Chapter 2: Chemical Changes and Reactions
Nearly three hundred years ago, a German professor named Johann Heinrich Schulze was working with a bottle of white paste — chalk stirred into a solution of silver in nitric acid. He left the bottle standing near a window. When he came back to it, one side of the paste had turned dark purple, almost black. It was the side that had faced the light. The other side was still white.
Schulze suspected the warmth of the sun. So he tested that idea. He warmed the bottle by a fire in a darkened room, and nothing happened at all. Light, and not heat, was doing this.
Then he did something clever. He cut letters out of paper, pasted them around the bottle as a stencil, and set it back in the sunlight. When he peeled the paper away, the letters stood out white against a dark background. He had written with sunlight.
Schulze did not know what was going on inside that paste. We do. The silver compound was being pulled apart, and one of the pieces left behind was silver metal itself — that was the dark colour. The light had not simply shone on the substance. It had turned it permanently into something new.
That is a chemical change. This chapter is about how such changes are started, how you recognise one, what kinds there are, and where their energy comes from.
Diagram
Fig. 2.0-1 — Writing With Sunlight A stoppered glass bottle standing on a windowsill with sunlight falling on it from the left; paper letters spelling a short word are pasted on the bottle’s surface; label the white paste inside, the pasted paper stencil, the darkened region of paste where light has fallen, and the pale letter shapes left behind where the paper shielded the paste.
2.1 Chemical Change
Leave a slice of apple on a plate for an hour and its cut face turns brown. Leave milk out of the refrigerator on a warm day and it curdles and smells sour. In both cases you cannot get the original substance back by cooling it, drying it or stirring it. Something new has taken its place.
A change of this kind is called a chemical change — a permanent change in which the composition of a substance is altered, and one or more new substances with different compositions and different properties are formed. The new substance is not a disguised form of the old one. It has its own colour, its own smell and its own behaviour.
The word ‘permanent’ is doing real work in that sentence. A chemical change is not easily reversible. You cannot un-brown the apple, and no amount of chilling will turn sour milk sweet again.
2.1.1 Characteristics Of A Chemical Change
Chemists recognise a chemical change by four things it always does.
- It is permanent and irreversible. The change cannot be undone by simply reversing the conditions that brought it about.
- The composition of the substance changes. The atoms present are rearranged into new arrangements, so the substance that results is a different one.
- It is either exothermic, where heat is given out, or endothermic, where heat is taken in. Energy is always involved one way or the other, and section 2.9 returns to this in detail.
- The total mass of the reactants and the total mass of the products are the same. This is the law of conservation of mass.
2.1.2 Chemical Bonds And Chemical Reactions
Behind every chemical change lies a chemical reaction. To see what a reaction is, you first need to know what holds a substance together in the first place.
The atoms in a molecule are not lying loose beside one another. They are held by a force of attraction called a chemical bond. Break that force and the atoms are free to join other partners.
A chemical reaction, then, is a matter of unmaking and remaking these bonds. The bonds in the starting substances are broken, and the freed atoms form new bonds in new arrangements. The starting substances are called the reactants, and the substances formed are called the products. This is set out in Fig. 2.1-1.
Diagram
Fig. 2.1-1 — Bonds Break, Then New Bonds Form Three panels in a row, using coloured spheres for atoms and short lines for bonds. Panel 1, labelled ‘reactants’: two molecules, one made of two identical dark spheres joined by a bond, the other of two identical light spheres joined by a bond. Panel 2, labelled ‘bonds break’: the same four spheres shown separated, with the broken bonds drawn as short torn lines. Panel 3, labelled ‘products’: two new molecules, each made of one dark sphere bonded to one light sphere. Add a caption strip beneath reading ‘same atoms, new partners, new substance’.
Notice how neatly this explains the law you met a moment ago. Every atom that goes into a reaction comes out of it, so not a single grain of mass can go missing.
Bonds do not break on their own, and atoms cannot form new bonds unless they meet. A chemical reaction therefore happens only when particles collide. Particles collide when the reactants are brought into close contact, or when energy is supplied to make them move and vibrate more violently.
That single idea tells you what to look for next. If collisions are what a reaction needs, then anything that brings particles together or feeds them energy can set a chemical change going — and the following section lists exactly those conditions.
2.2 Conditions Necessary For A Chemical Change
A reaction needs its particles to meet, and it needs bonds to break. Nothing else will do. So a chemical change begins only when the reactants are brought into close contact, or when energy is fed into them, or both.
That gives chemists a short list of practical conditions. One or more of the following is needed before a chemical change will take place: close contact, being in solution, heat, light, electricity, pressure, and the presence of a catalyst. The rest of this section takes them one at a time.
2.2.1 Close Contact Between Solids
Some substances react as soon as their solid particles are pressed together. No water, no flame, no current — simply contact.
Iodine and phosphorus are the extreme case. The two powders react explosively the moment they touch, which is why they are never brought together in a school laboratory and are stored well apart.
A gentler example uses two white powders. Lead nitrate and potassium iodide are both white, yet a bright yellow substance, lead iodide, appears where their particles meet.
lead nitrate + potassium iodide → potassium nitrate + lead iodide
\[\mathrm{Pb(NO_3)_2}\text{(s)} + 2\,\mathrm{KI}\text{(s)} \rightarrow 2\,\mathrm{KNO_3}\text{(s)} + \mathrm{PbI_2}\text{(s)}\]
The yellow colour is the evidence. Neither starting powder was yellow, so a new substance must have been formed.
2.2.2 Reaction In Solution
Many substances that ignore each other as dry solids react at once when they are dissolved. Water pulls the solid apart into free particles, and free particles can move, meet and collide. A molten substance behaves in the same way, for the same reason.
Oxalic acid crystals and sodium carbonate are a clear case. Shake the two dry powders together and nothing happens; add water and they react.
The second example is one you will meet again in this chapter. Sodium chloride solution and silver nitrate solution are both colourless, but on mixing, a white solid appears immediately.
sodium chloride + silver nitrate → silver chloride + sodium nitrate
\[\mathrm{NaCl}\text{(aq)} + \mathrm{AgNO_3}\text{(aq)} \rightarrow \mathrm{AgCl}\!\downarrow + \mathrm{NaNO_3}\text{(aq)}\]
The downward arrow marks the white silver chloride, which does not dissolve and settles out of the mixture.
2.2.3 Heat
Some reactions will not start at ordinary temperatures. Heating supplies the energy that breaks the existing bonds, and the symbol \(\Delta\) written over the arrow is the chemist’s shorthand for ‘heat this’.
Copper carbonate is a green powder. Heat it and it turns black, while a gas escapes.
copper carbonate → copper oxide + carbon dioxide
\[\mathrm{CuCO_3}\text{(s)} \xrightarrow{\ \Delta\ } \mathrm{CuO}\text{(s)} + \mathrm{CO_2}\text{(g)}\]
Lead nitrate gives a more dramatic result. On heating it leaves a yellow residue of lead monoxide, and gives off a brown gas, nitrogen dioxide, along with colourless oxygen.
lead nitrate → lead monoxide + nitrogen dioxide + oxygen
\[2\,\mathrm{Pb(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
Nitrogen dioxide is choking and poisonous, so this heating is done in a fume cupboard or a very well ventilated room.
2.2.4 Light
Light carries energy, and certain molecules absorb it. Once they have absorbed it they are activated — they hold more energy than before — and they then react rapidly. A reaction brought about by the action of light is called a photochemical reaction, and the process is also called photolysis.
The largest photochemical reaction on Earth happens in green leaves. Plants use carbon dioxide and water, in the presence of light, to build glucose.
\[6\,\mathrm{CO_2} + 12\,\mathrm{H_2O} \xrightarrow{\ \text{light}\ } \mathrm{C_6H_{12}O_6} + 6\,\mathrm{O_2} + 6\,\mathrm{H_2O}\]
Hydrogen and chlorine behave very differently. Kept in the dark they can be mixed with little effect, but in sunlight they combine to form hydrogen chloride.
\[\mathrm{H_2} + \mathrm{Cl_2} \xrightarrow{\ \text{sunlight}\ } 2\,\mathrm{HCl}\]
Chlorine water shows the same effect. Chlorine dissolved in water gives hydrochloric acid and hypochlorous acid, and sunlight then breaks the hypochlorous acid down, so bubbles of oxygen are seen rising through the liquid.
\[\mathrm{Cl_2} + \mathrm{H_2O} \rightarrow \mathrm{HCl} + \mathrm{HClO}\]
\[2\,\mathrm{HClO} \xrightarrow{\ \text{sunlight, UV radiation}\ } 2\,\mathrm{HCl} + \mathrm{O_2}\]
Two laboratory reagents have to be protected from light for exactly this reason, as Fig. 2.2-1 shows. Silver nitrate solution slowly decomposes to silver, nitrogen dioxide and oxygen, which is why the solution darkens and eventually turns black. Hydrogen peroxide decomposes to water and oxygen, a colourless, odourless gas.
\[2\,\mathrm{AgNO_3} \xrightarrow{\ \text{sunlight}\ } 2\,\mathrm{Ag} + 2\,\mathrm{NO_2} + \mathrm{O_2}\]
\[2\,\mathrm{H_2O_2} \xrightarrow{\ \text{sunlight}\ } 2\,\mathrm{H_2O} + \mathrm{O_2}\]
Diagram
Fig. 2.2-1 — Why Some Reagents Live In Brown Bottles Two reagent bottles side by side on a laboratory shelf with sunlight falling on both from above. Left bottle: clear glass, contents drawn darkening from colourless at the neck to grey-black at the base, labelled ‘clear glass — light passes through, silver nitrate decomposes’. Right bottle: brown glass, contents drawn clear and colourless, labelled ‘brown glass — light absorbed by the glass, solution unchanged’. Label both stoppers and add a shelf label reading ‘store away from direct sunlight’.
2.2.5 Electricity
Some compounds are broken up only when an electric current is passed through them. A reaction caused by, or accompanied by, the passage of an electric current is called an electrochemical reaction.
Water that has been slightly acidified conducts a current, and the current splits it into its two elements.
water → hydrogen + oxygen
\[2\,\mathrm{H_2O} \xrightarrow{\ \text{electricity}\ } 2\,\mathrm{H_2} + \mathrm{O_2}\]
Molten sodium chloride behaves in the same way, giving the metal at one electrode and the gas at the other.
\[2\,\mathrm{NaCl} \xrightarrow{\ \text{electricity}\ } 2\,\mathrm{Na} + \mathrm{Cl_2}\]
2.2.6 Pressure
Pressure forces particles closer together, so they collide more often. Some reactions take place only when the substances involved are put under pressure.
Mercuric chloride and potassium iodide are both white solids. Crushed together in a mortar, they give a scarlet substance, mercuric iodide. Mercury compounds are highly poisonous, so this is a demonstration to watch, never one to handle.
\[\mathrm{HgCl_2} + 2\,\mathrm{KI} \rightarrow \mathrm{HgI_2} + 2\,\mathrm{KCl}\]
Fig. 2.2-2 shows how sharply the colour announces the change. Industry uses pressure on a far larger scale. Nitrogen and hydrogen combine to form ammonia only above about 200 atmospheres, and even then iron must be present as a catalyst.
\[\mathrm{N_2} + 3\,\mathrm{H_2} \xrightleftharpoons[\ ]{\ \text{iron, above 200 atm}\ } 2\,\mathrm{NH_3}\]
Diagram
Fig. 2.2-2 — Two White Solids Give A Scarlet Product A porcelain mortar seen from slightly above, with a pestle resting in it. Left half of the mortar shows two small heaps of white powder side by side, labelled ‘mercuric chloride’ and ‘potassium iodide’. An arrow labelled ‘crushed together’ leads to a second mortar on the right in which the mixture is uniformly scarlet, labelled ‘mercuric iodide’. Add a warning tag reading ‘poisonous — demonstration only’.
2.2.7 Catalysts
Some reactions are so slow that they are of no practical use, and others run faster than we would like. Both can be managed with a catalyst — a substance that changes the rate at which a reaction occurs, speeding it up or slowing it down, without being used up itself.
Potassium chlorate is the standard illustration. On its own it decomposes only at about \(700^\circ\text{C}\), and even then oxygen comes off very slowly. Add manganese dioxide and decomposition begins at a much lower temperature, about \(300^\circ\text{C}\). When the reaction is over, the manganese dioxide is still there, unaffected.
potassium chlorate → potassium chloride + oxygen
\[2\,\mathrm{KClO_3} \xrightarrow{\ \mathrm{MnO_2}\ } 2\,\mathrm{KCl} + 3\,\mathrm{O_2}\]
A catalyst is chosen for a particular reaction, not for reactions in general. In industry, ammonia is oxidised to nitric oxide over platinum at about \(800^\circ\text{C}\).
\[4\,\mathrm{NH_3} + 5\,\mathrm{O_2} \xrightarrow{\ \mathrm{Pt},\ 800^\circ\text{C}\ } 4\,\mathrm{NO} + 6\,\mathrm{H_2O}\]
2.2.7.1 Positive Catalysts
When a catalyst accelerates a reaction, it is called a positive catalyst. Hydrogen peroxide decomposes slowly on its own; with manganese dioxide present, oxygen is released rapidly.
\[2\,\mathrm{H_2O_2} \xrightarrow{\ \mathrm{MnO_2}\ } 2\,\mathrm{H_2O} + \mathrm{O_2}\]
Living bodies depend on positive catalysts. Enzymes are proteins that act as catalysts in biochemical reactions, and digestion is their most familiar work. With enzymes, complex food material is broken down into simpler substances in two to three hours; without them the same job would take months.
2.2.7.2 Promoters
A catalyst can itself be helped. Substances that improve the efficiency of a catalyst, and so influence the rate of a reaction, are called promoters.
In the manufacture of ammonia, iron is the catalyst and molybdenum is the promoter. The molybdenum makes the iron work better than iron alone would.
\[\mathrm{N_2} + 3\,\mathrm{H_2} \xrightleftharpoons[\ ]{\ \text{iron} + \mathrm{Mo}\ } 2\,\mathrm{NH_3}\]
A promoter cannot do the job by itself. In the absence of a catalyst, a promoter alone cannot increase the rate of a reaction.
2.2.7.3 Negative Catalysts
A catalyst used to retard the rate of a reaction is called a negative catalyst, or an inhibitor. Slowing a reaction down is often just as valuable as speeding it up, because it keeps a substance usable for longer.
Phosphoric acid retards the decomposition of hydrogen peroxide. Alcohol slows down the oxidation of chloroform.
Heat, light, electricity and pressure are not the only forms of energy that can start a chemical change. Certain reactions proceed by absorbing sound energy; acetylene, for instance, breaks up into carbon and hydrogen.
\[\mathrm{C_2H_2} \xrightarrow{\ \text{sound energy}\ } 2\,\mathrm{C} + \mathrm{H_2}\]
2.3 How A Chemical Reaction Announces Itself
Bonds breaking and re-forming are far too small to watch. What you can watch are the effects, and certain effects are so typical that a chemist treats them as evidence. Seven of them are worth knowing well.
Remember the order of thinking. You observe the change first; only then do you say what it means.
2.3.1 Formation Of A New Substance
The most fundamental sign is the one the definition demands: a substance is there at the end that was not there at the start, and it behaves differently.
Drop a small piece of sodium into water and it fizzes across the surface. The soft grey metal disappears, a gas comes off, and the water is left containing sodium hydroxide.
sodium + water → sodium hydroxide + hydrogen
\[2\,\mathrm{Na} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH} + \mathrm{H_2}\]
Neither product resembles what went in. Sodium hydroxide is a soapy, corrosive solid dissolved in the water, and hydrogen is a light, colourless gas.
2.3.2 Change Of Colour
A colour is a property, and new substances bring new properties with them. A change of colour is therefore worth noticing carefully.
Drop a few clean iron filings into blue copper sulphate solution. Over a few minutes the blue fades and the solution turns light green, while a reddish brown deposit collects at the bottom. The green colour belongs to ferrous sulphate, and the deposit is copper. Fig. 2.3-1 records the two stages.
iron + copper sulphate → ferrous sulphate + copper
\[\mathrm{Fe} + \mathrm{CuSO_4} \rightarrow \mathrm{FeSO_4} + \mathrm{Cu}\]
Diagram
Fig. 2.3-1 — Blue Fades To Green Two identical beakers side by side on a bench. Beaker 1, labelled ‘at the start’: deep blue solution labelled ‘copper sulphate solution’, with a small heap of grey iron filings just being added, labelled ‘iron filings’. Beaker 2, labelled ‘after a few minutes’: pale green solution labelled ‘ferrous sulphate solution’, with a reddish brown deposit settled on the bottom, labelled ‘copper’. Draw an arrow from beaker 1 to beaker 2.
Copper itself can be made to change colour. When concentrated sulphuric acid is added to a piece of copper, the brownish pink metal is attacked and a blue solution of copper sulphate is formed, along with water and sulphur dioxide gas.
\[\mathrm{Cu}\text{(s)} + 2\,\mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{CuSO_4}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
Sulphur dioxide is choking and harmful, and concentrated sulphuric acid burns skin on contact. This reaction is carried out in a fume cupboard, with the acid handled by tongs and a dropper, never poured freely.
2.3.3 Evolution Of A Gas
In a great many reactions one of the products is a gas, and you see it leave as a stream of bubbles.
Zinc granules dropped into dilute sulphuric acid begin to fizz at once. The bubbles are hydrogen, and the brisk bubbling itself has a name.
zinc + dilute sulphuric acid → zinc sulphate + hydrogen
\[\mathrm{Zn} + \mathrm{H_2SO_4} \rightarrow \mathrm{ZnSO_4} + \mathrm{H_2}\!\uparrow\]
Marble chips do the same in dilute hydrochloric acid, but the gas this time is carbon dioxide.
\[\mathrm{CaCO_3}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{CaCl_2}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{CO_2}\text{(g)}\!\uparrow\]
2.3.4 Formation Of A Precipitate
Sometimes the new substance will not dissolve in the liquid it is formed in. It appears as a cloud of fine solid that slowly settles.
You have already seen one of these. When silver nitrate solution and sodium chloride solution are mixed in section 2.2.2, the white solid that appears is a precipitate of silver chloride.
Colour makes precipitates useful. Add sodium hydroxide solution to ferrous sulphate solution and a dirty green solid appears, which is ferrous hydroxide.
ferrous sulphate + sodium hydroxide → ferrous hydroxide + sodium sulphate
\[\mathrm{FeSO_4}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_2}\!\downarrow + \mathrm{Na_2SO_4}\text{(aq)}\]
The downward arrow after a formula is the chemist’s way of saying that this product leaves the solution as a solid.
2.3.5 Change In Temperature
Put a thermometer in the mixture and it will very often move. A reaction that releases heat and warms its surroundings is exothermic; one that absorbs heat and cools its surroundings is endothermic.
Add water to quicklime and the mixture becomes so hot that the water boils and hisses. The product is slaked lime, and the change is strongly exothermic.
\[\mathrm{CaO}\text{(s)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{Ca(OH)_2}\text{(aq)} + \text{heat}\]
Section 2.9 takes up the energy side of reactions properly. Here it is enough to treat a rise or a fall in temperature as one more sign that a chemical change is under way.
2.3.6 Change Of State
A reaction can begin with gases and end with a solid, or begin with solids and end with a liquid. Since state is a property, a change of state on reacting is a genuine clue.
Ammonia gas and hydrogen chloride gas meet and form dense white fumes that settle as solid ammonium chloride.
\[\mathrm{NH_3}\text{(g)} + \mathrm{HCl}\text{(g)} \rightleftharpoons \mathrm{NH_4Cl}\text{(s)}\]
The reverse happens with carbon and sulphur. Both are solids, yet on heating together they give carbon disulphide, which is a liquid.
\[\mathrm{C}\text{(s)} + 2\,\mathrm{S}\text{(s)} \xrightarrow{\ \Delta\ } \mathrm{CS_2}\text{(l)}\]
2.3.7 Change In Odour
Your nose can be evidence too, provided it is used safely. A gas is never sniffed directly; the vapour is wafted gently towards the face with the hand.
Add hydrochloric acid to iron sulphide and the smell of rotten eggs fills the room. That smell belongs to hydrogen sulphide.
iron sulphide + hydrochloric acid → iron chloride + hydrogen sulphide
\[\mathrm{FeS} + 2\,\mathrm{HCl} \rightarrow \mathrm{FeCl_2} + \mathrm{H_2S}\]
2.4 Types Of Chemical Reactions
You now know what a chemical reaction is and how to spot one. The next question is whether the thousands of known reactions fall into any kind of order — and they do.
Look again at what a reaction actually does. One set of chemical substances is transformed into another; the atoms of the reactant molecules are rearranged to give product molecules with different properties. No atoms are added and none are thrown away. Only the pattern in which they are joined is different.
That being so, there are only so many patterns the rearrangement can follow. Atoms can be brought together into fewer, larger units. They can be pulled apart into more, smaller units. One element can push another out of a compound. Or two compounds can swap partners.
Those four patterns are the four types into which chemical reactions are classified.
- Direct combination, also called synthesis
- Decomposition
- Displacement
- Double decomposition
Table 2.1 sets them side by side. Read it now as a map of the sections that follow, and come back to it once you have met all four.
Notice that the first two are opposites of each other, and so, in a sense, are the last two: one element changing places, or two changing places at once. Each of the four now gets a section of its own.
2.5 Direct Combination Or Synthesis
The simplest thing atoms can do is join up. When two or more substances combine to form a single substance, the change is called a combination reaction, or synthesis.
\[\mathrm{A} + \mathrm{B} \rightarrow \mathrm{AB}\]
Two points follow from that general form, and both are worth fixing in mind. First, only one substance is formed; if two products appear, the reaction is not a combination. Second, the reactants may be elements or compounds, but the product is always a compound.
Combination reactions may need heat, light or electricity to get them going. They come in three kinds, depending on what is being combined.
2.5.1 Two Elements Combine To Form A Compound
Heat a mixture of iron filings and sulphur powder in a test tube. A dull grey solid forms which is neither iron nor sulphur: it is iron(II) sulphide.
iron + sulphur → iron(II) sulphide
\[\mathrm{Fe}\text{(s)} + \mathrm{S}\text{(s)} \xrightarrow{\ \text{heat}\ } \mathrm{FeS}\text{(s)}\]
The heat is needed only to start this reaction. Once it has begun, it keeps itself going, because it gives out heat of its own — it is exothermic, in the sense you met in section 2.3.5.
Many salts can be made this way. Iron(III) chloride, a brown compound, is prepared by passing dry chlorine over heated iron.
\[2\,\mathrm{Fe}\text{(s)} + 3\,\mathrm{Cl_2}\text{(g)} \rightarrow 2\,\mathrm{FeCl_3}\text{(s)}\]
Nitrogen is only moderately reactive, so it needs particular conditions before it will combine. Under those conditions, metals and non-metals form nitrides with it. Burning magnesium takes nitrogen straight from the air.
\[3\,\mathrm{Mg} + \mathrm{N_2} \rightarrow \mathrm{Mg_3N_2}\]
Hydrogen combines with nitrogen only under pressure above 200 atmospheres, at a temperature of \(450^\circ\text{C}\), with iron present as a catalyst. The product is ammonia, and you met these conditions in section 2.2.6.
\[\mathrm{N_2} + 3\,\mathrm{H_2} \rightleftharpoons 2\,\mathrm{NH_3}\]
Oxygen is the great combiner. Metals and non-metals alike unite with it under suitable conditions to give their oxides, and Table 2.2 collects the common ones.
Table 2.2 — Combination of elements with oxygen
| Element | Reaction with oxygen | Oxide formed |
|---|---|---|
| Sodium | \(4\,\mathrm{Na} + \mathrm{O_2} \rightarrow 2\,\mathrm{Na_2O}\text{(s)}\) | Sodium oxide |
| Iron | \(4\,\mathrm{Fe} + 3\,\mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{Fe_2O_3}\text{(s)}\) | Iron(III) oxide |
| Iron | \(3\,\mathrm{Fe} + 2\,\mathrm{O_2} \xrightarrow{\ \Delta\ } \mathrm{Fe_3O_4}\text{(s)}\) | Magnetic oxide of iron |
| Mercury | \(2\,\mathrm{Hg} + \mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{HgO}\text{(s)}\) | Mercuric oxide |
| Nitrogen | \(\mathrm{N_2} + \mathrm{O_2} \xrightarrow{\ 3000^\circ\text{C}\ } 2\,\mathrm{NO}\text{(g)}\) | Nitric oxide |
| Phosphorus | \(4\,\mathrm{P} + 5\,\mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{P_2O_5}\text{(s)}\) | Phosphorus pentoxide |
| Carbon | \(\mathrm{C}\text{(s)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ \Delta\ } \mathrm{CO_2}\text{(g)}\) | Carbon dioxide |
The most spectacular of these is magnesium. Held in a pair of tongs over a flame, a strip of magnesium ribbon catches and burns with a dazzling white light, leaving a white powder of magnesium oxide, as in Fig. 2.5-1. Never look straight at the flame; the light is bright enough to hurt the eyes.
\[2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO}\]
Diagram
Fig. 2.5-1 — Magnesium Ribbon Burns In Air A hand holding a pair of tongs, gripping the end of a short strip of magnesium ribbon; the far end of the ribbon is alight with a bright white flame drawn as a starburst. A Bunsen burner stands below and to the left, its flame just reaching the ribbon. Beneath the ribbon sits a watch glass collecting a small heap of white powder. Label the tongs, the magnesium ribbon, the burner, and the white powder as magnesium oxide.
A second laboratory demonstration uses lead sulphide. Black lead sulphide heated in a test tube takes up oxygen and turns into white lead sulphate.
\[\mathrm{PbS} + 2\,\mathrm{O_2} \rightarrow \mathrm{PbSO_4}\]
2.5.2 An Element And A Compound Combine
Neither reactant has to be an element. A compound can combine with an element, provided the result is still a single product.
Carbon monoxide is a compound. It burns in oxygen, an element, and the only product is carbon dioxide.
carbon monoxide + oxygen → carbon dioxide
\[2\,\mathrm{CO}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ \text{heat}\ } 2\,\mathrm{CO_2}\text{(g)}\]
Sulphur dioxide behaves in the same way. Under suitable conditions it combines with oxygen to give sulphur trioxide.
\[2\,\mathrm{SO_2}\text{(g)} + \mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{SO_3}\text{(g)}\]
2.5.3 Two Or More Compounds Combine
Two compounds can also combine, and the product is again a single compound. Ammonia and hydrogen chloride, both compounds, give ammonium chloride — the reaction you saw as a change of state in section 2.3.6.
\[\mathrm{NH_3}\text{(g)} + \mathrm{HCl}\text{(g)} \rightarrow \mathrm{NH_4Cl}\text{(s)}\]
Soluble metal oxides combine with water to form their hydroxides.
\[\mathrm{K_2O} + \mathrm{H_2O} \rightarrow 2\,\mathrm{KOH}\]
\[\mathrm{Na_2O} + \mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH}\]
Potassium hydroxide is caustic potash and sodium hydroxide is caustic soda, and both are dangerously corrosive to skin.
Calcium oxide, or quicklime, reacts with water so vigorously that a great deal of heat is liberated. The product is calcium hydroxide, better known as slaked lime.
\[\mathrm{CaO}\text{(s)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{Ca(OH)_2}\text{(aq)}\]
Slaked lime is used for white washing, and there is more chemistry in a whitewashed wall than you might expect. Once the solution has been applied, the calcium hydroxide reacts slowly with carbon dioxide from the atmosphere, and a thin layer of calcium carbonate builds up on the wall.
\[\mathrm{Ca(OH)_2}\text{(aq)} + \mathrm{CO_2}\text{(g)} \rightarrow \mathrm{CaCO_3}\text{(s)} + \mathrm{H_2O}\text{(l)}\]
It is this layer of calcium carbonate that gives a whitewashed wall its shiny finish, and it forms over a day or two rather than at once.
Photosynthesis is also counted as a combination reaction, since carbon dioxide and water are used by plants to make a single food substance, glucose. Its equation is given in section 2.2.4.
2.5.4 Oxides, Bases, Alkalis And Acidic Anhydrides
The oxides in the last three subsections divide neatly into two families, and the division follows the elements they came from.
Metallic oxides such as sodium oxide, zinc oxide and calcium oxide are called bases. Those bases that dissolve in water are called alkalis. Sodium hydroxide, potassium hydroxide, ammonium hydroxide and calcium hydroxide are all alkalis.
Non-metallic oxides do the opposite. Those that dissolve in water form acids.
\[\mathrm{SO_2} + \mathrm{H_2O} \rightarrow \mathrm{H_2SO_3}\]
\[\mathrm{P_2O_5} + 3\,\mathrm{H_2O} \rightarrow 2\,\mathrm{H_3PO_4}\]
The first gives sulphurous acid and the second phosphoric acid. A non-metallic oxide that reacts with water to form an acid is called an acidic anhydride. Sulphur dioxide is the acidic anhydride of sulphurous acid, and phosphorus pentoxide is the acidic anhydride of phosphoric acid.
2.6 Decomposition
A decomposition reaction is combination run backwards. A single compound is pulled apart into two or more simpler substances, and those substances do not join up again on their own.
\[\mathrm{AB} \rightarrow \mathrm{C} + \mathrm{D} + \dots\]
Usually a single reactant decomposes with no second substance present at all. What it needs is energy, and very often heat alone is enough.
Decomposition may be brought about by heat, by light, or by the passage of an electric current. That list should look familiar. Heat, light and electricity are different forms of energy, and energy is what breaks the bonds inside the reactant molecules.
Decomposition is sorted by what the compound breaks into.
2.6.1 A Compound Breaks Into Elements
Sometimes a compound falls apart all the way, leaving nothing but elements.
Mercuric oxide is an orange-red powder. Heat it and two elements appear: liquid mercury and oxygen gas. The symbol \(\Delta\) written over the arrow signifies that heat has caused the reaction.
mercuric oxide → mercury + oxygen
\[2\,\mathrm{HgO}\text{(s)} \xrightarrow{\ \Delta\ } 2\,\mathrm{Hg}\text{(l)} + \mathrm{O_2}\text{(g)}\]
Mercury vapour is poisonous, so this decomposition is always carried out in a fume cupboard.
Light does the same work on the silver halides. Silver chloride is white, but leave it in sunlight and it darkens as grey silver metal is left behind, while greenish yellow chlorine escapes.
\[2\,\mathrm{AgCl}\text{(s)} \xrightarrow{\ \text{sunlight}\ } 2\,\mathrm{Ag}\text{(s)} + \mathrm{Cl_2}\text{(g)}\]
Silver bromide behaves in exactly the same way. The light yellow solid gives grey silver and brown bromine.
\[2\,\mathrm{AgBr}\text{(s)} \xrightarrow{\ \text{sunlight}\ } 2\,\mathrm{Ag}\text{(s)} + \mathrm{Br_2}\text{(g)}\]
An electric current does the same to acidulated water, splitting it into hydrogen and oxygen, as set out in section 2.2.5.
2.6.2 A Compound Breaks Into Elements And Compounds
The break need not go all the way. A compound may split into a simpler compound and an element.
Potassium nitrate, on heating, gives potassium nitrite, which is a compound, and oxygen, which is an element.
potassium nitrate → potassium nitrite + oxygen
\[2\,\mathrm{KNO_3}\text{(s)} \xrightarrow{\ \text{heat}\ } 2\,\mathrm{KNO_2}\text{(s)} + \mathrm{O_2}\text{(g)}\]
2.6.3 A Compound Breaks Into New Compounds
A compound may also break into two or more new compounds, with no free element at all.
Calcium carbonate is the substance of limestone and marble. Heated strongly, at about \(1000^\circ\text{C}\), it gives calcium oxide and carbon dioxide.
calcium carbonate → calcium oxide + carbon dioxide
\[\mathrm{CaCO_3}\text{(s)} \xrightarrow{\ 1000^\circ\text{C}\ } \mathrm{CaO}\text{(s)} + \mathrm{CO_2}\text{(g)}\]
Ferrous sulphate gives three products at once, and the changes are easy to follow with the eye. Ferrous sulphate crystals are really ferrous sulphate heptahydrate, \(\mathrm{FeSO_4\!\cdot\!7H_2O}\), and they are green. On heating in a test tube they first lose their water and turn white, and then a brown solid of ferric oxide is formed.
\[2\,\mathrm{FeSO_4}\text{(s)} \xrightarrow{\ \text{heat}\ } \mathrm{Fe_2O_3}\text{(s)} + \mathrm{SO_2}\text{(g)} + \mathrm{SO_3}\text{(g)}\]
The gas coming off is a mixture of sulphur dioxide and sulphur trioxide, with the suffocating, pungent smell of burning sulphur. Sulphur dioxide turns acidified potassium dichromate paper green, and that colour change is how the gas is confirmed. Fig. 2.6-1 shows the arrangement.
Diagram
Fig. 2.6-1 — Thermal Decomposition Of Ferrous Sulphate A boiling tube held at about forty-five degrees in a clamp on a retort stand, heated from below by a small lamp or burner. Green crystals sit at the closed end, labelled ‘ferrous sulphate’. Droplets of water are drawn condensed on the cooler upper wall of the tube, labelled ‘water droplets’. Wavy lines rising from the mouth are labelled ‘SO₃ and SO₂ gas’. A hand holds a strip of wet blue litmus paper at the mouth of the tube, labelled ‘wet blue litmus paper’. Label the clamp and the lamp.
Two more decompositions are worth carrying out. Zinc carbonate, heated in a test tube fitted with a cork and a bent delivery tube, gives off carbon dioxide, which turns lime water milky. The residue of zinc oxide is yellow while hot and white on cooling.
\[\mathrm{ZnCO_3} \xrightarrow{\ \Delta\ } \mathrm{ZnO} + \mathrm{CO_2}\!\uparrow\]
Orange ammonium dichromate is the most theatrical of them all. On heating it swells up, decomposes with flashes of light, gives off nitrogen and water vapour, and leaves a green solid, chromium oxide.
\[(\mathrm{NH_4})_2\mathrm{Cr_2O_7} \xrightarrow{\ \Delta\ } \mathrm{Cr_2O_3} + 4\,\mathrm{H_2O}\!\uparrow + \mathrm{N_2}\!\uparrow\]
Because glowing solid is thrown out of the tube, this is done in a fume cupboard, on a heat-resistant mat, with nothing flammable nearby.
2.6.4 Thermal Decomposition Of Metal Compounds
The decompositions in the last three subsections were all brought about by heat, and none of the products joined up again when the tube cooled. That combination has a name: thermal decomposition.
Not every metal compound decomposes equally readily, and the pattern is not random. It follows the reactivity of the metal. Compounds of reactive metals are more stable and do not decompose easily, while the compounds of less reactive metals decompose readily. Keep that one sentence in mind and the four tables below stop being lists to memorise and become a single idea applied four times.
2.6.4.1 Metal Hydroxides
Table 2.3 — Action of heat on metal hydroxides
| Metal | What the hydroxide does on heating | Example |
|---|---|---|
| K, Na | Stable to heat; they do not decompose | — |
| Ca, Mg, Al, Zn, Fe, Pb, Cu | Decompose to form the metal oxide and water vapour | \(\mathrm{Ca(OH)_2} \rightarrow \mathrm{CaO} + \mathrm{H_2O}\); \(\mathrm{Zn(OH)_2} \rightarrow \mathrm{ZnO} + \mathrm{H_2O}\); \(\mathrm{Pb(OH)_2} \rightarrow \mathrm{PbO} + \mathrm{H_2O}\); \(\mathrm{Cu(OH)_2} \rightarrow \mathrm{CuO} + \mathrm{H_2O}\) |
| Hg, Ag | Decompose to yield the metal, oxygen and water vapour | \(4\,\mathrm{AgOH} \rightarrow 4\,\mathrm{Ag} + \mathrm{O_2} + 2\,\mathrm{H_2O}\) |
2.6.4.2 Metal Carbonates
Table 2.4 — Action of heat on metal carbonates
| Metal | What the carbonate does on heating | Example |
|---|---|---|
| K, Na | Stable to heat, and soluble in water | — |
| Ca, Mg, Al, Zn, Fe, Pb, Cu | Decompose, with decreasing vigour down the list, to form the metal oxide and carbon dioxide | \(\mathrm{MgCO_3} \rightarrow \mathrm{MgO} + \mathrm{CO_2}\); \(\mathrm{ZnCO_3} \rightarrow \mathrm{ZnO} + \mathrm{CO_2}\); \(\mathrm{CuCO_3} \rightarrow \mathrm{CuO} + \mathrm{CO_2}\) |
| Hg, Ag | Decompose to form the metal, oxygen and carbon dioxide | \(2\,\mathrm{Ag_2CO_3} \rightarrow 4\,\mathrm{Ag} + \mathrm{O_2} + 2\,\mathrm{CO_2}\) |
2.6.4.3 Metal Bicarbonates
Metal bicarbonates, also called metal hydrogen carbonates, are simpler: they all behave alike. Each decomposes to give the metal carbonate, water vapour and carbon dioxide.
Table 2.5 — Action of heat on metal bicarbonates
| Bicarbonate | Equation |
|---|---|
| Sodium hydrogen carbonate | \(2\,\mathrm{NaHCO_3} \xrightarrow{\ \Delta\ } \mathrm{Na_2CO_3} + \mathrm{H_2O} + \mathrm{CO_2}\) |
| Calcium hydrogen carbonate | \(\mathrm{Ca(HCO_3)_2} \xrightarrow{\ \Delta\ } \mathrm{CaCO_3} + \mathrm{H_2O} + \mathrm{CO_2}\) |
| Magnesium bicarbonate | \(\mathrm{Mg(HCO_3)_2} \xrightarrow{\ \Delta\ } \mathrm{MgCO_3} + \mathrm{H_2O} + \mathrm{CO_2}\) |
2.6.4.4 Metal Nitrates
Every metal nitrate decomposes on heating, but what is left behind depends on the metal, as Table 2.6 shows.
Table 2.6 — Action of heat on metal nitrates
| Metal | What the nitrate does on heating | Example |
|---|---|---|
| K, Na | Melt, then decompose to give the metal nitrite and oxygen | \(2\,\mathrm{KNO_3} \rightarrow 2\,\mathrm{KNO_2} + \mathrm{O_2}\) |
| Ca, Mg, Al, Zn, Fe, Pb, Cu | Decompose to form the metal oxide, nitrogen dioxide and oxygen | \(2\,\mathrm{Ca(NO_3)_2} \rightarrow 2\,\mathrm{CaO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\); \(2\,\mathrm{Zn(NO_3)_2} \rightarrow 2\,\mathrm{ZnO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\); \(2\,\mathrm{Pb(NO_3)_2} \rightarrow 2\,\mathrm{PbO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\); \(2\,\mathrm{Cu(NO_3)_2} \rightarrow 2\,\mathrm{CuO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\) |
| Hg, Ag | Decompose to form the metal, nitrogen dioxide and oxygen | \(2\,\mathrm{AgNO_3} \xrightarrow{\ \Delta\ } 2\,\mathrm{Ag} + 2\,\mathrm{NO_2} + \mathrm{O_2}\) |
Decomposition Inside You
Digestion is decomposition, carried out at body temperature with enzymes doing the work of heat. The starch in the food you eat is broken down into glucose and sugar, proteins are broken down to amino acids, and fats and oils are broken down to fatty acids, which respiration finally oxidises to carbon dioxide and water. FC 2.1 traces the path of the starch in a chapati.
Flowchart
FC 2.1 — Starch To Energy In The Body Starch in food → broken down by enzymes → glucose → oxidised by respiration → carbon dioxide + water + energy
2.6.5 Dehydration Of A Hydrated Salt
Some crystals hold water inside the crystal itself, called water of crystallisation. Heating drives that water out, in what is called a dehydration reaction.
Blue crystals of hydrated copper(II) sulphate, heated in a test tube, turn into a white anhydrous salt while steam escapes.
\[\mathrm{CuSO_4\!\cdot\!5H_2O}\text{(s)} \rightarrow \mathrm{CuSO_4}\text{(s)} + 5\,\mathrm{H_2O}\text{(g)}\]
The white salt is not ruined. Add a few drops of water to it and the blue colour returns at once, which is exactly why chemists use it as a test for the presence of water.
\[\mathrm{CuSO_4}\text{(s)} + 5\,\mathrm{H_2O}\text{(l)} \rightarrow \mathrm{CuSO_4\!\cdot\!5H_2O}\text{(s)}\]
2.6.6 Reversible Reactions
Almost every reaction so far has run one way. Reactants went in, products came out, and it was clear which was which.
The copper sulphate you have just met does not behave like that. Heat drives the water out, water puts it back, and the direction depends entirely on the conditions. Reactions whose direction can be reversed by changing the conditions are called reversible reactions, and they are written with a double arrow.
\[\mathrm{CuSO_4\!\cdot\!5H_2O}\text{(s)} \rightleftharpoons \mathrm{CuSO_4}\text{(s)} + 5\,\mathrm{H_2O}\text{(g)}\]
Iron and steam behave the same way. Pass steam over red hot iron and the metal is partially converted into its magnetic oxide, releasing hydrogen; pass hydrogen over the heated oxide and it partially changes back to iron along with steam.
\[3\,\mathrm{Fe} + 4\,\mathrm{H_2O} \rightleftharpoons \mathrm{Fe_3O_4} + 4\,\mathrm{H_2}\]
The word ‘partially’ is the important one. In either case the vessel ends up holding a mixture of reactants and products, never one alone.
2.6.7 Thermal Dissociation
A reversible decomposition can be brought about by heat alone, and then it is called thermal dissociation.
Heat solid ammonium chloride in a test tube. Two colourless gases are produced, ammonia and hydrogen chloride. As they rise into the cooler upper part of the tube they recombine, and ammonium chloride appears again as a white sublimate on the cool glass.
\[\mathrm{NH_4Cl}\text{(s)} \rightleftharpoons \mathrm{NH_3}\text{(g)}\!\uparrow + \mathrm{HCl}\text{(g)}\!\uparrow\]
Nitrogen tetraoxide gives the same behaviour with a colour to watch. On heating, the colourless gas changes to nitrogen dioxide, a reddish brown gas; on cooling, the brown fades as the original compound re-forms.
\[\mathrm{N_2O_4} \xrightleftharpoons[\ \text{cool}\ ]{\ \text{heat}\ } 2\,\mathrm{NO_2}\]
2.7 Displacement
Elements are not equally eager to be in compounds. Put a keen element next to a compound of a less keen one, and the keen element will take its place.
That is a displacement reaction: a chemical change in which a more active element displaces a less active element from its salt solution.
\[\mathrm{AB} + \mathrm{C} \rightarrow \mathrm{AC} + \mathrm{B}\]
Here the more active element C has taken the place of the less active element B, and B has been turned out of the compound as a free element.
2.7.1 Displacement Of Hydrogen From Water
Water is a compound of hydrogen, so a metal active enough can displace hydrogen out of it. Metals that are more reactive than hydrogen do exactly that.
Potassium, sodium and calcium can displace hydrogen even from cold water.
\[2\,\mathrm{K} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{KOH} + \mathrm{H_2}\]
\[\mathrm{Ca} + 2\,\mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2} + \mathrm{H_2}\]
Sodium behaves in the same way, as you saw in section 2.3.1. These metals react so violently that only a very small piece is ever used, and always by a teacher, behind a safety screen.
Less reactive metals such as magnesium and zinc will not touch cold water. They need heat before they will liberate hydrogen from it.
\[\mathrm{Mg} + 2\,\mathrm{H_2O} \xrightarrow{\ \Delta\ } \mathrm{Mg(OH)_2} + \mathrm{H_2}\]
2.7.2 Displacement Of Hydrogen From Acids
An acid holds its hydrogen more loosely than water does, so the same metals displace hydrogen from acids more readily.
\[2\,\mathrm{K} + 2\,\mathrm{HCl} \rightarrow 2\,\mathrm{KCl} + \mathrm{H_2}\]
\[\mathrm{Zn} + 2\,\mathrm{HCl} \rightarrow \mathrm{ZnCl_2} + \mathrm{H_2}\]
Zinc and dilute sulphuric acid, which you met as an example of effervescence in section 2.3.3, is a displacement reaction of the same kind.
Magnesium shows it best. Drop a small piece of magnesium ribbon into dilute sulphuric acid in a test tube and brisk effervescence begins at once. Bring a burning splint to the mouth of the tube and the escaping gas burns with a squeaky pop, which is the test for hydrogen. Fig. 2.7-1 shows the arrangement.
magnesium + dilute sulphuric acid → magnesium sulphate + hydrogen
\[\mathrm{Mg} + \mathrm{H_2SO_4} \rightarrow \mathrm{MgSO_4} + \mathrm{H_2}\]
Diagram
Fig. 2.7-1 — Testing The Gas With A Burning Splint A test tube held in a clamp, about one-third full of a colourless liquid labelled ‘dilute sulphuric acid’, with a short strip of metal at the bottom labelled ‘magnesium ribbon’. Bubbles are drawn rising through the liquid, labelled ‘effervescence’. A hand holds a lighted wooden splint just at the mouth of the tube, with a small burst drawn at the splint tip labelled ‘gas burns with a pop — hydrogen’.
2.7.3 Displacement Of A Metal From Its Salt Solution
A metal can also be turned out of its own salt by a more active metal, and here the colours make everything visible.
Take copper sulphate solution in a beaker, add a few pieces of zinc, and stir with a glass rod. The blue colour gradually fades until the solution is almost colourless, while reddish brown particles of copper settle at the bottom. Zinc has displaced copper from copper sulphate.
\[\mathrm{CuSO_4}\text{(aq)} + \mathrm{Zn} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu}\!\downarrow\]
Now try it the other way round. Put a piece of copper into zinc sulphate solution and nothing at all happens. A less reactive element cannot displace a more reactive one.
\[\mathrm{ZnSO_4} + \mathrm{Cu} \rightarrow \text{no product}\]
That failed experiment teaches as much as the successful one. Displacement is a one-way street, and its direction is fixed by which element is the more active.
2.7.4 Displacement Among Non-Metals
Non-metals displace one another too, on exactly the same principle.
Pass chlorine gas through a solution of potassium iodide. The colourless solution turns yellow brown, because iodine has been set free. Chlorine is toxic, so this is done in a fume cupboard.
\[2\,\mathrm{KI} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{KCl} + \mathrm{I_2}\]
Put the three experiments of this section together and a pattern appears. Among the metals, zinc is more active than copper, and magnesium is more active than hydrogen. Among the non-metals, chlorine is more active than iodine.
2.7.5 The Activity Series Of Elements
By taking many such pairs and recording which displaces which, chemists have arranged the elements in order of activity. The result is the activity series, also called the reactivity series, given in Table 2.7.
Table 2.7 — Activity (reactivity) series of elements
| Metals, most active first | Non-metals, most active first | ||
|---|---|---|---|
| Potassium | K | Fluorine | F |
| Sodium | Na | Chlorine | Cl |
| Calcium | Ca | Bromine | Br |
| Magnesium | Mg | Iodine | I |
| Aluminium | Al | ||
| Zinc | Zn | ||
| Iron | Fe | ||
| Lead | Pb | ||
| [Hydrogen] | [H] | ||
| Copper | Cu | ||
| Mercury | Hg | ||
| Silver | Ag | ||
| Gold | Au | ||
| Platinum | Pt |
Hydrogen is set in brackets because it is not a metal. It is placed in the list as a marker: any metal above it displaces hydrogen from acids, and any metal below it does not.
Used properly, the series is a prediction tool rather than a list to be memorised. The more reactive element displaces the less reactive element from its salt solution — so once you know where two elements stand, you know which way the reaction will go, and whether it will go at all.
2.8 Double Decomposition
In a displacement reaction one element changes places. In the fourth and last type, two compounds change places with each other.
Double decomposition is a chemical change in which two compounds in a solution react to form two new compounds by a mutual exchange of radicals. It is also called double displacement.
\[\mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}\]
A radical here means the part of a compound that travels as a unit — the metal or the ammonium part on one side, the acidic part such as sulphate, chloride or nitrate on the other. In a double decomposition each compound keeps its own units but hands them to a new partner.
These reactions are of two kinds, and the rest of this section takes them in turn: precipitation reactions and neutralisation reactions.
2.8.1 Precipitation Reactions
When partners are exchanged, one of the new pairs may turn out to be insoluble. It cannot stay in solution, so it appears as a solid, and the change is called a precipitation reaction.
Three examples show the range of colours possible.
\[\mathrm{BaCl_2}\text{(aq)} + \mathrm{Na_2SO_4}\text{(aq)} \rightarrow \mathrm{BaSO_4}\!\downarrow + 2\,\mathrm{NaCl}\text{(aq)}\]
\[\mathrm{CuSO_4}\text{(aq)} + \mathrm{H_2S}\text{(g)} \rightarrow \mathrm{CuS}\!\downarrow + \mathrm{H_2SO_4}\text{(aq)}\]
\[\mathrm{FeCl_2}\text{(aq)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Fe(OH)_2}\!\downarrow + 2\,\mathrm{NaCl}\text{(aq)}\]
Barium sulphate is white, copper sulphide is black, and ferrous hydroxide is the same green solid you met in section 2.3.4. Notice the second one: a gas was bubbled into a solution, and a precipitate still formed.
Sodium Sulphate With Barium Chloride
Take an aqueous solution of barium chloride in a test tube and add a little aqueous sodium sulphate solution. A white precipitate of barium sulphate forms immediately, along with sodium chloride which stays in solution. The insoluble white solid, shown in Fig. 2.8-1, is the sign that a chemical change has occurred.
\[\mathrm{Na_2SO_4}\text{(aq)} + \mathrm{BaCl_2}\text{(aq)} \rightarrow \mathrm{BaSO_4}\!\downarrow + 2\,\mathrm{NaCl}\text{(aq)}\]
Diagram
Fig. 2.8-1 — A White Precipitate Of Barium Sulphate A hand tipping a small test tube of colourless liquid, labelled ‘sodium sulphate solution (Na₂SO₄)’, into a second test tube held upright in another hand. The lower tube contains a colourless liquid labelled ‘barium chloride solution (BaCl₂)’, with a dense white solid drawn settling at the bottom, labelled ‘barium sulphate (white ppt)’.
This reaction is put to work as a test. Any soluble sulphate, treated with barium chloride solution, gives the same white precipitate.
\[\mathrm{H_2SO_4}\text{(aq)} + \mathrm{BaCl_2}\text{(aq)} \rightarrow \mathrm{BaSO_4}\!\downarrow + 2\,\mathrm{HCl}\text{(aq)}\]
So a white precipitate on adding barium chloride tells you that sulphate ions, \(\mathrm{SO_4^{2-}}\), are present in the solution.
Silver Nitrate With Sodium Chloride
Sodium chloride solution and silver nitrate solution give a white curdy precipitate of silver chloride, the reaction first met in section 2.2.2. That precipitate dissolves in ammonium hydroxide solution, which distinguishes it from other white precipitates.
The reaction is the standard test for chloride ions, \(\mathrm{Cl^{-}}\). Any soluble chloride treated with silver nitrate gives the white curdy precipitate.
\[\mathrm{AgNO_3}\text{(aq)} + \mathrm{HCl}\text{(aq)} \rightarrow \mathrm{AgCl}\!\downarrow + \mathrm{HNO_3}\text{(aq)}\]
Lead Nitrate With Potassium Iodide
Take a few millilitres of lead nitrate solution in a test tube and add a few drops of potassium iodide solution from a dropper. A thick, bright yellow precipitate of lead(II) iodide forms the instant the solutions meet, as in Fig. 2.8-2.
\[\mathrm{Pb(NO_3)_2}\text{(aq)} + 2\,\mathrm{KI}\text{(aq)} \rightarrow \mathrm{PbI_2}\!\downarrow + 2\,\mathrm{KNO_3}\text{(aq)}\]
Diagram
Fig. 2.8-2 — A Yellow Precipitate Of Lead Iodide A hand squeezing a dropper over an upright test tube held in a second hand. The dropper is labelled ‘potassium iodide solution (KI)’. The tube contains a colourless liquid labelled ‘lead nitrate solution [Pb(NO₃)₂]’, with a thick bright yellow solid drawn gathering at the bottom, labelled ‘lead(II) iodide (yellow ppt)’.
Here the potassium and lead exchange their acidic radicals, exactly as the general form predicts. The yellow precipitate dissolves in water on heating and reappears on cooling.
Take the three experiments together and the point is striking. Two colourless solutions, both perfectly soluble in water, give a new compound that is coloured and insoluble.
2.8.2 Double Decomposition With Evolution Of A Gas
The new partner need not be an insoluble solid. Sometimes it is a gas, and it leaves the mixture just as decisively.
Add dilute sulphuric acid or dilute hydrochloric acid to a metal sulphide, and hydrogen sulphide is released with its offensive rotten-egg smell.
\[\mathrm{FeS}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{FeSO_4}\text{(aq)} + \mathrm{H_2S}\!\uparrow\]
\[\mathrm{ZnS}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + \mathrm{H_2S}\!\uparrow\]
Sodium carbonate does the same with carbon dioxide. Take a small amount of the solid in a test tube, add a few drops of dilute hydrochloric acid, and a gas is evolved with brisk effervescence.
\[\mathrm{Na_2CO_3} + 2\,\mathrm{HCl} \rightarrow 2\,\mathrm{NaCl} + \mathrm{H_2CO_3}\]
The carbonic acid formed is unstable and breaks up at once into carbon dioxide and water, which is why bubbles appear rather than a third solution.
Pass that gas through lime water, as in Fig. 2.8-3, and the lime water turns milky because insoluble calcium carbonate is formed.
\[\mathrm{Ca(OH)_2} + \mathrm{CO_2} \rightarrow \mathrm{CaCO_3} + \mathrm{H_2O}\]
Diagram
Fig. 2.8-3 — Carbon Dioxide Turns Lime Water Milky A test tube held at a slant in one hand, containing a white solid labelled ‘sodium carbonate’ with bubbles rising through it, labelled ‘brisk effervescence’; a second hand adds liquid from a dropper labelled ‘dil. HCl’. A bent delivery tube leads from a stopper in the first tube down into a second test tube containing a cloudy white liquid, labelled ‘lime water turns milky’. Label the gas in the delivery tube ‘CO₂(g)’.
Both of these are double displacement reactions. Keep passing carbon dioxide, though, and the milkiness slowly disappears, because the insoluble carbonate is converted into soluble calcium hydrogen carbonate.
\[\mathrm{CaCO_3} + \mathrm{H_2O} + \mathrm{CO_2} \rightarrow \mathrm{Ca(HCO_3)_2}\]
Milkiness that appears and then clears in excess gas is the accepted test for carbon dioxide.
2.8.3 Neutralisation Reactions
The second kind of double decomposition involves an acid and a base, and its products are always the same two things.
Neutralisation is the reaction between an acid and a base that forms salt and water only. A salt, in chemistry, means any compound formed when the hydrogen of an acid is replaced by a metal — common table salt is only one member of a very large family.
\[\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}\]
Why should salt and water always be the products? Write the same reaction in its ionic form and the answer appears.
\[\mathrm{Na^{+}}\mathrm{OH^{-}} + \mathrm{H^{+}}\mathrm{Cl^{-}} \rightarrow \mathrm{Na^{+}}\mathrm{Cl^{-}} + \mathrm{H^{+}}\mathrm{OH^{-}}\]
The sodium and chloride ions appear unchanged on both sides, so they can be cancelled. The only real change is that the hydrogen ion from the acid joins the hydroxyl ion from the base to form un-ionised water.
\[\mathrm{H^{+}} + \mathrm{OH^{-}} \rightarrow \mathrm{H_2O}\]
In a neutralisation reaction a soluble base, that is an alkali, or an insoluble base reacts with an acid to form salt and water. Both cases matter.
Alkalis with acids:
\[\mathrm{KOH} + \mathrm{HNO_3} \rightarrow \mathrm{KNO_3} + \mathrm{H_2O}\]
\[2\,\mathrm{NaOH} + \mathrm{H_2SO_4} \rightarrow \mathrm{Na_2SO_4} + 2\,\mathrm{H_2O}\]
\[\mathrm{NH_4OH} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl} + \mathrm{H_2O}\]
Insoluble bases with acids:
\[\mathrm{PbO} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Pb(NO_3)_2} + \mathrm{H_2O}\]
\[\mathrm{CuO} + \mathrm{H_2SO_4} \rightarrow \mathrm{CuSO_4} + \mathrm{H_2O}\]
\[\mathrm{Pb(OH)_2} + 2\,\mathrm{HNO_3} \rightarrow \mathrm{Pb(NO_3)_2} + 2\,\mathrm{H_2O}\]
\[\mathrm{Cu(OH)_2} + \mathrm{H_2SO_4} \rightarrow \mathrm{CuSO_4} + 2\,\mathrm{H_2O}\]
Double decomposition reactions occurring in aqueous solutions of acids and bases produce a neutral salt, and are therefore counted as neutralisation reactions. Table 2.8 sets the two kinds of double decomposition side by side.
2.8.4 Neutralisation In Everyday Life
Neutralisation is not confined to a test tube. Wherever there is unwanted acid, something basic is used to cancel it.
A bee sting injects formic acid into the skin, and that is what causes the pain. Rubbing the spot with slaked lime or with baking soda, both of which are basic, neutralises the acid and relieves it.
The stomach makes hydrochloric acid to digest food. When the glands secrete too much of it, the pain is relieved by taking milk of magnesia, which is magnesium hydroxide, or a solution of sodium hydrogen carbonate. Both are basic. If instead there is too little acid, a suitable organic acid is taken in dilute form to make up for it.
Acid spilled on clothes can be neutralised with ammonia solution. Acid spilled on skin is different: wash the area with plenty of water first, then apply a paste of sodium hydrogen carbonate, a mild base, to neutralise whatever acid remains.
Farmers use the same chemistry on a field scale. Soil that has turned somewhat acidic is unfavourable for certain crops, so slaked lime is added to neutralise the excess acid.
2.9 Energy Changes In Chemical Reactions
Every chemical reaction moves energy about. There is always a difference between the chemical energy of the reactants and the chemical energy of the products, and that difference has to go somewhere or come from somewhere.
\[\text{energy change} = \text{total energy of the products} - \text{total energy of the reactants}\]
The energy can appear as heat, as light, as sound or as electricity. Heat is simply the commonest form, not the only one.
Why should there be a difference at all? Think back to what a reaction does. Breaking the bonds of the reactants absorbs energy, and forming the new bonds of the products releases it. These two amounts are almost never equal. If more energy is released than absorbed there is a surplus, and it leaves the reaction; if less is released than absorbed there is a deficit, and it must be supplied from outside.
Depending on which way the surplus runs, chemical reactions are of two types: exothermic and endothermic. Fig. 2.9-1 shows the difference at a glance.
Diagram
Fig. 2.9-1 — Where The Energy Sits Two energy-level panels side by side, each with a vertical axis labelled ‘chemical energy’ and no numbers. Left panel, headed ‘exothermic’: a horizontal bar high up labelled ‘reactants’, a horizontal bar low down labelled ‘products’, and a downward arrow between them labelled ‘energy given out to the surroundings’. Right panel, headed ‘endothermic’: a horizontal bar low down labelled ‘reactants’, a horizontal bar high up labelled ‘products’, and an upward arrow between them labelled ‘energy taken in from the surroundings’. Add a small thermometer beside each panel, rising in the left and falling in the right.
2.9.1 Exothermic Reactions
A reaction in which heat is given out is called an exothermic reaction, and it causes a rise in temperature. In such a reaction the total energy content of the reactants is more than that of the products, so the surplus escapes as heat.
Burning carbon is the plainest case. A great deal of heat is produced when it burns in oxygen.
\[\mathrm{C} + \mathrm{O_2} \rightarrow \mathrm{CO_2} + \text{heat}\]
The combustion of natural gas is an important source of energy for homes and industry. Natural gas is mainly methane, and its complete combustion produces carbon dioxide and water vapour.
\[\mathrm{CH_4} + 2\,\mathrm{O_2} \rightarrow \mathrm{CO_2} + 2\,\mathrm{H_2O}\]
Substances such as methane, which undergo combustion readily and give out a large amount of energy, are called fuels.
Water added to quicklime, which you met in section 2.3.5, gives out so much heat that the water boils and hisses. The process is called slaking of lime.
Hydrogen burning in oxygen releases heat as well as forming water.
\[2\,\mathrm{H_2} + \mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{H_2O} + \text{heat}\]
The manufacture of ammonia from its elements is exothermic too, and so is every neutralisation reaction. Mix an acid and an alkali in a test tube and the tube grows warm in your hand.
\[\mathrm{N_2} + 3\,\mathrm{H_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{NH_3} + \text{heat}\]
\[\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O} + \text{heat}\]
Your own body runs on an exothermic reaction. Food is broken down during digestion into simpler substances — the carbohydrates in rice, potatoes and bread are broken down to glucose — and the glucose is then combusted with oxygen in the cells to provide energy.
\[\mathrm{C_6H_{12}O_6} + 6\,\mathrm{O_2} \rightarrow 6\,\mathrm{CO_2} + 6\,\mathrm{H_2O} + \text{energy}\]
This reaction is exothermic, and it has a special name of its own: respiration.
Combustion, neutralisation, respiration, rusting, the disintegration of radioactive material, the explosion of dynamite, and the burning of coal, petrol, kerosene and cooking gas are all common exothermic reactions.
2.9.2 Endothermic Reactions
A reaction in which heat is absorbed is called an endothermic reaction, and it causes a fall in temperature. Here the total energy content of the products is more than that of the reactants, so energy has to be supplied.
Such reactions cannot be sustained without a supply of energy from an external source. Take the burner away and the reaction stops.
Three examples from earlier in this chapter are all endothermic. Carbon heated with sulphur at high temperature gives liquid carbon disulphide. Nitrogen and oxygen heated together to about \(3000^\circ\text{C}\) give nitric oxide. Calcium carbonate heated to about \(1000^\circ\text{C}\) decomposes into calcium oxide and carbon dioxide.
\[\mathrm{C} + 2\,\mathrm{S} \xrightarrow{\ \Delta\ } \mathrm{CS_2}\]
\[\mathrm{N_2} + \mathrm{O_2} \xrightarrow{\ 3000^\circ\text{C}\ } 2\,\mathrm{NO}\]
\[\mathrm{CaCO_3}\text{(s)} \xrightarrow{\ 1000^\circ\text{C}\ } \mathrm{CaO}\text{(s)} + \mathrm{CO_2}\text{(g)}\]
Notice how high those temperatures are. That is the deficit being paid for. The electrolysis of water and the extraction of iron from its ore are endothermic for the same reason.
2.9.3 Reactions Involving Light
Some chemical changes occur with the absorption of light energy, and these are the photochemical reactions of section 2.2.4.
Photosynthesis is the great example, and it is essential for life on Earth. It takes place in the green leaves of plants and requires energy from sunlight. The green pigment chlorophyll is essential to it, because chlorophyll is what traps the energy of the sun.
The decomposition of silver nitrate in sunlight is a photochemical reaction of the same kind, and its equation is given in section 2.2.4.
Light can also travel the other way. In combustion reactions light energy is evolved, as when magnesium ribbon burns.
\[2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO} + \text{light}\]
2.9.4 Reactions Involving Electricity
Reactions that occur with the absorption of electrical energy are called electrochemical reactions.
Fused potassium chloride, on passing a current through it, breaks up into charged particles, the ions of potassium and chloride.
\[\mathrm{KCl} \xrightarrow{\ \text{electric current}\ } \mathrm{K^{+}} + \mathrm{Cl^{-}}\]
Acidulated water breaks into hydrogen and oxygen when an electric current is passed through it, as you saw in section 2.2.5.
Electrochemical reactions can be either exothermic or endothermic, and a battery shows both faces. When a battery is used to power a device it releases energy, so the reactions inside it are exothermic. Charging that same battery, and the electrolysis of water, absorb electrical energy, and are therefore endothermic.
2.10 Summary
- A chemical change is a permanent change in which the composition of a substance is altered and one or more new substances, with different properties, are formed. It is not easily reversible.
- A chemical change is permanent, alters composition, is either exothermic or endothermic, and conserves mass: the total mass of the products equals the total mass of the reactants.
- A chemical bond is the force that holds the atoms of a molecule together.
- A chemical reaction is the breaking of the bonds of the reactants and the making of new bonds to form the products. Since the same atoms are simply rearranged, no mass is gained or lost.
- Reactions occur when particles collide, which requires close contact or a supply of energy. The conditions that bring about a chemical change are close contact, solution, heat, light, electricity, pressure and the presence of a catalyst.
- A reaction brought about by light is a photochemical reaction, or photolysis; one caused or accompanied by an electric current is an electrochemical reaction.
- A catalyst alters the rate of a reaction without taking part in it, and emerges unchanged. A positive catalyst accelerates a reaction, a negative catalyst or inhibitor retards it, and a promoter improves the efficiency of a catalyst but does nothing on its own.
- A chemical reaction is recognised by seven typical changes: formation of a new substance, change of colour, evolution of a gas, formation of a precipitate, change in temperature, change of state and change in odour.
- Effervescence is the formation of gas bubbles in a liquid during a reaction. A precipitate is an insoluble solid formed in a solution during a reaction.
- Chemical reactions are classified into four types: direct combination, decomposition, displacement and double decomposition.
- In a direct combination or synthesis reaction, two or more substances combine to form a single substance. The reactants may be elements or compounds, but the product is always a compound. \[\mathrm{Fe}\text{(s)} + \mathrm{S}\text{(s)} \xrightarrow{\ \text{heat}\ } \mathrm{FeS}\text{(s)}\]
- Combination occurs between two elements, between an element and a compound, or between two or more compounds. \[2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO}\]
- Metallic oxides are bases; bases that dissolve in water are alkalis. Non-metallic oxides that dissolve in water form acids, and are called acidic anhydrides.
- In a decomposition reaction a compound splits into two or more simpler substances, brought about by heat, light or an electric current. \[2\,\mathrm{HgO}\text{(s)} \xrightarrow{\ \Delta\ } 2\,\mathrm{Hg}\text{(l)} + \mathrm{O_2}\text{(g)}\]
- A compound may decompose into elements, into an element and a compound, or into new compounds.
- Thermal decomposition is decomposition brought about by heat with no recombination on cooling. How readily a metal compound decomposes depends on the reactivity of the metal: compounds of reactive metals are the more stable.
- Hydroxides, carbonates, bicarbonates and nitrates each follow this pattern. Compounds of potassium and sodium resist heat; those of the middle metals give the metal oxide; those of mercury and silver give the free metal. \[2\,\mathrm{Pb(NO_3)_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO} + 4\,\mathrm{NO_2} + \mathrm{O_2}\]
- A dehydration reaction is one in which a hydrated salt loses its water of crystallisation. Blue hydrated copper(II) sulphate turns white on heating, and the white salt turns blue again with water, which is the test for water.
- A reversible reaction can be made to run either way by changing the conditions, and is written with a double arrow. \[3\,\mathrm{Fe} + 4\,\mathrm{H_2O} \rightleftharpoons \mathrm{Fe_3O_4} + 4\,\mathrm{H_2}\]
- Thermal dissociation is a reversible decomposition brought about only by heat, the products recombining on cooling. \[\mathrm{NH_4Cl}\text{(s)} \rightleftharpoons \mathrm{NH_3}\text{(g)} + \mathrm{HCl}\text{(g)}\]
- In a displacement reaction a more active element displaces a less active element from its salt solution. \[\mathrm{CuSO_4}\text{(aq)} + \mathrm{Zn} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu}\!\downarrow\]
- Metals above hydrogen in the activity series displace hydrogen from water and from acids; a less reactive element cannot displace a more reactive one. Non-metals displace one another on the same principle.
- The activity series arranges elements in order of decreasing activity and is used to predict whether a displacement will occur.
- In a double decomposition, or double displacement, two compounds in solution exchange their radicals to form two new compounds. It is of two types, precipitation and neutralisation.
- A precipitation reaction forms an insoluble salt as one of the products, either on mixing two solutions or on bubbling a gas into a solution. \[\mathrm{Pb(NO_3)_2}\text{(aq)} + 2\,\mathrm{KI}\text{(aq)} \rightarrow \mathrm{PbI_2}\!\downarrow + 2\,\mathrm{KNO_3}\text{(aq)}\]
- Barium chloride solution gives a white precipitate with any soluble sulphate, and silver nitrate solution gives a white curdy precipitate with any soluble chloride. These are the tests for the sulphate and chloride radicals.
- Double decomposition may also release a gas, as when an acid is added to a metal sulphide or a carbonate. Carbon dioxide turns lime water milky, and the milkiness clears when excess gas is passed.
- Neutralisation is the reaction between an acid and a base that forms salt and water only. In ionic terms, \[\mathrm{H^{+}} + \mathrm{OH^{-}} \rightarrow \mathrm{H_2O}\]
- Both soluble bases (alkalis) and insoluble bases neutralise acids. Neutralisation is used to treat a bee sting, to relieve excess stomach acid, to deal with spilled acid and to correct acidic soil.
- Every chemical reaction involves an energy change, because breaking bonds absorbs energy and forming bonds releases it. The energy may appear as heat, light, sound or electricity.
- An exothermic reaction gives out heat and raises the temperature; the reactants hold more energy than the products. Combustion, neutralisation, respiration and rusting are exothermic. \[\mathrm{C_6H_{12}O_6} + 6\,\mathrm{O_2} \rightarrow 6\,\mathrm{CO_2} + 6\,\mathrm{H_2O} + \text{energy}\]
- An endothermic reaction absorbs heat and lowers the temperature; the products hold more energy than the reactants, and the reaction cannot continue without an external supply of energy. \[\mathrm{CaCO_3}\text{(s)} \xrightarrow{\ 1000^\circ\text{C}\ } \mathrm{CaO}\text{(s)} + \mathrm{CO_2}\text{(g)}\]
- Photochemical reactions absorb light energy, as in photosynthesis; combustion reactions release it, as when magnesium burns. Electrochemical reactions absorb electrical energy, and may themselves be exothermic or endothermic.
Flowchart
FC 2.2 — Deciding The Type Of A Reaction Look at the number of reactants and products → one reactant giving two or more products: decomposition (if heat causes it and there is no recombination on cooling, thermal decomposition; if the products recombine on cooling, thermal dissociation) → two or more reactants giving one product: direct combination → an element plus a compound, with the element taking the place of another: displacement → two compounds exchanging radicals: double decomposition (an insoluble product: precipitation; an acid with a base giving salt and water: neutralisation)
2.11 Key Terms
- Acidic anhydride — a non-metallic oxide that reacts with water to form an acid.
- Activity series — the arrangement of elements in order of their activity, used to predict which element will displace which.
- Alkali — a base that dissolves in water.
- Base — a metallic oxide or hydroxide, which reacts with an acid to form a salt and water.
- Catalyst — a substance that alters the rate of a reaction without taking part in it, and is recovered unchanged.
- Chemical bond — the force that holds the atoms of a molecule together.
- Chemical change — a permanent change in which the composition of a substance is altered and new substances with different properties are formed.
- Chemical energy — the energy stored in a substance, held as potential energy in the bonds between its atoms.
- Chemical reaction — the breaking of the bonds of the reactants and the making of new bonds to form the products.
- Combination reaction — a reaction in which two or more substances join to form a single substance; also called synthesis.
- Decomposition reaction — a reaction in which one compound splits into two or more simpler substances.
- Dehydration reaction — a reaction in which a hydrated salt loses its water of crystallisation.
- Displacement reaction — a reaction in which a more active element displaces a less active element from its salt solution.
- Double decomposition — a reaction in which two compounds in solution exchange their radicals to form two new compounds; also called double displacement.
- Effervescence — the formation of gas bubbles in a liquid during a reaction.
- Electrochemical reaction — a reaction caused or accompanied by the passage of an electric current.
- Endothermic reaction — a reaction that absorbs heat and causes a fall in temperature.
- Exothermic reaction — a reaction that gives out heat and causes a rise in temperature.
- Fuel — a substance that burns readily and gives out a large amount of energy.
- Negative catalyst — a catalyst that retards the rate of a reaction; also called an inhibitor.
- Neutralisation — the reaction between an acid and a base that forms salt and water only.
- Photochemical reaction — a reaction brought about by the action of light; also called photolysis.
- Positive catalyst — a catalyst that accelerates the rate of a reaction.
- Precipitate — an insoluble solid formed in a solution during a chemical reaction.
- Precipitation reaction — a reaction in which two compounds in aqueous solution form an insoluble salt as one of the products.
- Product — a substance formed in a chemical reaction.
- Promoter — a substance that improves the efficiency of a catalyst but cannot alter a rate on its own.
- Reactant — a starting substance that takes part in a chemical reaction.
- Respiration — the exothermic reaction in which glucose is oxidised in the cells to release energy.
- Reversible reaction — a reaction whose direction can be reversed by changing the conditions, written with a double arrow.
- Salt — a compound formed when the hydrogen of an acid is replaced by a metal.
- Synthesis — another name for a direct combination reaction, in which a substance is built up from simpler ones.
- Thermal decomposition — decomposition brought about by heat, with no recombination of the products on cooling.
- Thermal dissociation — a reversible decomposition brought about only by heat, the products recombining on cooling.
2.12 Points To Ponder
- A chemical change is described as permanent and not easily reversible, yet thermal dissociation reverses itself the moment the tube cools. Does the definition need repairing, or does the word ‘easily’ do enough work? What test would you use to decide whether a particular change deserves to be called permanent?
- A catalyst is said not to take part in the reaction, yet it may undergo temporary changes during it. If a substance is chemically altered in the middle of a reaction and restored by the end, in what sense has it stayed out of the reaction?
- The activity series was built up by recording which element displaces which. Suppose two elements sit next to each other in the list and neither displaces the other under ordinary conditions. How could a chemist decide their order, and how confident should anyone be about the placing?
- Every reaction type in this chapter has an opposite: combination against decomposition, and displacement against a displacement running the other way. Does that mean every reaction could in principle be reversed if the conditions were extreme enough, or is something else stopping it?
- Compounds of reactive metals are the hardest to decompose by heat, and the metals themselves are the hardest to obtain in the free state. Are these two facts one fact seen twice, and what does that suggest about where a metal’s ores are likely to be found?
- Photosynthesis stores energy from sunlight in glucose, and respiration releases it again. If both are chemical reactions with the same substances on opposite sides, why can a plant not simply run one and then the other and get something for nothing?
2.13 Objective Questions
Multiple Choice
- A reaction of the type \(\mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}\) is called: [1]
- direct combination
- decomposition
- displacement
- double decomposition
- Silver nitrate, on strong heating, leaves behind: [1]
- silver nitrite and oxygen
- silver, nitrogen dioxide and oxygen
- silver oxide and nitrogen dioxide only
- no residue at all
- The gas evolved when dilute hydrochloric acid is added to zinc sulphide is: [1]
- hydrogen
- sulphur dioxide
- hydrogen sulphide
- carbon dioxide
- A substance added to retard the rate of a chemical reaction is called: [1]
- a promoter
- an inhibitor
- a positive catalyst
- an enzyme
- Which of these carbonates does not decompose on heating? [1]
- Calcium carbonate
- Sodium carbonate
- Zinc carbonate
- Copper carbonate
- Which of the following is a displacement reaction? [1]
- \(\mathrm{NH_3} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl}\)
- \(\mathrm{Zn} + \mathrm{CuSO_4} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu}\)
- \(\mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}\)
- \(\mathrm{CaCO_3} \rightarrow \mathrm{CaO} + \mathrm{CO_2}\)
- Adding a few drops of water to a white anhydrous salt turns it blue. The salt is: [1]
- barium sulphate
- silver chloride
- copper(II) sulphate
- lead iodide
Assertion–Reason
- For each of the following, choose from these codes:
- 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.
Assertion (A): A whitewashed wall gains a shiny finish a day or two after the whitewash is applied. Reason (R): Calcium hydroxide reacts slowly with carbon dioxide in the air to form a thin layer of calcium carbonate. [1]
Assertion (A): Compounds of potassium are difficult to decompose by heat. Reason (R): Potassium stands at the top of the activity series. [1]
Assertion (A): Anhydrous copper(II) sulphate is used to detect the presence of water. Reason (R): It changes from blue to white when water is added to it. [1]
Assertion (A): Copper does not displace zinc from zinc sulphate solution. Reason (R): A less active element cannot displace a more active element from its salt solution. [1]
Match The Following
- Match each reaction in column A with its type in column B. [4]
| Column A | Column B |
|---|---|
| a. \(\mathrm{NH_4Cl} \rightleftharpoons \mathrm{NH_3} + \mathrm{HCl}\) | i. Precipitation |
| b. \(\mathrm{Pb(NO_3)_2} + 2\,\mathrm{KI} \rightarrow \mathrm{PbI_2} + 2\,\mathrm{KNO_3}\) | ii. Direct combination |
| c. \(2\,\mathrm{CO} + \mathrm{O_2} \rightarrow 2\,\mathrm{CO_2}\) | iii. Neutralisation |
| d. \(\mathrm{CuO} + \mathrm{H_2SO_4} \rightarrow \mathrm{CuSO_4} + \mathrm{H_2O}\) | iv. Thermal dissociation |
Fill In The Blanks
- Complete each statement with a suitable word. [4]
- A non-metallic oxide that reacts with water to form an acid is called an …………… .
- A reaction brought about by the action of light is called a …………… reaction.
- A base that dissolves in water is called an …………… .
- In a …………… reaction, heat is absorbed and the temperature falls.
True Or False
- State whether each statement is true or false. [4]
- A promoter can increase the rate of a reaction even when no catalyst is present.
- Every alkali is a base, but not every base is an alkali.
- In thermal dissociation the products recombine on cooling.
- Barium chloride solution is used to test for chloride ions.
Name The Following
- Name the following. [4]
- The brown gas produced when lead nitrate is heated.
- The metal used as a catalyst in the manufacture of ammonia.
- The white sublimate that collects on the cooler part of a test tube in which ammonium chloride is heated.
- The green solid left behind when ammonium dichromate is heated.
Give One Word Or Phrase
- Give one word or phrase for each of the following. [3]
- The formation of gas bubbles in a liquid during a reaction.
- The loss of water of crystallisation by a hydrated salt on heating.
- The force that holds the atoms of a molecule together.
Odd One Out
- Pick the odd one out in each set and give a reason for your choice. [2]
- Zinc carbonate, copper carbonate, sodium carbonate, magnesium carbonate.
- Respiration, rusting, combustion of methane, decomposition of limestone.
2.14 Check Your Concepts
Define a chemical change and state any two of its characteristics. [2]
State what is meant by a catalyst, and name one used in the laboratory preparation of oxygen from potassium chlorate. [2]
Give a reason for each of the following. [3]
- Hydrogen peroxide is stored in brown bottles.
- Molybdenum is added along with iron in the manufacture of ammonia.
- Quicklime is stored in sealed containers in a dry place.
Differentiate between a positive catalyst and a negative catalyst, stating clearly the basis of difference, and give one example of each. [3]
Differentiate between thermal decomposition and thermal dissociation, stating the basis of difference, and give one balanced equation for each. [4]
Why does the blue colour of copper sulphate solution fade when a few pieces of zinc are added to it, while nothing happens when a piece of copper is added to zinc sulphate solution? [3]
Explain, with one example each, what is meant by a precipitation reaction and a neutralisation reaction, and state what the two have in common. [4]
What would happen if the following were carried out? Give a reason in each case. [4]
- Carbon dioxide is passed through lime water for a long time.
- Green ferrous sulphate crystals are heated strongly in a dry test tube.
- Silver chloride is left on a watch glass in bright sunlight.
- A piece of magnesium ribbon is dropped into dilute sulphuric acid.
Explain why a chemical reaction is always accompanied by an energy change, and use this to explain why some reactions warm their surroundings while others cool them. [4]
The compounds of reactive metals are more stable to heat than the compounds of less reactive metals. Using metal nitrates, explain this statement with reference to three different metals from different parts of the activity series. [5]
Four conditions can bring about a chemical change: close contact, heat, light and pressure. Explain, with one example of each, why each condition is able to start a reaction. [5]
2.15 Equations, Conversions & Identification
- Write balanced equations for the following, and state the type of reaction in each case. [5]
- Iron burns in chlorine gas.
- Calcium hydrogen carbonate is heated.
- Chlorine gas is passed through potassium iodide solution.
- Barium chloride solution is added to sodium sulphate solution.
- Potassium hydroxide solution is added to dilute nitric acid.
- Complete and balance the following equations. [4]
- \(\mathrm{Cu(OH)_2} + \mathrm{H_2SO_4} \rightarrow \dots + \dots\)
- \(\mathrm{Zn(NO_3)_2} \xrightarrow{\ \Delta\ } \dots + \dots + \dots\)
- \(\mathrm{FeS} + \mathrm{H_2SO_4} \rightarrow \dots + \dots\)
- \(\mathrm{K_2O} + \mathrm{H_2O} \rightarrow \dots\)
- Name the gas evolved in each case, and state one test by which you would confirm it. [3]
- Dilute hydrochloric acid is added to sodium carbonate.
- Zinc is added to dilute sulphuric acid.
- Dilute sulphuric acid is added to iron sulphide.
- State what you would observe in each of the following. [4]
- A strip of magnesium ribbon is burnt in air.
- Potassium iodide solution is added to lead nitrate solution.
- Water is added to a lump of quicklime.
- Blue crystals of hydrated copper(II) sulphate are heated in a dry test tube.
- How would you obtain the following? Give the balanced equation in each case and state the type of reaction. [4]
- Calcium oxide from calcium carbonate.
- Copper from copper sulphate solution.
- Silver chloride from sodium chloride solution.
- Sodium hydroxide from sodium oxide.
- Identify the substance in each case and give a reason for your answer. [4]
- A white solid which, on heating, gives a colourless gas that turns lime water milky, leaving a residue that is yellow when hot and white when cold.
- A white solid which, on heating in a test tube, gives two colourless gases that recombine on the cooler upper walls.
- A green crystalline solid which turns white and then brown on heating, giving off gases with the pungent smell of burning sulphur.
- A colourless solution which gives a thick bright yellow precipitate with potassium iodide solution.
- Fig. 2.7-1 shows a test being carried out on the gas coming from a metal and an acid. [3]
- Name the piece of apparatus being held at the mouth of the test tube, and state its purpose here.
- Name the gas being tested and state the observation that identifies it.
- Write the balanced equation for the reaction taking place in the tube.
- Fig. 2.8-3 shows a gas being passed from one test tube into another. [3]
- Name the bent tube connecting the two test tubes and state its purpose.
- State what is seen in the second test tube, and write the equation for the change taking place there.
- State what would be seen if the gas were passed for a long time, and explain why.
- A student is given three colourless solutions labelled X, Y and Z. Solution X gives a white precipitate with barium chloride solution. Solution Y gives a white curdy precipitate with silver nitrate solution, which dissolves in ammonium hydroxide. Solution Z gives no precipitate with either reagent, but gives brisk effervescence with dilute hydrochloric acid. [4]
- Name the acidic radical present in X and in Y.
- Name the gas evolved from Z and state the test for it.
- Write a balanced equation for the reaction of X with barium chloride, taking X to be sodium sulphate.
- State the type of reaction taking place in each of the three tests.