{H\_2O}\\text{(g)}\]
\[2,\\mathrm{Cu(NO\_3)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{CuO}\\text{(s)} + 4,\\mathrm{NO\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)}\]
Here the bluish green mass changes to a black residue of copper(II) oxide, a reddish brown gas is given off, and a gas that rekindles a glowing splinter is given off with it.
Lead nitrate is a heavy white crystalline solid, and it is not hydrated, so there are no steamy vapours at all. On strong heating it crumbles with a distinct crackling noise.
\[2,\\mathrm{Pb(NO\_3)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{PbO}\\text{(s)} + 4,\\mathrm{NO\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)}\]
The residue is the clue that names it. Lead oxide is reddish brown when hot; on cooling it changes to yellow, partly fuses into the glass, and stains it yellow. The colour changes of the two oxides are compared in Fig. 9.2-2.
Diagram
Fig. 9.2-2 — Residue Colours, Hot And Cold Two pairs of test tubes side by side. Left pair labelled ‘zinc oxide’: the hot tube shows a pale yellow residue with a flame beneath, the cold tube shows a white residue. Right pair labelled ‘lead oxide’: the hot tube shows a reddish brown residue with a flame beneath, the cold tube shows a yellow residue with a note that it stains the glass. Label each tube ‘hot’ or ‘cold’.
9.2.4 Solids That Sublime Or Decompose Without A Residue
Some solids leave nothing behind at all. On heating they pass into vapour, and on reaching the cool upper part of the test tube they form a solid again. The direct change from solid to vapour and back to solid, without a liquid stage, is sublimation.
Ammonium chloride is a white crystalline solid. On strong heating it does not simply vaporise — it splits into two gases, a basic one and an acidic one, which recombine on the cooler part of the tube.
\[\\mathrm{NH\_4Cl}\\text{(s)} \\xrightleftharpoons\[\\text{cool}]{\\text{heat}} \\mathrm{NH\_3}\\text{(g)} + \\mathrm{HCl}\\text{(g)}\]
Dense white fumes are seen, and they settle as a white mass on the cooler parts of the test tube. No residue is left behind. This is shown in Fig. 9.2-3.
Iodine is a violet crystalline solid, and its behaviour is true sublimation — the substance itself is unchanged.
\[\\mathrm{I\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{I\_2}\\text{(g)}\]
Violet vapours rise and settle on the cooler part of the tube as violet crystals, and nothing is left at the bottom. A filter paper dipped in starch solution turns blue when brought into contact with the vapour, and a paper dipped in silver nitrate solution turns yellow.
Ammonium dichromate behaves quite differently, and far more violently. It is an orange red crystalline solid, and on strong heating it swells up and decomposes with flashes of light.
\[\\mathrm{(NH\_4)\_2Cr\_2O\_7}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Cr\_2O\_3}\\text{(s)} + \\mathrm{N\_2}\\text{(g)} + 4,\\mathrm{H\_2O}\\text{(g)}\]
Steamy fumes condense as tiny droplets of water on the cooler parts of the tube, and the colourless liquid turns cobalt chloride paper pink. A greenish grey mass of chromium(III) oxide is left behind. A colourless, odourless gas is also given out which neither burns nor supports burning, and which does not turn lime water milky — it is nitrogen. Because this decomposition is violent and throws out hot solid, it is a demonstration for a teacher behind a screen, never a student activity.
Diagram
Fig. 9.2-3 — Sublimation In A Test Tube Test tube sloped in a holder over a flame, a white crystalline solid at the bottom, arrows showing vapour rising, and a white solid ring deposited on the cool upper wall of the tube. Label the solid being heated, the rising vapour, the cool upper part of the tube and the deposit re-formed there; add a note that no residue remains at the bottom.
9.2.5 Products Of Heating At A Glance
The ten substances of this section are collected in Table 9.2, with the residue each leaves and the gas each gives off.
Table 9.2 — Products Formed On Heating
| S. No. | Substance and its colour | Residue left | Colour of residue | Gas produced |
|---|---|---|---|---|
| 1. | Copper carbonate (\(\\mathrm{CuCO\_3}\)), green solid | Copper(II) oxide, \(\\mathrm{CuO}\) | Black solid | Carbon dioxide (\(\\mathrm{CO\_2}\)) |
| 2. | Zinc carbonate (\(\\mathrm{ZnCO\_3}\)), white solid | Zinc oxide, \(\\mathrm{ZnO}\) | Yellow when hot, white when cold | Carbon dioxide (\(\\mathrm{CO\_2}\)) |
| 3. | Washing soda (\(\\mathrm{Na\_2CO\_3!\\cdot!10H\_2O}\)), white crystals | Anhydrous \(\\mathrm{Na\_2CO\_3}\) | White powder | Water vapour (\(\\mathrm{H\_2O}\)) |
| 4. | Copper sulphate (\(\\mathrm{CuSO\_4!\\cdot!5H\_2O}\)), blue crystals | Anhydrous \(\\mathrm{CuSO\_4}\) | Dirty white | Water vapour (\(\\mathrm{H\_2O}\)) |
| 5. | Zinc nitrate (\(\\mathrm{Zn(NO\_3)\_2}\)), white solid | Zinc oxide, \(\\mathrm{ZnO}\) | Yellow when hot, white when cold | Nitrogen dioxide and oxygen (\(\\mathrm{NO\_2 + O\_2}\)) |
| 6. | Copper nitrate (\(\\mathrm{Cu(NO\_3)\_2}\)), green solid | Copper(II) oxide, \(\\mathrm{CuO}\) | Black solid | Nitrogen dioxide and oxygen (\(\\mathrm{NO\_2 + O\_2}\)) |
| 7. | Lead nitrate (\(\\mathrm{Pb(NO\_3)\_2}\)), white solid | Lead oxide, \(\\mathrm{PbO}\) | Yellow solid | Nitrogen dioxide and oxygen (\(\\mathrm{NO\_2 + O\_2}\)) |
| 8. | Ammonium chloride (\(\\mathrm{NH\_4Cl}\)), white solid | No residue | White sublimate | Hydrogen chloride and ammonia (\(\\mathrm{HCl + NH\_3}\)) |
| 9. | Iodine (\(\\mathrm{I\_2}\)), violet solid | No residue | Violet sublimate | Iodine vapours |
| 10. | Ammonium dichromate (\(\\mathrm{(NH\_4)\_2Cr\_2O\_7}\)), orange solid | Chromium(III) oxide, \(\\mathrm{Cr\_2O\_3}\) | Green solid | Water vapour and nitrogen (\(\\mathrm{H\_2O + N\_2}\)) |
Read as a working method rather than a list, these results give a short scheme for an unknown solid, set out in FC 9.1.
Flowchart
FC 9.1 — Identifying A Solid By The Action Of Heat Heat a little of the solid in a dry test tube → note any colour change while hot → test whatever is given off → branch A: colourless odourless gas turns lime water milky → a carbonate; branch B: steamy vapour condensing to a colourless liquid that turns cobalt chloride paper pink → a hydrated salt; branch C: reddish brown gas together with a gas that relights a glowing splint → a heavy metal nitrate; branch D: solid vanishes and re-forms on the cool wall, no residue → a substance that sublimes → in every branch, finish by noting the colour of the residue while hot and again when cold
9.3 Recognising A Substance From Colour, Odour, State And Action Of Heat
You have now seen a good many individual results. It is time to gather them into working rules, so that an unknown substance can be narrowed down before any reagent is opened.
Four kinds of clue are available before you add anything at all: the colour of the substance, its odour, its physical state, and its behaviour on heating. None of these on its own names a substance. Taken together, they usually leave only one or two possibilities.
9.3.1 Clues From Colour
The colour of a salt usually comes from the metal ion in it. A few metal ions are strongly coloured, and once you know them, a glance at the bottle narrows the field at once. These colours are collected in Table 9.3, and some of them are shown in Fig. 9.3-1.
Table 9.3 — Colour Of A Substance And The Ion Suspected
| Colour of the substance | Ion suspected |
|---|---|
| Blue or bluish green | \(\\mathrm{Cu^{2+}}\) |
| Light green | \(\\mathrm{Fe^{2+}}\) |
| Yellow or yellowish brown | \(\\mathrm{Fe^{3+}}\) |
| White or colourless | \(\\mathrm{Pb^{2+}}\), \(\\mathrm{Zn^{2+}}\), \(\\mathrm{Ca^{2+}}\), \(\\mathrm{Na^{+}}\), \(\\mathrm{K^{+}}\) or \(\\mathrm{NH\_4^{+}}\) |
The last row deserves attention. A white or colourless substance has told you almost nothing, because six common ions are all colourless. Colour is a strong clue when it is present and a weak one when it is absent.
Diagram
Fig. 9.3-1 — Colours Of Some Common Salts Four watch glasses in a row on a white bench, each holding a small heap of solid, with a hand lens resting beside them. From left: a bright blue solid labelled ‘copper ion, blue or bluish green’; a light green solid labelled ‘iron(II) ion, light green’; a yellowish brown solid labelled ‘iron(III) ion, yellow to yellowish brown’; and a white solid labelled ‘white or colourless — several ions possible’.
9.3.2 Clues From Odour
An odour, where there is one, points to the negative part of the salt rather than to the metal. Three smells are worth learning, and each is wafted towards the nose, never inhaled directly.
The smell of ammonia suggests an ammonium salt, containing \(\\mathrm{NH\_4^{+}}\).
The smell of hydrogen sulphide — rotten eggs — suggests a sulphide, containing \(\\mathrm{S^{2-}}\).
The smell of sulphur dioxide — burning sulphur — suggests a sulphite, containing \(\\mathrm{SO\_3^{2-}}\).
In each case the smell appears because a little of the salt has already reacted with moisture or acid in the air and released a trace of its gas.
9.3.3 Clues From Physical State
How the solid sits in the bottle is itself an observation. Two states are worth noting.
An amorphous salt — a powder with no regular crystal shape — suggests a carbonate, containing \(\\mathrm{CO\_3^{2-}}\).
A salt that is hygroscopic, meaning that it absorbs moisture from the air and turns damp without dissolving, or deliquescent, meaning that it absorbs enough to dissolve in the water taken up, suggests a chloride, \(\\mathrm{Cl^{-}}\), or a nitrate, \(\\mathrm{NO\_3^{-}}\).
9.3.4 Rules For The Action Of Heat
The individual results of 9.2 can now be stated as general rules. Each rule says which family of compounds behaves in a given way, so that a new substance from that family can be predicted rather than looked up.
An ammonium salt heated with an alkali, other than ammonium hydroxide, produces ammonia gas. Ammonium chloride leaves no residue on heating, and neither does ammonium nitrate, which decomposes violently and is never heated in a school laboratory. Ammonium dichromate leaves a greenish grey mass, as you saw in 9.2.4.
Lead dioxide, red lead, mercury(II) oxide, potassium nitrate and sodium nitrate all produce oxygen on heating. Notice that the two nitrates in this list give oxygen but no nitrogen dioxide, which is exactly what was said of them in 9.1.10.
Carbonates and hydrogen carbonates evolve carbon dioxide on heating, except potassium carbonate and sodium carbonate, which are stable.
Hydrated salts produce water vapour on heating.
Sulphites, and some sulphates, produce sulphur dioxide gas on heating.
Lead compounds decompose to give lead monoxide, \(\\mathrm{PbO}\), also called litharge. It is brown when hot, yellow when cold, and it sticks to the glass.
Zinc compounds decompose on heating to give zinc oxide, which is yellow when hot and white when cold.
\[\\mathrm{ZnO}\\text{(s, white)} \\xrightleftharpoons\[\\text{cool}]{\\text{heat}} \\mathrm{ZnO}\\text{(s, yellow)}\]
- Copper compounds decompose to give black copper oxide, \(\\mathrm{CuO}\).
These results are collected in Table 9.4.
Table 9.4 — Identification Of Solids By The Action Of Heat
| Compound | Result of the test |
|---|---|
| Ammonium salt | Ammonia gas is evolved |
| \(\\mathrm{PbO\_2}\), \(\\mathrm{Pb\_3O\_4}\), \(\\mathrm{HgO}\), \(\\mathrm{KNO\_3}\), \(\\mathrm{NaNO\_3}\) | Oxygen gas is evolved |
| Carbonate and hydrogen carbonate (except \(\\mathrm{K\_2CO\_3}\) and \(\\mathrm{Na\_2CO\_3}\)) | Carbon dioxide is evolved |
| Hydrated salt | Water vapour evolved |
| Sulphite or sulphate | Sulphur dioxide gas is evolved |
| Lead compounds, which give lead monoxide, \(\\mathrm{PbO}\) | Brown when hot, yellow when cold |
| Zinc compounds, which give zinc oxide | Yellow when hot, white when cold |
| Copper compounds, which give black copper oxide, \(\\mathrm{CuO}\) | Black residue |
9.3.5 Dry Tests And Wet Tests
Everything in this section so far has been done without adding a reagent. Such tests have a name.
A wet test, by contrast, is one in which reagents are added, usually to a solution of the substance, and a precipitate or a colour change is looked for. The two kinds of test are compared in Table 9.5.
The sensible order is dry tests first, wet tests afterwards. A dry test costs a pinch of substance and a minute of heating, and it can rule out whole families at once.
9.3.6 Solubility Of Salts And Bases In Water
A wet test depends on something being insoluble. When two solutions are mixed and a precipitate appears, the precipitate is the one product that will not stay dissolved. To predict such a reaction, you need to know which salts are soluble and which are not.
Table 9.6 sets this out. Read it as a grid: choose the row for the negative ion and the column for the metal, and the cell tells you whether that salt is soluble. The Roman numeral under each metal is the valency of that metal in the salt.
Table 9.6 — Solubility Of Salts And Bases In Water
| 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 | Sp | S |
| \(\\mathrm{OH}\) (I) | S | S | S | Sp | Sp | N | N | N | N | N | N | N | — | — |
| \(\\mathrm{S}\) (II) | S | S | S | Sp | S | — | 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 |
S : soluble N : insoluble Sp : sparingly soluble — : does not exist
Two patterns in the table are worth carrying in your head. Every nitrate is soluble, and every sodium and potassium salt is soluble — so a precipitate can never be a nitrate, nor a salt of sodium or potassium. Carbonates, sulphides, sulphites and phosphates, on the other hand, are insoluble for almost every metal except sodium, potassium and ammonium.
9.4 The Flame Test
Compounds of certain metals impart a characteristic colour to a flame. The colour comes from the metal, not from the rest of the salt, so any two salts of the same metal give the same colour.
There is a reason behind it. In the heat of the flame, the electrons of the metal ion take up energy and then give it out again as light — and the energy each metal gives out is fixed, so the colour of that light is fixed too. That is why the test works at all, and why it is a flame test for the metal.
9.4.1 Cleaning And Using The Platinum Wire
The test is carried out on a thin platinum wire. Platinum is chosen for two reasons: it is unreactive, so the salts do not attack it, and it gives no colour of its own to the flame.
The wire is first cleaned thoroughly. It is dipped in concentrated hydrochloric acid and then heated in the non-luminous flame of the burner. The process is repeated until the wire imparts no colour at all to the flame. Only then is it ready for use, because any trace of an earlier salt would give a false result.
Now the clean wire is dipped in concentrated hydrochloric acid again, and then into a small amount of the substance being investigated, so that a little of it sticks to the wire. The acid is not just there to moisten the wire — it converts the salt into its chloride, and chlorides vaporise easily in the flame.
The wire is then introduced into the non-luminous part of the flame, and the colour imparted to the flame is observed. The whole procedure is shown in Fig. 9.4-1. Concentrated hydrochloric acid is corrosive, so it is used in a small watch glass, with eye protection, and the wire is never waved about while wet.
Diagram
Fig. 9.4-1 — Performing A Flame Test Bunsen burner standing on a bench with a tall non-luminous flame; a hand holding a thin platinum wire fixed in a holder, the tip of the wire placed at the edge of the non-luminous part of the flame and imparting colour to it. Below and to the left, a watch glass holding a small heap of the salt moistened with concentrated hydrochloric acid, labelled ‘salt + concentrated HCl’. Label the flame, the non-luminous part of the flame, the platinum wire, the wire holder and the burner.
9.4.2 Colours Imparted By Salts Of Different Metals
The colours imparted by salts of four metals are given in Table 9.7, and are shown in Fig. 9.4-2. Read the second column carefully — it is the one students most often skip.
Table 9.7 — Colour Imparted To A Flame By Salts Of Different Metals
| Colour imparted to the flame (observed with the naked eye) | Colour seen through a blue glass | Name of metal |
|---|---|---|
| Golden yellow (persistent) | Yellow colour vanishes | Sodium (\(\\mathrm{Na^{+}}\)) |
| Violet (lilac) | Violet or pink | Potassium (\(\\mathrm{K^{+}}\)) |
| Brick red (fugitive, that is, it comes and goes) | Light green | Calcium (\(\\mathrm{Ca^{2+}}\)) |
| Peacock bluish green (the wire slowly gets corroded) | Bluish green | Copper (\(\\mathrm{Cu^{2+}}\)) |
The blue glass is used because sodium is everywhere. A trace of sodium from a fingerprint or from dust gives a strong, persistent golden yellow that easily swamps the pale lilac of potassium. Blue glass absorbs that yellow light, so the sodium colour vanishes and the potassium colour comes through as violet or pink.
Two other words in the table are worth pausing on. The sodium colour is described as persistent, meaning it lasts as long as the salt is in the flame. The calcium colour is described as fugitive, meaning it flashes and fades, so you must watch for it rather than wait for it.
Diagram
Fig. 9.4-2 — Colours Imparted To The Flame By Salts Four flame photographs or paintings side by side against a dark background, each showing a burner flame coloured by a salt. From left: a golden yellow flame captioned ‘sodium salt — golden yellow’; a pale violet flame captioned ‘potassium salt — lilac’; a red flame captioned ‘calcium salt — brick red’; a green flame captioned ‘copper salt — bluish green’. Keep all four flames the same size and shape so only the colour differs.
9.5 Some Naturally Occurring Acids
Not every acid comes out of a bottle in a laboratory. A good many are made by living things, and you meet them daily — in the kitchen, in fruit, and even in an insect sting.
Table 9.8 lists some of these acids with the natural source in which each is found. A few are drawn in Fig. 9.5-1.
Table 9.8 — Some Naturally Occurring Acids And Their Sources
| Acid | Natural source |
|---|---|
| Acetic acid | Vinegar |
| Citric acid | Citrus fruits: oranges, lemons |
| Butyric acid | Rancid butter |
| Hydrochloric acid | Gastric juice |
| Formic acid | Sting of bees and ants |
| Lactic acid | Sour milk |
| Malic acid | Apples |
| Oleic acid | Olive oil |
| Stearic acid | Fats |
| Tartaric acid | Grapes, tamarind, apples |
| Uric acid | Urine |
The list explains a few everyday observations. Milk that has been left out turns sour because lactic acid has formed in it. Butter that has gone off smells sharp because of butyric acid. And the sting of an ant burns because formic acid has been injected into the skin.
Diagram
Fig. 9.5-1 — Where Some Natural Acids Are Found Five small labelled drawings arranged in a row on a plain background: a cut lemon labelled ‘citric acid’; a bottle of vinegar labelled ‘acetic acid’; a glass of curdled milk labelled ‘lactic acid’; an apple labelled ‘malic acid’; and an ant labelled ‘formic acid’. Keep the drawings simple and the labels below each item.
9.6 Hard Water And Soft Water
Rub a little soap between wet hands and you expect lather. With some water you get it at once; with other water you rub and rub and get only a greyish, curdy mess. The difference has nothing to do with the soap. It is in the water.
9.6.1 What Makes Water Hard
Water is a good solvent, and as it moves over rock and soil it dissolves a little of whatever it passes through. Where it has picked up certain calcium and magnesium salts, it becomes hard.
Sea water, river water and lake water are all examples of hard water. Each has been in long contact with rock and soil.
Distilled water and rain water are examples of soft water. Rain water has just condensed from vapour and has touched no rock at all; distilled water has been boiled and condensed deliberately, leaving its dissolved salts behind.
9.6.2 Temporary And Permanent Hardness
Not all hardness is the same, because not all the salts behave alike when the water is boiled. Hardness is therefore of two kinds, set out in FC 9.2 and compared in Table 9.9.
Flowchart
FC 9.2 — The Two Kinds Of Hardness Hard water → branch 1: temporary hardness → due to calcium and magnesium bicarbonates → removed by boiling, followed by filtration; branch 2: permanent hardness → due to calcium and magnesium chlorides and sulphates → removed by adding washing soda or caustic soda
Temporary hardness is hardness that boiling can remove. The bicarbonates responsible are unstable to heat, and on boiling they break down into an insoluble carbonate, which can then be filtered off.
Permanent hardness is hardness that boiling cannot remove, because the chlorides and sulphates responsible are unaffected by heat. Removing it needs a chemical, not a kettle.
9.6.3 Distinguishing Hard Water From Soft Water
Two unknown samples, A and B, one containing hard water and the other soft water, are taken separately in beakers. Ordinary soap is rubbed by the hands into the two samples. The result is shown in Fig. 9.6-1.
Observation: one sample forms lather easily, while the other sample does not lather.
Conclusion: the sample that lathers easily with ordinary soap is soft water. The other is hard.
Diagram
Fig. 9.6-1 — Soap In Soft Water And In Hard Water Two identical beakers side by side, each half full of water, with a hand rubbing a piece of soap into each. Left beaker labelled ‘soft water’, drawn with a thick white foam of lather on the surface. Right beaker labelled ‘hard water’, drawn with almost no foam and grey curdy solid floating on the surface and clinging to the sides. Label the lather in the left beaker and the scum in the right.
9.6.4 Distinguishing Temporary From Permanent Hard Water
Two unknown samples are provided, one containing temporary hard water and the other permanent hard water, in separate beakers.
The given samples of water are boiled slowly, and then filtered. Ordinary soap is rubbed by the hands into each filtered sample.
Observation: one sample lathers with soap, but the other one does not.
Conclusion: the boiled and filtered sample that lathers with soap is temporary hard water. The other sample is permanent hard water.
The reason lies in what boiling does to the bicarbonate.
\[\\mathrm{Ca(HCO\_3)\_2}\\text{(aq)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CaCO\_3}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)} + \\mathrm{H\_2O}\\text{(l)}\]
Temporary hard water contains bicarbonate, which gives a precipitate of insoluble carbonate on boiling. This precipitate is removed by filtration, and the water becomes soft.
9.6.5 Softening Hard Water With Washing Soda
Samples of temporary and permanent hard water are taken in separate beakers, and washing soda is added to both. The two solutions are filtered to remove any precipitate formed. Ordinary soap is then rubbed into the filtered solutions.
Observation: both samples lather easily with ordinary soap.
Conclusion, first, that temporary hard water is softened by using washing soda. The carbonate from the washing soda takes the calcium out of solution as insoluble calcium carbonate.
\[\\mathrm{Ca(HCO\_3)\_2}\\text{(aq)} + \\mathrm{Na\_2CO\_3}\\text{(aq)} \\rightarrow \\mathrm{CaCO\_3}\\text{(s)} + 2,\\mathrm{NaHCO\_3}\\text{(aq)}\]
Conclusion, second, that permanent hard water is also softened by using washing soda, since permanent hardness is due to the presence of chlorides and sulphates of calcium and magnesium, and the carbonate removes the calcium from these in exactly the same way.
\[\\mathrm{CaCl\_2}\\text{(aq)} + \\mathrm{Na\_2CO\_3}\\text{(aq)} \\rightarrow \\mathrm{CaCO\_3}\\text{(s)} + 2,\\mathrm{NaCl}\\text{(aq)}\]
Both equations do the same job. Whatever the calcium arrived with, the carbonate takes it out of solution as a precipitate, and what is left behind is a sodium salt, which does not cause hardness. This is why washing soda works on both kinds of hardness, while boiling works on only one.
9.6.6 Soaps And Detergents
Soap is the sodium or potassium salt of an organic fatty acid. It reacts with hard water forming scum, an insoluble curdy solid, which is why ordinary soap is wasted in hard water.
\[\\mathrm{CaSO\_4}\\text{(aq)} + 2,\\mathrm{NaSt}\\text{(aq)} \\rightarrow \\mathrm{CaSt\_2}\\text{(s)} + \\mathrm{Na\_2SO\_4}\\text{(aq)}\]
In this equation St stands for the stearate group, so NaSt is sodium stearate, a common soap, and \(\\mathrm{CaSt\_2}\) is calcium stearate, the scum. Notice what has happened: the soap has not washed anything. It has been used up in reacting with the calcium sulphate present in the hard water. Only when all the calcium and magnesium has been removed in this way can any further soap produce lather.
Detergents are the sodium salts of alkyl sulphonic acids. Detergents contain a sulphonic acid group, \(\\mathrm{-SO\_3H}\), instead of a carboxylic acid group, \(\\mathrm{-COOH}\).
That one difference decides everything. The calcium and magnesium salts of a detergent are soluble in water, so they never come out as a solid. Detergents therefore lather even with hard water, and no scum is formed.
9.6.7 Comparing Soap And Detergent In Hard Water
Two beakers are taken, both containing samples of hard water. In one beaker, ordinary soap is rubbed into the water, while in the other beaker, a detergent is rubbed into the water.
Observation: lather is formed in the beaker in which detergent is used, but lather does not form when ordinary soap is rubbed.
Conclusion: detergents form lather even with hard water, while ordinary soap is wasted due to the formation of scum.
9.7 Experiments On Double Displacement Reactions
Mix two clear solutions and, in certain cases, a solid appears at once. Nothing has been heated and no gas has escaped. What has happened is an exchange.
Each of the salts in solution is made of two parts — a metal and a radical, which is a group of atoms that stays together as a single unit and carries a charge, such as sulphate, \(\\mathrm{SO\_4^{2-}}\), or nitrate, \(\\mathrm{NO\_3^{-}}\). When two such solutions meet, the partners can swap.
Whether a precipitate appears is decided entirely by solubility, which you can read off Table 9.6. If both new compounds are soluble, nothing is seen at all.
9.7.1 Sodium Sulphate With Barium Chloride
Take about 3 grams of sodium sulphate in a test tube, add 5 ml of distilled water to it and stir it well to dissolve. This is an aqueous solution of sodium sulphate. In another test tube, make a solution by adding distilled water to about 3 g of barium chloride.
To the sodium sulphate solution, add barium chloride solution slowly, as shown in Fig. 9.7-1.
\[\\mathrm{Na\_2SO\_4}\\text{(aq)} + \\mathrm{BaCl\_2}\\text{(aq)} \\rightarrow \\mathrm{BaSO\_4}\\text{(s)} + 2,\\mathrm{NaCl}\\text{(aq)}\]
Observation: a white, heavy, fine crystalline precipitate of barium sulphate is formed immediately, due to a mutual exchange of radicals between the reactants. This is a double displacement reaction.
On adding dilute hydrochloric acid to the white precipitate of barium sulphate, the precipitate remains insoluble. That second observation matters, because several other white precipitates do dissolve in acid. Barium chloride solution is therefore a reagent which can be used to test for sulphate radicals in any soluble sulphate.
Diagram
Fig. 9.7-1 — Formation Of Barium Sulphate By Double Displacement Test tube held in a hand at a slight slant, containing a colourless solution labelled ‘sodium sulphate solution (Na₂SO₄)’; a second tube tilted above it pouring in a colourless solution labelled ‘barium chloride solution (BaCl₂)’; a dense white solid drawn settling at the bottom of the first tube, labelled ‘barium sulphate, white precipitate’.
9.7.2 Silver Nitrate With Sodium Chloride
Make an aqueous solution of sodium chloride in a test tube. To this, add a few drops of silver nitrate solution, as shown in Fig. 9.7-2.
\[\\mathrm{NaCl}\\text{(aq)} + \\mathrm{AgNO\_3}\\text{(aq)} \\rightarrow \\mathrm{AgCl}\\text{(s)} + \\mathrm{NaNO\_3}\\text{(aq)}\]
Observation: a white precipitate of silver chloride, curd-like in appearance, is immediately formed as the mutual exchange of radicals occurs between the reactants sodium chloride and silver nitrate. A double displacement reaction occurs.
This white precipitate of silver chloride dissolves on adding ammonium hydroxide solution. You have met that pair of observations before, in 9.1.5, where they identified hydrogen chloride gas; here the same reaction is used to test for chloride radicals in solution.
Diagram
Fig. 9.7-2 — Formation Of Silver Chloride By Double Displacement Test tube held in a hand, containing a colourless solution labelled ‘sodium chloride solution (NaCl)’; a dropper or second tube above it adding a colourless solution labelled ‘silver nitrate solution (AgNO₃)’; a white curdy solid drawn at the bottom of the tube, labelled ‘silver chloride, white precipitate’.
9.7.3 Lead Nitrate With Potassium Iodide
Dissolve a small amount of lead nitrate in a test tube or beaker in distilled water to make a clear, colourless solution. Take about 5 ml of this solution in a clean test tube. To this, add potassium iodide solution with the help of a dropper, as in Fig. 9.7-3.
\[\\mathrm{Pb(NO\_3)\_2}\\text{(aq)} + 2,\\mathrm{KI}\\text{(aq)} \\rightarrow \\mathrm{PbI\_2}\\text{(s)} + 2,\\mathrm{KNO\_3}\\text{(aq)}\]
Observation: a brilliant yellow precipitate of lead iodide is immediately formed, which settles at the bottom of the test tube. Both solutions were colourless before mixing, so the colour has arrived with the new compound.
This is due to the exchange of radicals of the reactants lead nitrate and potassium iodide, thus forming lead iodide — again a double displacement reaction.
If the solutions are heated, the lead iodide dissolves. Upon cooling, it recrystallises into shiny golden yellow crystals.
Lead nitrate is toxic and harmful. While doing this experiment, wear eye protection and gloves, handle the chemical with care, and do not pour the contents of the tube down the sink.
Diagram
Fig. 9.7-3 — Yellow Precipitate Of Lead Iodide Test tube held in a hand, containing a colourless solution labelled ‘lead nitrate solution (Pb(NO₃)₂)’; a dropper above it releasing drops of a colourless solution labelled ‘potassium iodide solution (KI)’; a bright yellow solid drawn settling at the bottom, labelled ‘lead(II) iodide, yellow precipitate’.
The three reactions of this section are compared in Table 9.11. Read the last row: the further test is what turns a precipitate into a positive identification.
9.8 Water Pollution
A river does not stop being water when waste is tipped into it. It stops being useful. That is the idea behind the definition used here.
Water pollution affects the health of living organisms and may offend the sense of sight, taste or smell. It refers to any type of aquatic contamination which renders the water unusable for living organisms.
9.8.1 What Water Pollution Means
Polluted water can often be identified without any apparatus at all. Look for the following.
Foul smell from water.
Bad taste of drinking water.
Oil and grease floating on the surface of water.
Excessive growth of algae and weeds in water.
All over the world, freshwater resources are threatened not only by over-exploitation and poor management but also by ecological degradation. The main reasons for fresh water pollution are these.
- Discharge of untreated waste.
- Dumping of industrial effluents.
- Run-off from agricultural fields.
- Industrial growth and urbanization.
- Increasing use of synthetic organic substances.
It is generally accepted that developed countries suffer more from problems of chemical discharge into water sources, while developing countries face more problems of agricultural run-off into water sources. Polluted water causes water-borne diseases and leads to other health-related problems.
9.8.2 Sources Of Water Pollution
The main sources of water pollution are given below, and several of them appear in Fig. 9.8-1.
Household detergents that are used for washing purposes.
Sewage and domestic wastes.
Industrial wastes, which contain heavy metals, their salts and organic compounds. All these get accumulated in marine food chains, so a substance released in small amounts can end up concentrated in the fish that people eat.
Oil spills, which poison sea life.
Thermal pollution, which raises the temperature of water and depletes the oxygen present in it, killing fishes and other aquatic organisms. Warm water cannot hold as much dissolved oxygen as cold water, so heat by itself is enough to make a stretch of river uninhabitable.
Diagram
Fig. 9.8-1 — Sources Of Water Pollution Wide landscape scene of a river running from left to right. On the far bank, a factory with chimneys and a pipe discharging effluent into the river, and beside it a cooling tower returning warm water. On the near bank, a drain carrying household waste water, a field sloping to the water’s edge, and floating rubbish with a film of oil on the surface. Label the industrial effluent pipe, the warm water outflow, the domestic drain, the run-off from the field and the floating oil and waste.
9.8.3 Judging Water Quality
Saying that water is polluted is not enough. To act on it, you have to measure something. Determination of water quality may be done in two ways.
Select a water body and survey its specific parameters throughout the year.
Select a specific parameter and measure it in different types of water bodies.
Six parameters are used at this level: pH value, bacteria present, hardness of water, dissolved oxygen, biological oxygen demand and turbidity. Each is described below.
9.8.3.1 pH Value
The acidic and basic strengths of solutions are measured on the hydrogen ion scale, or pH scale, which gives them a value called the pH value. The concentration of \(\\mathrm{H^{+}}\) ions present in a solution decides whether it is acidic, alkaline or neutral.
The pH value is obtained with the help of a universal indicator, which changes colour at different pH values, as shown in Fig. 9.8-2. The pH can also be studied with the help of a pH meter, or read by matching a wet pH paper against a standard printed colour scale.
- A solution of pH 7 is neutral, as pure water is.
- A solution of pH less than 7 is acidic.
- A solution of pH more than 7 is basic.
- The higher the pH above 7, the stronger the base; the lower the pH below 7, the stronger the acid.
Pollution can change the pH of water, which in turn can harm animals and plants living in it. For instance, water coming out of an abandoned coal mine can have a pH value as low as 2, which means it is very acidic and would affect any creature surviving in or on it.
Non-polluted water is water that has a pH value of 7 or close to 7. If the pH value of water is much higher or much lower than 7, that water is polluted and potentially harmful.
The pH values of some familiar substances are given in Table 9.12.
Diagram
Fig. 9.8-2 — Colour Changes In A Universal Indicator At Different pH Values Horizontal scale numbered 0 to 14 at equal intervals, with a coloured band above it: red at 0 to 1, pink at 2 to 3, orange and yellow at 4 to 6, green at 7, blue at 8 to 9, indigo at 10 to 12 and violet at 13 to 14. Mark 7 as ‘neutral’. Draw an arrow to the left of 7 labelled ‘acidic nature increases’ and an arrow to the right of 7 labelled ‘basic nature increases’. Print the colour names below the band.
Table 9.12 — pH Values Of Some Common Substances
| Substance | pH value |
|---|---|
| Hydrochloric acid | 1.0 |
| Sulphuric acid | 1.2 |
| Citric acid | 2.3 |
| Lime juice | 2.4 |
| Acetic acid | 2.9 |
| Carbonic acid | 3.0 |
| Pure water | 7.0 |
| Human blood | 7.3 |
| Sodium bicarbonate solution | 8.4 |
| Milk of magnesia | 10.5 |
| Sodium carbonate solution | 11.6 |
| Sodium hydroxide | 13.0 |
9.8.3.2 Bacteria
Biological water quality assessment can be done by measuring the growth of bacteria in a water body.
Growth of bacteria can be measured by taking a sample of water in a petri dish and observing it under a microscope three or four times at regular intervals. Repeating the observation is the point: one look tells you what is present, but only a series of looks tells you whether it is multiplying.
Water free of bacteria is non-polluted water, whereas water with bacterial growth is considered polluted.
9.8.3.3 Hardness
When water produces little or no lather with soap, it is considered hard water, as you saw in 9.6. Hard water is unsuitable for washing clothes, for geysers, and for drinking. Hard water is thus regarded as polluted water.
9.8.3.4 Dissolved Oxygen
Dissolved oxygen is oxygen from the air that has dissolved in the water, and it is responsible for the survival of aquatic life. Fish and other aquatic creatures breathe it through their gills; they cannot use the oxygen locked inside the water molecules themselves.
Flowing water, that we see in mountain streams and large rivers, tends to contain a lot of dissolved oxygen, because it is constantly churned against the air. Stagnant water contains little of it.
If the quantity of oxygen dissolved in stream or river water is less than the quantity of dissolved oxygen in normal water, the water is said to be polluted. The quantity of dissolved oxygen is measured with the instrument shown in Fig. 9.8-3.
Diagram
Fig. 9.8-3 — Instrument For Measuring Dissolved Oxygen Hand-held dissolved oxygen meter on a plain background: a rectangular display unit with a small screen and a few buttons, a long coiled cable running from it, and a cylindrical probe at the far end of the cable. Label the display unit, the cable and the probe, with a note that the probe is lowered into the water sample.
9.8.3.5 Biological Oxygen Demand
Biological oxygen demand, written BOD, is a measure of the oxygen utilized by micro-organisms during the oxidation of organic substances.
Micro-organisms present in a polluted water sample use oxygen to oxidize waste, and thus BOD increases in polluted water. The amount of oxygen present in pure water is measured first, and then in the polluted water. The decrease in the amount of oxygen is the BOD value of the polluted water.
Biological oxygen demand increases in polluted water. Read the two quantities together and the picture is complete: dissolved oxygen tells you what the water still holds, while BOD tells you how fast it is being used up.
9.8.3.6 Turbidity
Turbidity is the amount of particulate matter suspended in water.
Turbidity measures the scattering effect that suspended solids have on light. The higher the intensity of the scattered light, the higher the turbidity. Materials that are responsible for turbidity include the following.
- Clay.
- Waste matter containing insoluble substances.
- Microscopic organisms.
As the turbidity of water increases, its ability to support aquatic life decreases. Suspended particles cut off the light that water plants need, and those plants are where much of the oxygen in a healthy pond comes from.
9.8.4 Ways To Control Water Pollution
Each measure below removes a particular kind of pollutant, and each is chosen for the kind of matter it has to deal with. A sensible order of treatment is set out in FC 9.3.
- Sewage should be treated fully before its contents are released into water bodies.
- Chemicals released from factories should be treated to neutralize their toxic contents.
- Heavier floating solids can be removed by gravity settlement and screening.
- Organic matter should be removed by oxidation or precipitation.
- Pathogens can be destroyed by ultraviolet radiation.
Flowchart
FC 9.3 — Treating Waste Water Before It Is Released Untreated waste water → screening to hold back large floating solids → gravity settlement of heavier solids → removal of organic matter by oxidation or precipitation → treatment of factory chemicals to neutralize their toxic contents → destruction of pathogens by ultraviolet radiation → treated water released into the water body
9.8.5 Surveying A Water Body In Your Locality
The object of such a survey is to find out the sources of pollution of water bodies in one’s locality and to determine the quality of the water.
For the survey described here, a locality where a drain exists was chosen. The major sources of water pollution found were household wastes; detergents that people wash clothes with; wastes from shops; animal excreta; wastes generated by a restaurant near the drain, including plastic packaging, bottles, cans, food stuff and clothing; washes from a nearby hospital; and wastes from factories.
In order to determine the quality of the water, \(100\\ \\text{cm}^3\) of water was taken in a beaker for finding its pH value, and \(500\\ \\text{cm}^3\) of water was taken in another beaker for measuring the dissolved oxygen in it. The data obtained is given in Table 9.13.
Table 9.13 — pH And Dissolved Oxygen Of The Water Sampled, April To September
| Month | April | May | June | July | Aug. | Sept. |
|---|---|---|---|---|---|---|
| pH | 6.0 | 5.9 | 6.6 | 7.8 | 7.0 | 6.8 |
| DO (mg/l) | 4.1 | 3.5 | 2.1 | 1.7 | 1.7 | 1.5 |
Two conclusions follow from these figures.
If the pH of the water is less than 7, that is, the water is acidic, it would affect all the living creatures present there. Acidity results mainly from the discharge of industrial effluents.
When the dissolved oxygen decreases in a water body, mainly due to the carbon-containing matter present in it, it becomes difficult for aquatic organisms that take dissolved oxygen to survive.
Thus water quality is considered to be good when there is more dissolved oxygen, and bad when there is less dissolved oxygen. Notice that in Table 9.13 the dissolved oxygen falls steadily from April to September while the pH moves up and down. That is exactly why more than one parameter is measured, and why a survey is spread across months rather than done once.
9.9 Summary
An unknown substance is identified from what is observed: its colour, its odour, its physical state, what it does on heating, and how it behaves towards a few reagents.
Ten gases are met at this level: hydrogen, oxygen, carbon dioxide, chlorine, hydrogen chloride, sulphur dioxide, hydrogen sulphide, ammonia, water vapour and nitrogen dioxide.
Hydrogen is colourless, odourless and neutral to litmus. It burns with a pale blue flame, and mixed with air it burns with a pop sound.
Oxygen is colourless, odourless and neutral. It rekindles a glowing splinter and is absorbed by alkaline pyrogallol solution, turning it dark brown.
Carbon dioxide is colourless and odourless, turns moist blue litmus faint red, and turns lime water milky. The milkiness disappears with excess gas.
\[\\mathrm{Ca(OH)\_2}\\text{(aq)} + \\mathrm{CO\_2}\\text{(g)} \\rightarrow \\mathrm{CaCO\_3}\\text{(s)} + \\mathrm{H\_2O}\\text{(l)}\]
Chlorine is greenish yellow with a pungent choking odour. It turns moist blue litmus red and then bleaches it permanently, and turns moist starch iodide paper blue black.
Hydrogen chloride is colourless and pungent, turns moist blue litmus red, and gives dense white fumes of ammonium chloride with a rod dipped in ammonia solution.
Sulphur dioxide is colourless with the smell of burning sulphur. It turns lime water milky like carbon dioxide, but it also decolourises acidified potassium permanganate and turns acidified potassium dichromate from orange to green.
Hydrogen sulphide is colourless with a rotten-egg smell and turns lead acetate paper black.
Ammonia is colourless with a sharp pungent smell and is the only basic gas of the ten; it turns moist red litmus blue and turns Nessler’s reagent brown.
Water vapour condenses to a colourless neutral liquid that turns white anhydrous copper sulphate blue and blue cobalt chloride paper pink.
Nitrogen dioxide is brown, has an irritating odour, is non-combustible, and liberates iodine from potassium iodide solution.
On heating, a solid may lose its water of crystallisation, decompose with the loss of a gas, change colour, or sublime.
Carbonates other than sodium and potassium carbonate give carbon dioxide on heating, leaving the metal oxide.
\[\\mathrm{ZnCO\_3}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } \\mathrm{ZnO}\\text{(s)} + \\mathrm{CO\_2}\\text{(g)}\]
Hydrated salts give water vapour on heating. Copper sulphate crystals lose water first and, on stronger heating, decompose further to black copper oxide with sulphur dioxide and oxygen.
Heavy metal nitrates give the metal oxide, nitrogen dioxide and oxygen on heating. Sodium nitrate and potassium nitrate give oxygen but no nitrogen dioxide.
\[2,\\mathrm{Pb(NO\_3)\_2}\\text{(s)} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{PbO}\\text{(s)} + 4,\\mathrm{NO\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)}\]
Ammonium chloride and iodine leave no residue on heating. Ammonium dichromate decomposes violently, leaving a greenish grey mass of chromium(III) oxide, with water vapour and nitrogen.
The colour of the residue while hot and again when cold is an important clue: zinc oxide is yellow when hot and white when cold; lead oxide is reddish brown when hot and yellow when cold, and stains the glass; copper oxide is black.
Colour points to the metal ion: blue or bluish green to \(\\mathrm{Cu^{2+}}\), light green to \(\\mathrm{Fe^{2+}}\), yellow or yellowish brown to \(\\mathrm{Fe^{3+}}\); a white or colourless solid may contain any of several ions.
Odour points to the negative radical: ammonia to an ammonium salt, rotten eggs to a sulphide, burning sulphur to a sulphite.
Physical state is a clue too: an amorphous salt suggests a carbonate, and a hygroscopic or deliquescent salt suggests a chloride or a nitrate.
A dry test examines the solid itself — colour, density, physical state, dry heating and the flame test. A wet test adds reagents, usually to a solution, and looks for a precipitate or a colour change.
All nitrates, and all sodium and potassium salts, are soluble in water. Carbonates, sulphides, sulphites and phosphates are insoluble for most metals other than sodium, potassium and ammonium.
In the flame test, a clean platinum wire is moistened with concentrated hydrochloric acid, dipped in the substance and held in the non-luminous flame. Sodium gives golden yellow, potassium lilac, calcium brick red and copper bluish green.
A blue glass is used because it absorbs the strong yellow of sodium, allowing the violet of potassium to be seen.
Many acids occur naturally — acetic acid in vinegar, citric acid in citrus fruits, lactic acid in sour milk, formic acid in the sting of bees and ants, and tartaric acid in tamarind and grapes.
Hard water contains dissolved bicarbonates, sulphates and chlorides of calcium and magnesium, and does not form lather with ordinary soap. Soft water is free of these salts and lathers readily.
Temporary hardness is due to bicarbonates of calcium and magnesium and is removed by boiling followed by filtration. Permanent hardness is due to their chlorides and sulphates and is removed by adding washing soda or caustic soda.
\[\\mathrm{CaCl\_2}\\text{(aq)} + \\mathrm{Na\_2CO\_3}\\text{(aq)} \\rightarrow \\mathrm{CaCO\_3}\\text{(s)} + 2,\\mathrm{NaCl}\\text{(aq)}\]
Soap is the sodium or potassium salt of an organic fatty acid, and forms an insoluble scum with hard water. Detergents are sodium salts of alkyl sulphonic acids, carrying a \(\\mathrm{-SO\_3H}\) group instead of a \(\\mathrm{-COOH}\) group; their calcium and magnesium salts are soluble, so they lather even in hard water.
In a double displacement reaction, two compounds exchange their radicals, and a precipitate appears when one of the new compounds is insoluble.
\[\\mathrm{Na\_2SO\_4}\\text{(aq)} + \\mathrm{BaCl\_2}\\text{(aq)} \\rightarrow \\mathrm{BaSO\_4}\\text{(s)} + 2,\\mathrm{NaCl}\\text{(aq)}\]
Barium chloride solution tests for a sulphate radical, giving a white precipitate that is insoluble in dilute hydrochloric acid. Silver nitrate solution tests for a chloride radical, giving a white curdy precipitate that dissolves in ammonium hydroxide. Lead nitrate with potassium iodide gives a brilliant yellow precipitate of lead iodide.
Water pollution is the introduction of substances of such character and in such quantity that they alter the natural quality of water and impair its usefulness.
Its main sources are household detergents, sewage and domestic wastes, industrial wastes, oil spills and thermal pollution.
Water quality is judged from six parameters: pH value, bacteria present, hardness, dissolved oxygen, biological oxygen demand and turbidity.
Non-polluted water has a pH of 7 or close to it. Pollution lowers the dissolved oxygen and raises the biological oxygen demand, and both changes make survival harder for aquatic life.
Water pollution is controlled by treating sewage fully, neutralizing toxic factory chemicals, removing heavy solids by screening and gravity settlement, removing organic matter by oxidation or precipitation, and destroying pathogens with ultraviolet radiation.
Flowchart
FC 9.4 — The Route To Identifying An Unknown Substance Unknown substance → dry tests: note colour → note odour → note physical state → heat a little in a dry test tube → identify the gas or vapour evolved and the colour of the residue, hot and cold → apply the flame test for the metal → wet tests: dissolve in water → add a reagent such as barium chloride or silver nitrate → identify the precipitate formed → name the substance
9.10 Key Terms
Anhydrous — describing a salt from which the water of crystallisation has been driven out.
Biological oxygen demand — a measure of the oxygen used up by micro-organisms while oxidising organic matter in water; it rises as water becomes more polluted.
Bleaching — the destruction of a colour; permanent in the case of chlorine, temporary in the case of sulphur dioxide.
Deliquescent — describing a substance that absorbs so much moisture from the air that it finally dissolves in the water it has taken up.
Detergents — the sodium salts of alkyl sulphonic acids, which carry a \(\\mathrm{-SO\_3H}\) group and lather even in hard water.
Dissolved oxygen — the oxygen present in solution in water, on which aquatic life depends.
Double displacement reaction — a reaction in which two compounds exchange their radicals to form two new compounds.
Flame test — a dry test in which a salt on a clean platinum wire is held in a non-luminous flame and identified from the colour it imparts.
Hard water — water containing dissolved bicarbonates, sulphates and chlorides of calcium and magnesium, which does not lather with ordinary soap.
Hygroscopic — describing a substance that absorbs moisture from the air and becomes damp, but remains solid.
Lime water — a clear, saturated solution of calcium hydroxide, used to test for carbon dioxide.
Permanent hardness — hardness caused by chlorides and sulphates of calcium and magnesium, which boiling cannot remove.
pH value — the value on the hydrogen ion scale that expresses how acidic or basic a solution is; 7 is neutral, below 7 acidic and above 7 basic.
Precipitate — an insoluble solid that separates out when two solutions react.
Radical — a group of atoms that moves as a single unit in a reaction and carries a charge.
Scum — the insoluble curdy solid formed when soap reacts with the calcium and magnesium salts in hard water.
Soap — the sodium or potassium salt of an organic fatty acid.
Soft water — water free of dissolved calcium and magnesium salts, which lathers readily with ordinary soap.
Sublimation — the direct change of a solid into vapour and back into solid, without passing through the liquid state.
Temporary hardness — hardness caused by bicarbonates of calcium and magnesium, which boiling followed by filtration removes.
Turbidity — the amount of particulate matter suspended in water, judged from the scattering of light.
Water of crystallisation — the fixed quantity of water built into the structure of a crystal.
Water pollution — the introduction of substances of such character and in such quantity that they alter the natural quality of water and impair its usefulness.
Wet test — a test in which reagents are added, usually to a solution of the substance, and a precipitate or colour change is looked for; the counterpart of a dry test, which examines the solid itself.
9.11 Points To Ponder
Every test in this chapter turns something invisible into something you can see, hear or smell. Choose three tests and decide, in each case, what the chemistry has actually converted the evidence into — a colour, a solid, a sound, or a smell. Which kind of evidence do you trust most, and why?
A single observation almost never identifies a substance, yet three or four together usually do. Why should that be so? Try to state what a good confirmatory test must do that a first test need not.
Boiling removes temporary hardness but not permanent hardness, while washing soda removes both. What does that tell you about the difference between changing a substance by heat and changing it by adding another substance?
Hard water is described in this chapter as polluted water, even though the calcium and magnesium salts in it are dissolved by natural processes and are not toxic. Is ‘polluted’ the right word for it? Consider what the word is being used to mean here.
A precipitate forms only when one product of an exchange happens to be insoluble. If every compound in the world were soluble in water, which of the tests in this chapter would still work, and which would become useless?
Sodium’s golden yellow flame is so strong that it hides the colours of other metals, which is why a blue glass is used. Where else in this chapter does one substance interfere with the detection of another, and how is that interference dealt with?
9.12 Objective Questions
- The gas that turns moist starch iodide paper blue black is: [1]
- sulphur dioxide
- chlorine
- hydrogen sulphide
- ammonia
- A residue that is yellow when hot and white when cold is: [1]
- lead oxide
- copper oxide
- zinc oxide
- chromium(III) oxide
- The only basic gas among the ten gases studied in this chapter is: [1]
- hydrogen
- ammonia
- nitrogen dioxide
- water vapour
- The reagent used to test for a sulphate radical in a soluble sulphate is: [1]
- silver nitrate solution
- lime water
- barium chloride solution
- potassium iodide solution
- The substance that leaves no residue at all on strong heating is: [1]
- lead nitrate
- copper carbonate
- washing soda
- ammonium chloride
- Seen through a blue glass, the flame colour given by a calcium salt appears: [1]
- light green
- violet or pink
- bluish green
- colourless
- Choose the correct option. [1] Assertion (A): Detergents form lather even with hard water. Reason (R): The calcium and magnesium salts of detergents are soluble in water.
- Both A and R are true and R is the correct explanation of A.
- Both A and R are true but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
- Choose the correct option. [1] Assertion (A): Temporary hard water becomes soft on boiling and filtering. Reason (R): Calcium bicarbonate decomposes on boiling to give insoluble calcium carbonate.
- Both A and R are true and R is the correct explanation of A.
- Both A and R are true but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
- Choose the correct option. [1] Assertion (A): Carbon dioxide turns acidified potassium dichromate paper green. Reason (R): Carbon dioxide is an acidic gas.
- Both A and R are true and R is the correct explanation of A.
- Both A and R are true but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
- Match each gas in Column A with its identification in Column B. [5]
| Column A | Column B |
|---|---|
| a. Hydrogen sulphide | i. Turns blue cobalt chloride paper pink |
| b. Chlorine | ii. Gives dense white fumes with a rod dipped in concentrated hydrochloric acid |
| c. Ammonia | iii. Turns lead acetate paper black |
| d. Nitrogen dioxide | iv. Greenish yellow, and bleaches moist blue litmus |
| e. Water vapour | v. Brown gas, non-combustible |
- Complete the following. [4]
- Zinc oxide is ______ when hot and white when cold.
- The gas that rekindles a glowing wooden splinter is ______.
- Permanent hardness is due to the ______ and sulphates of calcium and magnesium.
- A white precipitate of silver chloride dissolves in excess ______.
- State whether each statement is true or false. [4]
- Sodium carbonate gives off carbon dioxide on strong heating.
- The bleaching action of sulphur dioxide is permanent.
- All nitrates are soluble in water.
- Rain water is an example of soft water.
- Name each of these. [5]
- The gas with the smell of rotten eggs.
- The residue left when copper carbonate is heated strongly.
- The reagent used to identify a chloride radical in solution.
- The metal whose salt imparts a lilac colour to a flame.
- The gas evolved on heating ammonium dichromate, which neither burns nor supports combustion.
- Give one word or phrase for each of the following. [4]
- The water built into the structure of a crystal.
- A substance that absorbs so much moisture from the air that it dissolves in it.
- The direct change of a solid into vapour without melting.
- The insoluble curdy solid formed when soap is used in hard water.
- Pick the odd one out in each group and give a reason for your choice. [3]
- zinc carbonate, copper carbonate, sodium carbonate, calcium carbonate
- hydrogen, oxygen, water vapour, chlorine
- sodium chloride, barium chloride, silver chloride, potassium chloride
9.13 Check Your Concepts
Define sublimation, and give one example of a substance from this chapter that sublimes. [2]
State what you would observe when a burning splinter is brought to the mouth of a jar of hydrogen, and when a glowing splinter is lowered into a jar of oxygen. [2]
Give a reason why a platinum wire, and not an iron wire, is used for the flame test. [2]
Give a reason why the flame given by a potassium salt is viewed through a blue glass. [2]
Differentiate between temporary hardness and permanent hardness, stating the basis of difference. [3]
Differentiate between a dry test and a wet test, stating the basis of difference. [3]
Why does carbon dioxide turn lime water milky, and why does the milkiness disappear when excess gas is passed? [3]
Explain, with one example, what is meant by a double displacement reaction. [3]
Explain why ordinary soap is wasted in hard water while a detergent is not. Refer to the group each contains. [4]
What would happen if copper sulphate crystals were heated first gently and then very strongly? Describe the observations at each stage and name the products. [4]
Give a reason for each of the following. [4]
- A glowing splint bursts into flame in the gases obtained by heating lead nitrate, although nitrogen dioxide does not support combustion.
- Hydrogen sulphide leaves acidified potassium permanganate solution cloudy, while sulphur dioxide leaves it clear.
You are given three unlabelled samples: soft water, temporary hard water and permanent hard water. Using only soap, a burner and washing soda, describe how you would identify each. [5]
Name any five parameters used to judge the quality of water, and state briefly what each one tells you about the sample. [5]
9.14 Equations, Conversions & Identification
Write balanced equations for the two stages in the action of heat on zinc nitrate hexahydrate. [3]
Complete and balance the following equations. [3]
- \(\\mathrm{MnO\_2} + \\mathrm{HCl} \\rightarrow\)
- \(\\mathrm{NH\_4Cl} + \\mathrm{Ca(OH)\_2} \\rightarrow\)
- \(\\mathrm{FeS} + \\mathrm{H\_2SO\_4} \\rightarrow\)
Name the gas evolved in each case below, and give one test for it. [3]
- Dilute hydrochloric acid is added to sodium sulphite.
- Ammonium chloride is heated with sodium hydroxide.
- Dilute sulphuric acid is added to zinc.
State what you would see in each of the following. [3]
- Chlorine is passed through potassium iodide solution.
- Hydrogen sulphide is passed through lead nitrate solution.
- Water vapour is brought into contact with blue cobalt chloride paper.
How would you obtain each of the following? Write the balanced equation in each case. [3]
- Copper oxide from copper nitrate.
- Soft water from temporary hard water.
- Barium sulphate from a solution of sodium sulphate.
A white crystalline solid is heated in a dry test tube. Dense white fumes are given off, a white solid forms on the cooler part of the tube, and no residue is left. The vapour contains one acidic and one basic gas. Identify the solid and write the equation for its behaviour on heating. [2]
Look at Fig. 9.4-1. Name the labelled parts of the apparatus and state the purpose of each. [3]
Look at Fig. 9.2-1. State why the mouth of the test tube is pointed away from everyone, and name the part of the tube on which a sublimate would be deposited. [2]