Chemicalreactions

C H A P T E R 2

Stimulus questions1 How can we tell that a

chemical reaction hasoccurred?

2 What makes glow-wormsglow?

3 What puts the blue in bluejeans?

4 How are hydrogen and oxygenmade from water?

5 What is the chemical name ofrust?

6 How do antacids work?

OutcomesChemical science

6.1Relate the properties of fundamental groupings of substances to thenature of their constituent particles.

6.3Specify the characteristics, chemical reactions and usefulness tosociety of groups of similar substances.

6.4Represent chemical change, using chemical symbols and formulas.

6.5 ext.Describe the production and uses of substances with unusual andspecialised properties.

Fig 2.1.2 Melting is a physical change. Burning methane is a chemicalchange.

Chemical reactions are occurringconstantly: inside us, around us, in the soil,the air, on distant stars and planets—absolutely everywhere! Some reactions, likefireworks, are quick and violent. Otherreactions, like the reactions that occur inevery cell in your body, are more difficult tosee. So how do we know whether or not achemical reaction has taken place?

Signs of chemicalnchangen

We say that a chemical reaction has occurred if at leastone new substance has been formed. This differs from

SCI3 32

physical change, where no new substances are created.The melting of ice is an example of physical change. Itis simply a change of state; solid water turns into liquidwater. Burning methane gas is an example of chemicalchange because water vapour and carbon dioxide areformed, and they weren’t present at the start of thereaction.

The substances present at the start of a reaction arecalled the reactants, and the new substances formed arecalled the products. We can write word equations torepresent chemical reactions. These are written in theform:

reactants → productsChemical equations are useful because they provide

a quick and easy way to represent complex reactions.For example, the reaction of limestone (calcium

carbonate) and sulfuric acid produces calcium sulfate,carbon dioxide and water. This can be written as:

calcium carbonate + sulfuric acid →calcium sulfate + carbon dioxide + waterUsing chemical symbols, this reaction is written:CaCO3 + H2SO4 → CaSO4 + CO2 + H2O

Unit 2.1 What arechemical reactions andwhy do they happen?

Fig 2.1.1 Chemical reactions occur constantly in this nebula.

When ice melts nonew substances form.

When methane burns,carbon dioxide andwater are formed.

change is observed, but then theoriginal colour reappears shortly after.This can mean that a new (coloured)substance formed, but then thereaction reversed and the reactantsreformed, so overall no changeoccurred in the system. Manyreactions occur in more than one step.Sometimes a temporary colour changewill be observed as the reactantsrearrange to form the products. If aproduct breaks down easily, it is saidto be unstable. We will limit ourdiscussion to reactions that producestable, coloured products, so we willlook only for permanent changes in colourto indicate that a reaction has taken place.

A gas is given offIf a reaction is taking place in solution (in a liquid), it is very

easy to see a gas beingproduced becausebubbling will be observed.With other reactions, itcan be more difficult tosee a gas being producedbecause most gases, likeoxygen, hydrogen,nitrogen andcarbon dioxide,are colourless andodourless.

Fig 2.1.5 The chemical structure of indigo

33 Chapter 2 Chemical reactions

U N I T 2 . 1

Blue jeansMany colouredsubstances havepractical uses. Theearly Britons used ablue dye called woad todye cloth. Theyextracted this from theIsatis tinctoria plant.Indigo is the maincomponent of woad,and is used today to dyeblue denim jeans.

Sciencesnippet

Prac 1p. 38

Fig 2.1.3 The formation of a gas is a sign of chemical change.

In this case, the reactants are calcium carbonate(CaCO3) and sulfuric acid (H2SO4). The products arecalcium sulfate (CaSO4), carbon dioxide (CO2) andwater (H2O).

There are several signs which will tell us whether ornot a chemical reaction has occurred.

A chemical reaction has definitely occurred if one ormore of the following is observed:• There is a permanent colour change.• A gas is given off.• Energy is produced or absorbed.• A precipitate (solid) forms in solution.• One metal deposits on another.

Permanent colour changeIt is important to distinguish permanent colour changesfrom temporary colour changes. Sometimes a colour

Fig 2.1.4 Left:The iron in cars can react with oxygen to form iron(III) oxide—rust.The effects can be very damaging.Right:This patina protects the copper from further corrosion, unlike rust which does not protect iron at all.

C C

O

OH

H

N

N

dilute H2SO4

C02 gas

limestone

Energy is produced or absorbedMany reactions absorb or produce energy, usually in theform of heat. Imagine a reaction taking place in abeaker. If you feel the beaker as the reaction is takingplace and it gets colder, the reaction is absorbing energyfrom the surroundings. Reactions that absorb heatenergy are called endothermic. If you feel the beakerand it gets warmer, the reaction is releasing energy tothe surroundings. Reactions that produce heat energyare called exothermic.

Endothermic: energy + reactants → productsExothermic: reactants → products + energyOne example of an exothermic reaction is the

burning of fuels. The heat produced can then be usedto do things like make cars move, or to produceelectricity in power stations.

Plants make their own food by the process ofphotosynthesis. This reaction can’t proceed unlesssunlight is present. Chlorophyll in plants captures theSun’s energy and uses it to start photosynthesis, anendothermic reaction. Photosynthesis is actually a seriesof reactions. The overall reaction can be written as:

energy + carbon dioxide + water →glucose + oxygenThe energy that is taken in is stored in the chemical

bonds of the products. When the products are brokendown, the energy is released. Respiration is thechemical reaction that takes place in our cells and givesus energy to stay alive. The chemical equation forrespiration is the opposite of photosynthesis:

glucose + oxygen →carbon dioxide + water + energyThis is an exothermic reaction.

SCI3 34

A precipitate formsSoluble substances are those that dissolve easily in aliquid, while insoluble substances will not dissolve. Asolution is made up of a solute and a solvent. A soluteis a soluble substance that dissolves in the solvent(usually water). For example, a solution of sodiumchloride (table salt) is made up of solid sodium chloridedissolved in water. The salt dissolves because the sodiumand chloride ions that it is made of are more attracted tothe water than to each other, so they break away fromeach other and spread out to form a clear solution. Aprecipitate is an insoluble substance that forms whentwo clear (but not always colourless) solutions are addedtogether. The precipitate is observed as cloudiness.

If left, the cloudiness will diminish as the precipitatecollects at the bottom of the container. The precipitateforms because the ions are more attracted to each other

Carbon monoxide poisoning

Carbon monoxide (chemical formula CO) is a deadly gas emitted by cars. It is

produced when carbon-based fuels like petrol burn in a limited supply of

oxygen. Haemoglobin is a molecule that transports oxygen around your body.

Carbon monoxide, which is odourless and colourless, is extremely toxic because

it binds to haemoglobin 200 times more strongly than oxygen does. This leaves

no space for the oxygen to bind onto and move around the body, so the cells

quickly become starved of oxygen and die.

Science snippet

BioluminescenceBioluminescence is the light produced by chemical reactions in livingorganisms. It is a type of chemiluminescence, which is light producedby a chemical reaction. Bioluminescence is sometimes called ‘coldlight’ because it doesn’t give out much heat. This is good news for theanimals that use this light, like glow-worms, fireflies and some deep-sea jellyfish, because they would quickly overheat and die if much heatwas produced along with the light.

Science snippet

Fig 2.1.6 This firefly beetle shows bioluminescence on its lowerabdomen. It is believed to attract other beetles.

Prac 2p. 38

Fig 2.1.10 Copper is displaced by zinc.

Fig 2.1.9 Activity series of metals

Fig 2.1.8 Mixing solutions of barium nitrate and sodium sulfate produces a precipitate of barium sulfate.

than to the water molecules. This means that they tendto clump together, rather than spreading out in thewater. Figure 2.1.8 shows barium sulfate precipitatingfrom two clear solutions. Clear solutions may becoloured. ‘Clear’ just means that you can see rightthrough them.

One metal deposits on anotherDifferent metals have different degrees of reactivity.Outer-shell electrons are also called valence electrons.Some metals give up their valence electrons very easilyand so are very reactive. The alkali metals (Group I) arethe most reactive metals, with francium being the mostreactive of the group. Francium is the most reactivebecause it is the biggest. Its single valence electron isquite a long way from the positive nucleus so it is heldvery weakly. A very small amount of energy will makefrancium react and give up its outer electron. Othermetals, like gold and silver, are very unreactive. Thismeans that they don’t give up their electrons as easily.

The activity series of metals lists the metals in orderfrom most reactive to least reactive.

35 Chapter 2 Chemical reactions

U N I T 2 . 1

When one metal deposits on (settles on top of)another, the reaction is known as a displacementreaction. The basic rule is that a more reactive metal willdisplace a less reactive metal in solution. This means thatif you have a solution of a particular metal salt, and youplace a solid piece of a more reactive metal in thesolution, a reaction will take place. The electrons fromthe more reactive metal will be transferred to the ions ofthe less reactive metal, which will become solid anddeposit on the surface of the more reactive metal. Figure2.1.10 shows what happens when a piece of zinc isplaced in a solution of copper sulfate.

Why do chemicalnreactions occur?n

Some reactions happen naturally. These are calledspontaneous reactions, and one example is the rusting ofiron. Spontaneous reactions also include those that need

Fig 2.1.7 A solution of coppersulfate

NaNO3 solutionwith BaSO4precipitateNO3–

NO3–

Ba2+

SO42–

Na+

Na+

Ba(NO3)2 solution Na2SO4 solution

Na+

Na++ NO3–

NO3–

Ba SO4

zinc ions formwhen zinc atomsgive up electronsto copper ions

SO42–

SO42–

SO42–

SO42–

Cu2+ Cu2+

Cu2+

Zn2+

a coating of solidcopper forms onthe surface of thezinc as copper ionsaccept electrons tobecome copper atoms.

copper sulfatesolution

solid zinc

most reactive least reactive

K N a C a M g A l Z n Fe P b C u H g A g P t A u

a little help getting started, but then will keep going bythemselves, like burning gas in a Bunsen burner, whichneeds a spark or flame to get it going. Other reactionsneed a continual energy input to keep them going, likethe electrolysis of water, where an electric current isused to split water up into hydrogen and oxygen gas.These are called non-spontaneous reactions. So why dosome reactions proceed with no help from us, whileothers are very difficult to start? What are the drivingforces behind chemical reactions?

The energy content of substancesChemicals have a certain amount of energy stored intheir bonds. This energy varies from substance tosubstance. Reactions are more likely to occur if thereactants have more stored energy than the products.This means that exothermic reactions are more likely tooccur than endothermic reactions. The burning ofmethane, the gas that comes out of the gas taps in thelaboratory, is an exothermic reaction. We have to add aspark to start the reaction, but then it will just keepburning by itself. The equation for the reaction is:

methane + oxygen →carbon dioxide + water + energyCH4 + 2O2 → CO2 + 2H2O + energyNotice the ‘2’ in front of the oxygen and water.

These are put there to balance the equation. Becausematter cannot be created or destroyed, there must bethe same number and types of atoms on each side of theequation. This is achieved by putting whole numbers infront of the chemical formulas in the equation. At thisstage, you don’t need to know how to balance equations.You will learn more about this in SCI 4, Chapter 1. For now, we’ll just get used to seeing what the equationslook like.

The products, carbon dioxide and water, have lessstored energy than the reactants, so this reaction is likelyto proceed as written.

The hydrolysis (splitting up) of water doesn’t happenspontaneously. We have to run an electric currentthrough water to make this reaction happen. The equation for this reaction is:

water + energy → oxygen + hydrogen2H2O + energy → O2 + 2H2

SCI3 36

The products, oxygen and hydrogen, have more storedenergy than the reactants, so this reaction is not likely toproceed easily. Both oxygen and hydrogen are diatomicgases. They exist as two atoms joined together to make amolecule. This is why they are written as O2 and H2

rather than just O and H.

The drive towards increaseddisorderThe other driving force behind chemical reactions is thedrive towards increased disorder. Imagine your bedroomall neat and tidy. What would it look like if you didn’t putany energy into cleaning it? All systems, including yourroom, have a tendency to become more disordered. Ifyou think about the particles in solids, liquids and gases,it should be easy to see that solids are more ordered thanliquids, which are more ordered than gases. This meansthat reactions which cause solids to turn into liquids, orliquids to turn into gases, are more likely to occur.

This drive towards disorder is related to the probabilityof a reaction occurring. The more ways there are for areaction to occur, the more likely that it will. Forexample, there are more ways to arrange water moleculesin the disordered liquid state than in the ordered icestate, so liquid water is more probable than ice.

Fig 2.1.11 The hydrolysis of water produces twice as much hydrogen gasas oxygen gas. Why do you think that is?

O2 H2 gas

water

electriccurrent

So why doesn’t the electrolysis of water occurnaturally? After all, this is a case where a liquid (water)turns into two very disordered gases (oxygen andhydrogen). The answer lies in the fact that one of thedriving forces can be stronger than the other. With the

electrolysis of water, the fact that the products havemuch more stored energy than the reactants overridesthe fact that the products are more disordered than thereactants. Hence, overall this becomes a non-spontaneous reaction.

37 Chapter 2 Chemical reactions

U N I T 2 . 1

1 Give three examples of chemical reactions that happenin your home.

2 Are the following examples of chemical change orphysical change?a Cutting up cheeseb Making toastc Burning gasd Melting chocolatee Freezing cordialf Water evaporatingg Putting a soluble aspirin tablet in water

3 How are chemical equations useful?4 What three things are produced when calcium

carbonate reacts with sulfuric acid?5 What are four signs of chemical change?6 For each of the following, state if a chemical change

has occurred. Give a reason for your answer.a A student mixes two unknown solutions together and

notices a cloudiness forming.b Solid purple iodine crystals are heated slightly and a

purple cloud of iodine gas is observed.c When nitric acid is poured onto limestone, bubbling

is seen.d Two colourless solutions at room temperature are

mixed. After a minute, the temperature of the mixtureis 60°C.

7 What is a solution?8 Explain what is meant by ‘a solution is clear, but not

always colourless’.9 Describe how precipitates form.

10 Why is francium the most reactive metal in Group I?11 Arrange these metals in order from most reactive to least

reactive: gold, copper, magnesium, zinc, potassium,silver.

12 If solid magnesium is placed in a silver nitrate solution,solid silver is produced. Explain how this happens.

U n i t 2 . 1 Questions13 Write word equations for the following reactions:

a When copper is added to nitric acid, copper nitrate,nitrogen monoxide and water are formed.

b If sulfuric acid is poured onto solid sodium carbonate,bubbles of carbon dioxide are produced, as well aswater and sodium sulfate.

c Magnesium burns easily in oxygen, producingmagnesium oxide.

14 If burning methane is a favourable reaction, why do wehave to light a match to make it burn?

15 What are the two driving forces behind chemicalreactions?

16 Which are more ordered, solids, liquids or gases? Explain.17 For each of the following reactions:

a water + energy → hydrogen + oxygenb methane + oxygen → carbon dioxide + water + energystate:

i What are the reactants?ii What are the products?iii Is the reaction exothermic or endothermic?

1 Paints have special pigments (chemicals) in them that givethem their unique colours. See if you can find out the nameof the coloured chemical in monastral blue.

2 What is the chemical composition of a copper patina?3 What is the function of the bioluminescence given out by

deep-sea animals? What are some of the animals that usethis light?

4 Research the most reactive metal, francium. How was itdiscovered? Is it found in nature? What are somecompounds containing francium?

5 Find out how light sticks work.6 Find out what entropy is.

U n i t 2 . 1 Research /Extension

Questions1 What gas was formed in reaction 2?2 What do you think would happen if the zinc in reaction

5 was replaced with silver?

Unit 2.1 Prac 2Light sticks—chemiluminescence

You will needTwo 250 mL beakers, ice, hot tap water, 2 light sticks (fromscuba-diving store)

What to do1 Set up your two beakers, one with a mixture of ice and

water, the other with hot water. Your beakers should befilled to the 200 mL mark.

2 Activate your light sticks by bending them. This snapsthe capsule inside and allows the chemicals to mix.

3 Place one light stick in ice and the other in hot water.4 After a few minutes, take them out and compare the

intensity of the light.

Questions1 Were the light sticks as bright as each other or different?

Why do you think that is?2 What would happen to an activated light stick if you put

it into the freezer?3 What are some uses for light sticks?

SCI3 38

Unit 2.1 Prac 1Signs of chemical change

You will needSolid copper carbonate, magnesium, dilute nitric acid, splint,matches, Bunsen burner, test-tube holder, test-tube rack, dilutesodium hydroxide, thermometer, dilute barium nitrate, dilutesodium sulfate, dilute copper sulfate, solid zinc, 5 test tubes(1 with stopper), lab coat, safety glassesNote: 0.1 M is an appropriate concentration for thesesolutions, but anything between 0.1 M and 2.0 M would besuitable.

What to do1 Copy a larger version of the following results table into

your book:

2 Carefully heat a small amount of copper carbonate in atest tube. Ensure that the test tube is pointed away fromother people. Stop as soon as you see a colour change.Record your observations.

3 Add a small piece of magnesium to 2 cm of nitric acidin a test tube. Stopper the tube to collect some gas.Have your lab partner light a splint and place it near themouth of the test tube.

4 Record the temperature of 2 cm of the nitric acidsolution. Add 2 cm of sodium hydroxide and record thenew temperature.

5 Place a small piece of zinc into 2 cm of dilute coppersulfate solution. Record your observations.

U n i t 2 . 1 Practicalactivities

Reaction Reactant/s Observations Conclusionnumber

1 Copper carbonate

2 Nitric acid and magnesium

3 Dilute barium nitrate andsodium sulfate

4 Dilute nitric acid and sodiumhydroxide

5 Zinc and dilute copper sulfate

light sticklight stick

hot waterice water

Fig 2.1.12

Fig 2.2.1 The formation of sodium chloride

39 Chapter 2 Chemical reactions

It is not possible to write or understandchemical equations if you don’t understandhow compounds are named, so before wecontinue talking about reactions, it is worthspending a bit of time learning how toname compounds. At this stage, we won’tlearn how to name compounds that containboth carbon and hydrogen. That is part of abranch of chemistry called organicchemistry, which you will learn about inSCI 4, Chapter 1. For now, we will lookonly at ionic compounds and inorganiccovalent compounds.

Ionic compoundsnIonic compounds are made up of positive and negativeions (atoms that have lost or gained electrons) heldtogether by electrostatic attraction—the attractionbetween opposite charges. Ionic compounds are alsocalled salts. In almost all cases, the positive ions aremetal ions. The only exception which we willencounter is the ammonium ion, NH4

+. The formulaNH4

+ means that this ion is made up of one nitrogenatom and four hydrogen atoms chemically bondedtogether. It has an overall +1 charge.

As you learned in Chapter 1, ions are formed whenelectrons are completely transferred from one atom toanother. The reason they do this is so that they can gainelectronic configurations that look like that of thenearest noble gas, which makes them more stable (and

hence less reactive). For example, sodium is in Group Iand has the electron configuration 2,8,1. This meansthat it has one electron in its third energy level, which iscalled the outer shell or valence shell. If it loses this oneelectron, its electron configuration becomes 2,8 whichis the same as neon’s, the nearest noble gas. It now has11 protons but only 10 electrons. This is why sodiumalways forms an ion with a +1 charge.

When the ions come together to form compounds,they combine in a ratio that allows the compound tohave a net charge of zero. There must be sufficientnegative charge to balance the positive charge and viceversa.

Sodium and chloride ions combine in a 1:1 ratiobecause sodium ions have a +1 charge and chlorideions have a –1 charge. Add these charges together: +1 +(–1) = 0. Thus, one of each ion joins with the other one

Unit 2.2 Namingcompounds

Cl Cl–Na Na+

+ +Na+

An electron is transferredfrom sodium to chlorine.These positive and negativeions are attracted to eachother and form a crystalwhere the ions are stackedto maximise attraction.

sodium chlorine sodium ion chloride(2, 8, 1) (2, 8, 7) (2, 8) (2, 8, 8)

––

––

–

–––

––

––

–

++

++

+++

+

+

++

+

• Metals in Group II (the alkaline earth metals) alwaysform ions with a +2 charge.

• Metals in Group III always form ions with a +3charge.

• Metals in Group IV may have either +2 or +4charges. Assume +2 unless you are told otherwise.Transition metals have variable valencies, but most

form ions with a +2 charge. If a metal has more thanone common ion, it is denoted with Roman numerals.For example, copper(I) = Cu+, copper(II) = Cu2+,iron(II) = Fe2+, iron(III) = Fe3+.

We will not look at compounds of metals in GroupsV and VI. They can also have variable valencies.

Recall from the previous chapter that Group VIIIelements have stable outer electron shells and do noteasily react.

Non-metal ionsElements in Group VII (the halogens) always form ionswith a –1 charge.

Elements in Group VI always form ions with a –2charge.

Elements in Group V always form ions with a –3charge.

Non-metal elements in Group IV (carbon andsilicon) may form –4 ions.

Fig 2.2.3 These atoms lose electrons to get a noble gas electronicconfiguration.

to give a compound with a net charge of zero. Theformula is NaCl, and the name of this compound issodium chloride. Other examples are given in the tablebelow.

Positive ion Negative ionand electron and electronconfiguration configuration Formula Name

Mg2+ (2, 8) Cl– (2,8,8) MgCl2 Magnesium chloride

Na+ (2, 8) O2– (2,8) Na2O Sodium oxide

Al3+ (2,8) S2– (2,8,8) Al2S3 Aluminium sulfide

Ca2+ (2,8,8) N3– (2,8) Ca3N2 Calcium nitride

Note that only subscript numbers appear informulas. These indicate the number of the atom or iondirectly before the number in the formula. For instance,MgCl2 indicates that there are one magnesium ion andtwo chloride ions in the formula. No charges appear inthe overall formula because once they are balanced,there is zero charge.

Metal ionsThe formation of ions was covered in Chapter 1. Recallthat:• Metals in Group I (the alkali metals) always form

ions with a +1 charge .

SCI3 40

Fig 2.2.2 Formulas of ionic compounds must be balanced so that the netcharge on the compound is zero.

Cl–Cl– Na+Na+

1 magnesium ion joins with 2 chloride ions to formmagnesium chloride.

+

1 sodium ion joins with 1 chloride ion to form sodium chloride.

Mg2++ Cl– + Cl– Cl– Mg2+ Cl–

Na+ + Cl– NaCl

Mg2+ + 2Cl– MgCl2

Mg

Mg2+ + 2e–

2, 8, 2

2, 8

Al

Al3+ + 3e–

2, 8, 3

2, 8

e–e–

e–

e–

e–

e–

e–e–

e–

e–

e–e–

e–e–

e–

e–

e–e–

e–

e–

e–e–

e–e–

e–

e–

e–

e–

e–

e–e–

e–

e–

e–e–

e–e–

e–

e–

e–e–

e–

e–

e–e–

Fig 2.2.4 These atoms gain electrons to get a noble gas electronicconfiguration.

Remember, elements inGroup VIII (the noble gases) donot form ions.

Polyatomic ionsIons made up of more than onetype of atom are calledpolyatomic ions or radicals.These have special names. Thetable below shows some of themore common polyatomic ions.

When more than onepolyatomic ion is required in aformula, brackets are used. Forexample, in sodium sulfate,Na2SO4, only one sulfate ion isneeded to balance thecharge so no brackets areused. For aluminiumsulfate, Al2(SO4)3, threesulfate ions are required sobrackets are used.

41 Chapter 2 Chemical reactions

U N I T 2 . 2

Franciumfluoride

The most ioniccompound possible is

francium fluoride,

because francium is the

most reactive metal

and fluorine is the most

reactive non-metal. If

these two elements

ever came together, the

result would beexplosive indeed!

Sciencesnippet

Ionic or covalent?Some compounds contain both ionic and covalent bonding. Theammonium ion is held together by covalent bonding (the sharing ofelectrons). However, when it joins to a negative ion like Cl–, thechloride ion, it forms an ionic bond.

Science snippet

Inorganic covalentncompoundsn

When non-metals come together, covalent bonds areformed. In these bonds, no electrons are transferred.Atoms simply share electrons in order to gain a noblegas electronconfiguration. Figure2.2.5 shows how somecovalent moleculesshare electrons.

When namingthese compounds,prefixes are used toshow how many ofeach atom we have—see the table on theright.

Ion name Formula

Hydroxide OH–

Sulfate SO2–

4

Carbonate CO2–

3

Hydrogen carbonate HCO3–

Ammonium NH+

4

Nitrate NO–

3

Number of atoms Prefix

1 mono

2 di

3 tri

4 tetra

5 pent

6 hex

7 hept

8 oct

9 non

10 dec

Cl–

Cl + e–

2, 8, 7

2, 8, 8

O2–

O + 2e–

2, 8, 6

2, 8, 8

e–e–e–e–

e–e–

e–e–

e–e–

e–

e–

e–

e–e–

e–e–

e–e–e–e–

e– e–e–e–e–

e–e–e–

e–

e–

e–e–

e–e–

e–e–e–e–

e–e–

e–e–e–e–

e–

e–

e–

e–e–

e–

e–e–e–e–

e– e–e–e–e–

e–e–e–

e–

e–

e–e–

e–e–

WS 2.1

Fig 2.2.5 These atoms share electrons to gain noble gas electronicconfigurations.

OC

H

H

H

HH H

NH H

H

ammonia, NH3

water, H2Omethane, CH4

8 Name the following compounds:a NH4Clb LiOHc Ag2CO3

d ZnSO4

9 Write formulas for:a iron(III) chlorideb iron(II) chloridec copper(I) nitrated copper(II) nitrate

10 What is the total charge of:a four sodium ions?b eight manganese(IV) ions?c three nitride ions?

11 How many atoms of each type are in the followingformulas?a (NH4)2SO4

b K2Cr2O7

c Ca(OH)2

12 What is the difference between ionic bonding andcovalent bonding?

13 Name the following inorganic covalent compounds:a CO2

b N2O5

c SF6

d H2O2

14 Write chemical formulas for:a phosphorus trihydrideb oxygen dichloride

If there is only one of the firstatom, no prefix is used for that atom.

Examples are:CO2 = carbon dioxideCO = carbon monoxideN2O5 = dinitrogen pentoxideCCl4 = carbon tetrachlorideH2O = dihydrogen monoxide

SCI3 42

1 What holds the ions together in ionic compounds?2 Why is sodium more stable as a +1 ion than as a neutral

atom?3 Magnesium oxide has a higher melting point than sodium

chloride. What does this tell you about the strength of theattractive forces between ions in these compounds?

4 What is a polyatomic ion? Give an example.5 Using the Periodic Table on page 8 write chemical

formulas for:a sodium bromideb magnesium sulfidec calcium fluorided lithium nitridee aluminium carbide

6 Name the following ionic compounds:a RbBrb K2Sc BeOd Na3N

7 Write formulas for the following compounds containingpolyatomic ions:a Sodium sulfateb Magnesium hydroxidec Strontium carbonated Lithium nitratee Ammonium oxide

Fig 2.2.6 Some ionic and covalent compounds you might know

Prac 1p. 43

U n i t 2 . 2 Questions

43 Chapter 2 Chemical reactions

U N I T 2 . 215 When covalent molecular compounds melt, only the

bonds between molecules are broken. The moleculesthemselves stay intact. Ammonia (NH3) has a meltingpoint of –78°C while dinitrogen monoxide (N2O) has amelting point of –91°C. What does this tell you aboutthe relative strength of their intermolecular bonds?

1 Some transition elements form brightly colouredcomplexes. Find out what a complex is and someexamples of such complexes. What special purposes arethey used for?

2 Even though the noble gases are usually unreactive, theydo form compounds under certain conditions. Find outwhat compounds the noble gases form.

3 Even though Al2O3 is an ionic compound, its bonding ispartially covalent. Elements close to the metals/non-metals border often have properties of both groups. Findout which elements are classed as metalloids. Chooseone example and state how it is like a metal and how itis like a non-metal.

4 Some compounds go by a common name, rather thantheir systematic chemical names. Water is an example ofthis. Its chemical name is dihydrogen monoxide. Whatother compounds can you find that go by a commonname rather than their chemical names? See if you canfind out where the common name came from.

Unit 2.2 Prac 1Action of heat on ionicand covalent compounds

You will needSolid samples of various ionic and covalent compounds thatdo not produce toxic fumes on heating (e.g. wax—candle orparaffin), graphite, sodium chloride, potassium nitrate,

Bunsen burner, heat mat, metal spatulas, wooden pegs,safety glasses, lab coat

What to do1 Draw up a suitable table to record your results. The

table below is an example.

2 Take a small amount of your first sample on the metalspatula. Hold it in the hottest part of the Bunsen burnerflame for no more than 5 seconds.

3 Record your results.4 Thoroughly clean the spatula. Ask your teacher before

disposing of any sample down the sink.5 Repeat steps 2 to 4 for your other samples.

Questions1 Write a paragraph summarising your results.2 Explain your observations in terms of the strengths of

bonds between ions or molecules.3 Do you think the following substances would have high

or low melting points?a Sulfur (covalent)b Magnesium carbonate (ionic)c Iodine (covalent)d Lithium nitrate (ionic)

U n i t 2 . 2 Research /Extension

U n i t 2 . 2 Practicalactivity

sample

heat-proof mat

Bunsen burner

spatula

Fig 2.2.7

Substance Ionic or covalent Melts on heating?

Sodium chloride

Wax

Potassium nitrate

Fig 2.3.1 One atom of carbon joins with one molecule of oxygen to formone molecule of carbon dioxide.

Even though each substance is unique,similar substances behave similarly inchemical reactions. This allows us to groupreactions into several general categories.This unit introduces some of the mainreaction types. You can find informationabout other types in SCI 4, Chapter 1.

Combinationnreactionsn

In combination reactions, two or more substancescombine to form one new substance. For twosubstances combining, these reactions have the generalequation:

X + Y → XYFor example, carbon and oxygen can combine to

form carbon dioxide:carbon + oxygen → carbon dioxideNow that we have been through some naming, we

will start to use the correct chemical formula equation(called ‘chemical equation’ for short) as well as the wordequation. The above reaction is written as:

C + O2 → CO2

O2 is used instead of O by itself because the oxygenin the air around us exists as diatomic molecules.‘Diatomic’ means that two oxygen atoms bond togetherto form a molecule. This makes them more stable. In chemical equations, we always write the formula ofsubstances as they would be if they were in the roomwith us now.

Another example of a combination reaction is:sodium + chlorine → sodium chloride2Na + Cl2 → 2NaCl

SCI3 44

Decompositionnreactionsn

Decomposition reactions are the opposite ofcombination reactions. One substance breaks down toform two or more new substances. For substances thatbreak down to form two new substances, the generalequation can be written:

XY → X + YExamples are:

• calcium carbonate → calcium oxide + carbondioxideCaCO3 → CaO + CO2

• magnesium hydroxide → magnesium oxide + waterMg(OH)2 → MgO + H2O

Precipitationnreactionsn

Precipitation reactions result in an insoluble solidbeing formed when two clear solutions are mixed. Theycan be written as:

soluble salt A + soluble salt B →insoluble salt C + soluble salt D

Unit 2.3 Reaction types

O

C + O2 CO2

C

O

O

+ C O

Prac 1p. 47

Fig 2.3.3 Van Gogh (A Wheat Field with Cypresses, 1889) used chrome yellowin his series of wheatfield paintings.

A subscript ‘(s)’ can be used in the chemicalequation to show which is the insoluble salt. This samesubscript is used to show any solid substance in achemical reaction. When solutions of the soluble saltssilver nitrate and sodium chloride are mixed, silverchloride precipitates. The sodium and nitrate ions areleft dissolved in solution. A subscript ‘(aq)’ is used toshow aqueous (from the Latin word aqua, meaning‘water’) species—those dissolved in water. The word andchemical equations for this reaction are:

silver nitrate + sodium chloride →silver chloride + sodium nitrateAgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

Solubility rules (see the table below) help us to workout which substance in the mixture is precipitating. Forexample, in the above reaction, it can’t possibly besodium nitrate because all sodium salts are soluble andall nitrate salts are soluble.

45 Chapter 2 Chemical reactions

U N I T 2 . 3

Solubility of common inorganic compounds in water

Negative ions Positive ions Compunds with (anions) + (cations) → solubility:

Acetate CH3COO– All Soluble

All Alkali ions, Li+, Na+, SolubleK+, Rb+, Cs+, Fr+

All Ammonium ion NH4+ Soluble

All Hydrogen ion H+(aq) Soluble

Chloride Cl– Ag+, Pb2+, Hg2+, Low solubilityBromide Br– Cu+

Iodide I– All others Soluble

Hydroxide OH– Alkali ions, H+(aq), NH4

+, SolubleSr2+, Ba2+, Ra2+

All others Low solubility

Nitrate NO3– All Soluble

Phosphate PO43– Alkali ions, H+

(aq), NH4+ Soluble

Carbonate CO32– All others Low solubility

Sulfate SO42– Ca2+, Sr2+, Ba2+, Low solubility

Pb2+, Ra2+

All others Soluble

Sulfide S2– Alkali ions, H+(aq), NH4

+, Soluble

Be2+, Mg2+, Ca2+,

Sr2+, Ba2+, Ra2+

All others Low solubility

Fig 2.3.2 Lead iodide has a disctinctive yellow colour.

Chrome yellow

Chrome yellow was a pigment discovered by the French chemist

Vauquelin in the nineteenth century. Chrome yellow is a bright yellow

substance, lead chromate, obtained through precipitation: lead nitrate

+ sodium chromate → sodium nitrate + lead chromate. Chrome yellow

is the pigment used by Vincent Van Gogh to create the darkest yellow

of the wheatfield in his painting, A Wheat Field with Cypresses.

Science snippet

Neutralisationnreactionsn

Neutralisation reactions occur when an acid is addedto a base, forming water and at least one othersubstance. A base is anything that can neutralise anacid, like metal hydroxides. For example,

acid + metal hydroxide → salt + water:hydrochloric acid + potassium hydroxide →potassium chloride + waterHCl + KOH → KCl + H2ONitric acid and sulfuric acid are two other common

acids.Acids, bases and salts will be explained in more

detail in Unit 2.4.

Combustion reactionsnA combustion reaction is simply burning a substance inoxygen, so O2 is always a reactant. The products willvary, depending on what substance is being burned.Examples of combustion reactions are:• propane + oxygen → carbon dioxide + water

C3H8 + 5O2 → 3CO2 + 4H2O• magnesium + oxygen → magnesium oxide

2Mg + O2 → 2MgOYou can see that the combustion of magnesium is

also an example of a combination reaction. Reactionscan sometimes fall into more than one general category.

Displacementreactionsn

As mentioned in Unit 2.1, displacement reactions causeone metal to be deposited on another. The reason theydo this is that the less reactive metal ion will acceptelectrons from the more reactive metal. This is the basisof batteries. In batteries, a similar exchange of electronsoccurs. This flow of electrons forms an electric current,which can then be used to power things like torchesand portable CD players. This type of reaction is alsoknown as a redox reaction, because it involves bothoxidation and reduction. Oxidation is the loss of

SCI3 46

electrons. Reduction is the gain of electrons. Displacement reactions can be used to coat metals.

For instance, if you wanted to coat a piece of iron withsilver, you could simply dip it into a solution of silvernitrate. The silver ions would accept electrons from theiron atoms on the surface, because silver is less reactivethan iron. Some of the iron atoms would form ions anddissolve, while some of the silver ions would becomesolid silver and coat the iron. In the reaction betweenAg+ and Fe(s), the silver ions are reduced to Ag(s) and thesolid iron is oxidised to Fe3+.

This reaction could be written as:iron + silver nitrate → iron(III) nitrate + silverFe(s) + 3AgNO3(aq) → Fe(NO3)3(aq) + 3Ag(s)

This would not produce a very even coating,however. Industrial silver plating is carried out quitedifferently. Electroplating is the name for the processthat causes metal atoms to be deposited on the surfaceof a conducting substance. By using an outside sourceof electricity, non-spontaneous reactions can be made toproduce decorative or useful surfaces. Forexample, a zinc coating on copper helpsprotect the copper from oxidising.

How apistol works

A combustion reactionoccurs when a pistol isfired. A hammer orfiring pin detonates acap at the base of thebullet cartridge. Thisignites the propellantcharge. The combustionof the propellantproduces gas, whichforces the bullet out ofthe barrel.

Sciencesnippet

Prac 2p. 48

47 Chapter 2 Chemical reactions

U N I T 2 . 3

1 What is a combination reaction? Give an example.2 Why is oxygen written as O2 in chemical reactions,

rather than just O?3 What is a decomposition reaction? Give an example.4 What gas is given off when calcium carbonate is

heated?5 Describe what observations are made when a

precipitation occurs.6 a What do the subscripts (s) and (aq) mean?

b Sometimes the subscripts (g) and (l) are also used.What do you think they might mean?

7 What precipitate is formed when the following solutionsare mixed?a Silver nitrate and sodium chlorideb Mercury(I) nitrate and potassium iodidec Calcium nitrate and lithium carbonated Barium nitrate and sodium sulfate

8 Refer to the table of solubility rules on page 45. Whichof the following substances would be soluble in water?a BaSO4

b LiNO3

c CaCO3

d MgCl29 What is a neutralisation reaction? Give an example.

10 What products are formed when an acid is added to ametal hydroxide?

11 Name three different acids and give their chemicalformulas.

12 What products are formed when iron is added to silvernitrate?

13 What is a combustion reaction? How is it possible torecognise a combustion reaction?

14 Give an example of a reaction that is in two differentreaction categories.

15 What is a displacement reaction?16 What are oxidation and reduction?

1 Find out what combustion reaction drives the spaceshuttle. Research a NASA mission that did not achieveits objective.

2 Does the solubility of ionic compounds increase ordecrease as solutions get hotter? Explain.

3 Find out the main steps in the carbon cycle. What typesof reactions occur in this cycle?

4 How is industrial silver plating carried out?

Unit 2.3 Prac 1Precipitation of unknowns

You will needThe table of solubility rules on page 45; unknown 0.1 Msolutions labelled A, B, C, D, E—these are (not in order)sodium iodide, sodium chloride, sodium sulfate, sodiumcarbonate and sodium nitrate; 0.1 M solutions of silver,lead, calcium and barium nitrate; 20 semi-micro test tubes;Pasteur pipettes; lab coat; safety glasses; glovesHint: Cu2+ ions are blue in aqueous solution. Lead iodide isbright yellow.Note: Solutions from this experiment must not be washeddown the sink. They should be placed in a clearly markedwaste bottle. Gloves must be worn at all times.

U n i t 2 . 3 Questions U n i t 2 . 3 Research /Extension

U n i t 2 . 3 Practicalactivities

Add 10 dropstest solutionand mix.

step 1

10 drops ofunknownsolution

step 2 step 3

Check forcloudiness—holdit up to the lightif not sure.

+

Fig 2.3.4

Unit 2.3 Prac 2Electroplating

You will need6 V DC power source, 250 mL beaker, 2 insulated wireswith crocodile clips on one end, 1 very thin 7 × 4 cm stripof copper metal (coiled copper wire may also be used), 1stainless steel electrode, sand paper, tongs, washbottle ofdistilled water, 1 M zinc sulfate solution, 2 M nitric acid, labcoat, safety glasses, gloves

What to do1 Clean the copper with the sand paper.2 Dip the copper in the acid and then rinse with distilled

water. Don’t touch the part that will go in the solution.3 Attach the copper to the negative terminal of the power

source. Rest it in the beaker.4 Put 150 mL of 1 M zinc sulfate solution in the beaker.5 Attach the stainless steel electrode to the positive power

terminal.6 Turn the power on for about 3 minutes.7 Remove the copper and rinse with distilled water.

Questions1 What did the copper look like before and after the

experiment?2 Write a word equation for the reaction that occurred.3 Describe how you could have coated the copper with

nickel.

SCI3 48

What to do1 Draw up a suitable table, similar to the one below, to

record your results.

Unknown Silver Lead Calcium Bariumnitrate nitrate nitrate nitrate

A

B

C

D

E

2 Put about 10 drops of unknown A into each of 4 semi-micro test tubes.

3 Add 10 drops of silver nitrate solution to the first, 10drops of lead nitrate to the second, 10 drops of bariumnitrate to the third, and 10 drops of magnesium nitrateto the fourth. Record your results.

4 Repeat steps 2 and 3 for each unknown solution.5 Use the table of solubility rules to work out which

solution is which.

Questions1 Were any of your results inconclusive? If so, suggest a

reason.2 If you wanted to test a clear solution for the presence of

lead, what could you add?

6 VOLTS

DC

power source

1 M zinc sulfatesolution

beaker

stainlesssteel

copper

wire

Fig 2.3.5

Fig 2.4.2 Strong acids break apart completely in water, while weak acidstend to stay together.

49 Chapter 2 Chemical reactions

Many familiar substances can be classifiedas acids or bases. You can probably evenname a few, like the acid in citrus fruits(citric acid) or the acid in vinegar (ethanoicor acetic acid). Can you name any bases?You’ve no doubt heardof ammonia, which is inmany cleaning products.Did you know it was abase? So what makes asubstance acidic or basic,anyway?

AcidsnAcids contain the elementhydrogen in combination withother non-metal elements. Forexample, hydrochloric acid hasthe formula HCl. It containshydrogen in combination with

chlorine. Sulfuric acid, H2SO4, contains hydrogen incombination with the sulfate ion, which is made up of onesulfur and four oxygen atoms. When an acid is placed inwater, the hydrogen breaks away from the other elements.

With strong acids, the hydrogen breaks away veryeasily. Weak acids tend to hold onto their hydrogen, andvery little hydrogen breaks away. Strong acids arecorrosive. This means they will destroy living tissue and‘eat through’ some surfaces. Hydrochloric acid, whichhelps digestion of food in our stomachs, is a strong acid.The reason it doesn’t eat through the stomach lining isbecause the lining secretes a protective mucus. Otherexamples of strong acids are nitric acid (HNO3) andsulfuric acid. Weak acids include the ammonium ion(NH4

+) and citric acid (C6H8O7).Some properties of acids are:

• They have a sour taste (don’t try this).• They turn blue litmus red. Litmus is an indicator. See

page 53 for information about indicators.• They conduct electricity in aqueous (water) solution.

Earlier in this chapter, a solution was defined as asolute plus a solvent. A dilute solution is one that has veryfew solute particles in a relatively large volume of solvent.Concentrated solutions have many solute particlesdissolved in the solvent. Concentrated acid solutions arevery dangerous, while very dilute ones may be harmlessenough to drink. The acids that we eat and drink, like

Unit 2.4 Acids and bases

Fig 2.4.1 All of these common substances contain acids and bases.

Acid burnsSulfuric acid is a very

good dehydrating agent;

it removes water from

substances very easily.

That is why it chars

paper if they come into

contact. It is also this

property that makes

sulfuric acid sodamaging to living

tissue. It used to be

known as oil of vitriol,

and was sometimes

used by jealous women

to destroy the beauty of

their competitors for a

man’s affections.

Sciencesnippet

H+H+

H+ CH3COOH

H+

H+

H+

H+

Cl–Cl–

Cl–

Cl–Cl–

Cl–

CH3COO–

CH3COOH

CH3COOH

CH3COOH

CH3COOH

weak acid—CH3COOHacetic acid

strong acid—HCl hydrochloric acid

a strong base, because it reacts with oils to form soap,which then washes away easily.

Bases:• taste bitter• have a soapy feel (but remember, it is not advisable

to touch chemicals)• turn red litmus blue

Bases include ionic compounds like hydroxides,oxides, carbonates and hydrogen carbonates.

The table below shows some bases and their uses.

Base Common name Common use

Sodium hydroxide Caustic soda Making soaps

Calcium hydroxide Slaked lime Reducing acidity in soil

Ammonium hydroxide ‘Cleaning’ ammonia Cleaning products

Sodium hydrogen Baking soda, Cooking—makes cakescarbonate bicarbonate of soda rise

Sodium carbonate Washing soda, soda ash Washing powders

Acids and metalsWhen an acid reacts with a metal, a salt and hydrogengas are produced. A salt is an ionic compoundcontaining the ions left over after reaction. The generalreaction can be written as:

acid + metal → salt + hydrogene.g. nitric acid + zinc → zinc nitrate + hydrogen2HNO3(aq) + Zn(s) → Zn(NO3)2(aq) + H2(g)

Note that a subscript ‘(g)’ is used to show a gas.Most metals will react with acids. Some, like the

Group I metals, will react violently with even cold,dilute acid. Other metals, like lead, need hotter or moreconcentrated acid solutions to make them react.

The table below shows the reactions between someacids and metals.

SCI3 50

citric acid in oranges and sherbet, and lactic acid inyoghurt, are both weak and dilute, so don’t harm us.They are still strong enough to affect sensitive tissuesthough, like lemon juice on a cut or in the eye—ouch!

The table below shows some acids and the commonuses of each.

Acid Common name Common use

Acetylsalicylic acid Aspirin Pain reliever

Benzoic acid Sorbic acid Preservatives in foods

Ascorbic acid Vitamin C Vitamin supplement,antioxidant

Sulfuric acid Battery acid Car batteries,manufacturing fertilisers

4-chloro-2- MCPA Herbicidemethylphenoxyacetic acid

Hydrochloric acid Spirit of salts Brick cleaners, cleaningmetals

Ethanoic (acetic) acid Vinegar Flavour and preservingfood

BasesYou can think of bases as the chemical opposites ofacids. They neutralise acids. Bases react with acids to

produce water and othersubstances. Any reaction of anacid with a base is calledneutralisation. Strong bases, likestrong acids, attack living tissueand cause serious burns. The waythey react to skin is different toacids, so while strong acids arecorrosive, we say that strong basesare caustic. Bases thatdissolve in water are calledalkalis. Many householdcleaners contain basesbecause they are excellentat dissolving grease. Ovencleaners usually containsodium hydroxide (NaOH),

Sour wineIn addition to its mainingredients of water,ethanol (alcohol),sugars, tannins andadditives, wine alsocontains a variety ofacids. These includetartaric, malic, lacticand succinic acids. Theacidity level has to becontrolled carefully.Too much and the wineacquires a nasty, sourtaste. Too little and thewine will easily go ‘off’.

Sciencesnippet

Acid Metal Reaction equation Salt produced

Nitric acid Calcium 2HNO3(aq) + Ca(s) → H2(g) + Ca(NO3)2(aq) Calcium nitrate

Sulfuric acid Magnesium H2SO4(aq) + Mg(s) → H2(g) + MgSO4(aq) Magnesium sulfate

Hydrochloric Iron 2HCl(aq) + Fe(s) → H2(g) + FeCl2(aq) Iron(II) chlorideacid

You can test for the hydrogen given off using the‘pop’ test. A spark in the presence of H2 causes apopping sound as the gas combines with O2 to formwater.

NeutralisationAs stated earlier, a neutralisation reaction is when anacid reacts with a base. Water is always a product inneutralisation reactions, as is a salt:

acid + base → salt + waterWhen the base used is a carbonate or a hydrogen

carbonate, a third product, carbon dioxide, is observedas bubbles in the solution. Neutralisation reactions arevery common. Every time we brush our teeth, thetoothpaste, which contains a base, neutralises thedamaging acids left on our teeth by bacteria. Farmerscan reduce the effects of acid rain on soil by adding thebase calcium hydroxide. Indigestion caused by toomuch acid in the stomach can be fixed with antacids,which are just bases in solid or liquid form.

Acids and hydroxidesThe general reaction equation for an acid combiningwith a hydroxide is:

acid + hydroxide → salt + waterExamples of dilute acids reacting with dilute

hydroxide solutions are:• hydrochloric acid + sodium hydroxide → sodium

chloride + waterHCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

• nitric acid + lithium hydroxide → lithium nitrate +waterHNO3(aq) + LiOH(aq) → LiNO3(aq) + H2O(l)

The above acids react in the same way with solidhydroxides. The only difference is that the ‘(aq)’subscript next to the hydroxide would be replaced by‘(s)’.

The subscript ‘(l)’ is always used for water to show itis a liquid.

51 Chapter 2 Chemical reactions

U N I T 2 . 4

Prac 1p. 55

Acids, bases and astronauts

The astronauts of the Apollo 13 space mission faced, among other

things, a serious build-up of carbon dioxide on board their damaged

craft. Unless a way could be found to reduce the CO2 levels, the

astronauts would quickly have asphyxiated. By using lithium hydroxide

containers, they were able to keep the air breathable. Lithium

hydroxide reacts with carbon dioxide, producing lithium carbonate

and water.

Science snippet

Fig 2.4.4 Staff at Mission Control discuss the use of lithium hydroxide

canisters to reduce the dangerous levels of carbon dioxide

aboard the crippled Apollo 13 spacecraft, 15 April, 1970.

Fig 2.4.3 Hydrochloric acid is neutralised by sodium hydroxide, a strong base.

H+

H+

H+

H+

Cl–

Cl–Cl–

Cl– Cl–

Cl–Cl–

Cl–Na+

Na+

Na+Na+

Na+

Na+

Na+

Na+

OH–

OH–

OH–

OH–

H2O (l) + NaCl (aq)NaOH (aq)HCl (aq) +

waterH20

waterH20

Acids and oxidesThis general reaction of an acid with an ionic oxide is:

acid + oxide → salt + waterExamples of dilute acids reacting with solid oxides

are:• hydrochloric acid + calcium oxide → calcium

chloride + water2HCl(aq) + CaO(s) → CaCl2(aq) + H2O(l)

• sulfuric acid + lithium oxide → lithium sulfate +waterH2SO4(aq) + Li2O(s) → Li2SO4(aq) + H2O(l)

Acids and carbonatesLike the last two neutralisation reactions that we’velooked at, the reaction of an acid with a carbonateproduces a salt and water. It also produces a thirdproduct, carbon dioxide. The reaction of an acid with ahydrogen carbonate produces the same three things:• acid + carbonate → salt + water + carbon dioxide• acid + hydrogen carbonate →

salt + water + carbon dioxideExamples are:

• nitric acid + sodium carbonate → sodium nitrate +water + carbon dioxide2HNO3(aq) + Na2CO3(s) → 2NaNO3(aq) + H2O(l) + CO2(g)

• hydrochloric acid + ammonium carbonate →ammonium chloride + water + carbon dioxide2HCl(aq) + (NH4)2CO3(s) → 2NH4Cl(aq) + H2O(l) +CO2(g)

• sulfuric acid + copper(I) hydrogen carbonate →copper(I) sulfate + water + carbon dioxideH2SO4(aq) + 2CuHCO3(s) → Cu2SO4(aq) + H2O(l) +CO2(g)

It is the carbon dioxide produced that can make youburp after having antacids. Baking powder is a mix oftartaric acid and sodium hydrogen carbonate. Whenwater is added, they dissolve and mix. This producescarbon dioxide which makes cakes rise.

You can test for carbon dioxide by bubblingthe gas through limewater. The limewaterchanges from clear to milky if carbon dioxideis present. Another test is that a lit match willgo out in the presence of carbon dioxide.

SCI3 52

Fig 2.4.5 The pain of a bee sting is produced by the methanoic (formic)acid the bee injects

Things that stingAs everyone knows, bees can give very nasty stings. The painful sting isproduced by the methanoic (formic) acid they inject. This is the sameacid that puts the sting into stinging ants and stinging nettles. If youget stung by a stinging nettle, a dock leaf can give some relief. This isbecause the dock leaf contains a base that neutralises the acid fromthe nettle. Other stinging creatures, like wasps and some jellyfish,inject a base into the skin of their victims. This can be neutralised bywashing the wound with vinegar.

Science snippet

Fig 2.4.6 Limewater is a solution of calcium hydroxide that turns cloudyin the presence of carbon dioxide, due to the formation of acalcium carbonate precipitate

Prac 2p. 55

limewater

calciumcarbonatehydrochloric

acid

WS 2.2

Fig 2.4.7 The pH scale

53 Chapter 2 Chemical reactions

U N I T 2 . 4

The pH scaleDissolving a substance in water produces an aqueoussolution. To describe how strongly acidic or basic asolution is, we use the pH scale (pH is short for ‘powerof hydrogen’). At 25°C, the pH scale goes from 0 to 14,with acidic solutions having pH less than 7, and basicsolutions having pH greater than 7. Neutral solutionshave a pH equal to 7. Strongly acidic solutions have pHcloser to 0. Very basic solutions have pH closer to 14.Basic solutions are said to be alkaline. The pH is ameasure of how much free hydrogen is present in asolution. If there is a lot, the pH is very low. If there ishardly any, the pH is higher. Every time you take a stepalong the pH scale, say from pH 3 to pH 4, thehydrogen present decreases by a factor of 10. Say youhave 10 mL of a solution with a pH of 1. If you add 90 mL of water, the new volume is 100 mL and youhave diluted the solution by a factor of 10. The pH ofthe new solution will be 2. Note that very strong basicsolutions may have pH greater than 14, and very strongacid solutions may have pH less than zero.

IndicatorsIndicators are chemicals that may be used to show thepH of a particular solution. They are very useful in awide variety of chemical analyses. Some indicators cangive a good estimation of pH. Universal indicator is anexample of this because it can undergo many colourchanges. It is a mixture of dyes. Other indicators are lessprecise. Litmus, which is made from plants calledlichens, is red in acidic solution and blue in alkalinesolution. This doesn’t really allow for a precise pH valueto be assigned, but can help distinguish between an acidand a base. Figure 2.4.8 shows the colour changes ofsome common indicators.

Many plants, like beetroot, red cabbage, hydrangeasand hibiscus, produce dyes that can be used asindicators. For example, hydrangeas have blue flowersin acidic soil and pink flowers in alkaline soil.

Fig 2.4.9 Colour changes of universal indicator with pH

Fig 2.4.8 The colour changes of some indicators

Prac 3p. 56

Prac 4p. 56

Prac 5p. 57

Prac 6p. 57

strong acids neutral strong bases0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

pH

dish

was

hing

pow

der

dete

rgen

ts

baki

ng so

dase

a w

ater

bloo

dpu

re w

ater

tap

wat

er

coffe

ew

ine

oran

ge ju

ice

vine

gar

stom

ach

acid

bromothymolblue

litmus

methyl orange

phenolphthalein

universalindicator

colourless change pink

red-orangechange yellow

red change blue

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

deepviolet

violetbluegreenyelloworangereddeep red

pHindicator

yellow change blue

deep red red red orangeyellow

greengreen

green blue blue violet red orange blue violet

1 2 3 4 5 6 7 8 9 10 11 12 13 14pH

1 Name three fruits containing citric acid.2 What is an acid?3 What is a base?4 Give two properties of an acid and two properties of a

base.5 A substance X reacts with hydrochloric acid, producing

water and sodium chloride. What is the chemical nameof X?

6 What is the difference between a dilute solution of nitricacid and a concentrated solution of nitric acid?

7 How does sulfuric acid react with living tissue?8 When sodium hydroxide reacts with fats, soap is

produced. Why do bases feel soapy to touch?9 What is a salt?

10 What is an alkali?11 Write word equations for the reactions of the following

metals with nitric acid:a Aluminiumb Zincc Irond Lithium

12 How could you test for:a hydrogen gas?b carbon dioxide gas?

13 What salt would be produced by the followingneutralisation reactions?a nitric acid + strontium hydroxideb sulfuric acid + copper carbonatec hydrochloric acid + silver oxided nitric acid + magnesium hydrogen carbonate

14 What is the approximate pH of:a a strong acid?b a weak acid?c pure water?d a weak base?e a strong base?

15 A certain food is found to be slightly acidic. It containseither hydrochloric acid or acetic acid. Which do youthink is more likely and why?

16 At pH 8, what colour is:a universal indicator?b red litmus?c blue litmus?

SCI3 54

17 Does a solution at pH 7 contain more or less freehydrogen than a solution at pH 4?

18 Write word equations (and formula equations if youthink you can) for the following reactions:a hydrochloric acid + iron(II) hydrogen carbonateb nitric acid + silver hydroxidec sulfuric acid + barium oxide

19 What acid and base could you combine to make thesesalts?a Barium chlorideb Calcium nitratec Iron(III) sulfate

20 The normal pH in the mouth is about 6.5. The pH in thestomach is around 2 to 3. Why do you get a burningsensation in the oesophagus, throat and mouth whenyou throw up?

21 Azaleas grow only in soil with pH less than 7. Whatsubstance could you add to basic soil to lower its pH—water, an acid or a base? Explain.

22 You are given 10 mL each of two solutions. Solution Ahas a pH of 2. Solution B has a pH of 4. How muchwater would you have to add to solution A to make itspH the same as that of solution B? This is a hard one!

1 Many advertisements for shampoos and skin lotionsmention their pH. Which pH is best for your skin andwhich is best for your hair? Does the age of a person,and their type of skin or hair change the answer to thisquestion?

2 Sulfuric acid is one of the most important chemicals inthe world. The sulfuric acid production of a country issaid to be a good indicator of the state of its economy.Why?

3 Find some acids and bases not mentioned in thischapter that are used either in cooking or medicine.

4 The ideal pH of a swimming pool is around 7.2. At thispH, most bacteria and green algae can’t grow. Describehow pH levels of pools are tested and how the pH iskept at 7.2.

U n i t 2 . 4 Questions

U n i t 2 . 4 Research /Extension

WS 2.3

Unit 2.4 Prac 1Acids and metals

You will need5 test tubes with stoppers, test-tube rack, matches, 100 mLbeaker, small pieces of aluminium, magnesium, zinc, ironand tin, 0.1 M solutions of hydrochloric, sulfuric and aceticacids, lab coat, safety glasses

What to do1 Copy the following results table into your book:

Hydrochloric acid Sulfuric acid Acetic(ethanoic) acid

Aluminium

Magnesium

Zinc

Iron

Tin

2 Pour 2 cm of hydrochloric acid into each test tube.3 To the first test tube add one of the metals. If there is an

obvious reaction, hold the stopper on the tube for about15 seconds. Light a match and, removing the stopperquickly, hold the lit match to the mouth of the tube.Record your observations.

55 Chapter 2 Chemical reactions

U N I T 2 . 44 Repeat step 3 for the other metals. You do not have to

repeat the gas test for every reaction.5 Repeat steps 2–4 for the other acids.

Questions1 Write word equations for the reactions of the metals with

one of the acids tested.2 For the other two acids, name the salts produced in

each reaction.3 From the speed of the reaction with each metal, arrange

the metals tested in order from most active to leastactive.

Unit 2.4 Prac 2Acids and metalcarbonates

You will need4 test tubes, test-tube rack, stopper, 100 mL beaker, matches,limewater, solid samples of sodium hydrogen carbonate,lithium carbonate, sodium carbonate, ammonium carbonate,spatula, 0.1 M solutions of nitric and hydrochloric acids, lab coat, safety glasses

What to do1 You will be combining each acid with each solid.

Draw a suitable results table, similar to that used in Prac 1.

2 Add a small amount of each solid (about the tip of aspatula full) to four different test tubes.

3 Add 2 cm of nitric acid to the first tube and quicklystopper. Light a match and, removing the stopper

U n i t 2 . 4 Practicalactivities

Hold stopperon tube for15 seconds.

A ‘pop’indicateshydrogenis present.

A second person lightsa match and holds itto the mouth of the tube as the stopper isremoved.

‘pop’

i ii iii

Fig 2.4.10

matchstopper

Test 1If carbon dioxide is present, a lit match goes out.

Test 2If carbon dioxide is present, limewatergoes fromclear tomilky.

Fig 2.4.11

quickly, put the lit match in the mouth of the tube. Recordyour observations.

4 Add 2 cm of nitric acid to the second tube and quicklystopper. Remove the stopper and add a small amount oflimewater. Restopper the tube, but don’t let too muchgas build up. Record your observations.

5 Add nitric acid to the other two tubes, but don’t stopper.Record your observations.

6 Repeat steps 2–5 using hydrochloric acid.

Questions1 Write word equations for all reactions.2 Circle the salt in each equation.3 How could you set this experiment up so that the gas

produced bubbles through a separate beaker oflimewater as it is produced? Draw a diagram.

Unit 2.4 Prac 3Common indicators

You will need0.1 M solutions of sodium hydroxide and hydrochloric acid,distilled water, 3 test tubes, test-tube rack, 3 × 100 mLbeakers, liquid red and blue litmus, universal indicator,methyl orange, methyl red, bromothymol blue,phenolphthalein, lab coat, safety glasses

What to do1 Copy the following results table into your books:

Colour in Colour in Colour instrong acid strong base neutral solution

(hydrochloric (sodium (water)acid) hydroxide)

Red litmus

Blue litmus

Universal indicator

Methyl orange

Methyl red

Phenolphthalein

2 Using the beakers, pour 2 cm of acid into one test tube,2 cm of sodium hydroxide (base) into another test tube,and 2 cm of distilled water into the third test tube.

3 Add 3 drops of red litmus to each tube. Record yourresults.

4 Repeat steps 2 and 3 for the other indicators.

SCI3 56

Questions1 Why was distilled water used for this experiment, rather

than tap water?2 Summarise your results in a paragraph.

Unit 2.4 Prac 4Natural indicators

You will need0.1 M hydrochloric acid, 0.1 M sodium hydroxide, distilledwater, pink or red flower petals, beetroot juice, tea bag, 3 ×100 mL beakers, filter paper cut into strips, Bunsen burner,heat mat, tripod, gauze, matches, 3 test tubes, test-tube rack,Pasteur pipettes, lab coat, safety glasses

What to do1 Gently boil 50 mL of water in a beaker, combined with

the flower petals until the water becomes stronglycoloured, then remove from heat.

2 Place 20 mL of beetroot juice in another beaker.3 Boil 50 mL of water in another beaker and add a tea

bag.4 Using tongs, dip a strip of filter paper into each solution

and lay them on paper towel to dry.5 Place 2 cm of hydrochloric acid in a test tube, 2 cm of

sodium hydroxide in a second test tube and 2 cm ofdistilled water in a third test tube.

6 Add about 10 drops of flower petal water to each andrecord the colour of each solution.

7 Clean the test tubes and repeat steps 5 and 6, first usingbeetroot juice, then tea.

beaker

distilledwater

sodiumchloride

hydrochloricacid

test-tuberack

indicator

making flower petal indicator

heat-proof mat

Bunsenburner

tripod

50 mL water

red flowerpetals pipette

Fig 2.4.12

8 Carefully dry the freshly made indicator paper over aBunsen burner flame, being careful not to burn it.

9 When dry, put a drop of acid on one end of each pieceof paper, and a drop of base on the other end. Allowthem to dry, then stick them in your book.

Questions1 Write a suitable conclusion for this experiment.2 Which of the three indicators was best and why?3 How did the colours seen with the paper compare to the

colours seen when using the liquid indicators?4 Can you think of any other substances that might be

natural indicators? If so, explain why you think theywould work.

Unit 2.4 Prac 5Universal indicator

You will need0.1 M hydrochloric acid, 1.0 M sodium hydroxide, distilledwater, Pasteur pipette, 10 mL measuring cylinder, waterproofTexta, 14 large test tubes, universal indicator, lab coat,safety glasses

What to do1 Label your test tubes 1 to 14.2 Place 10 mL of 0.1 M hydrochloric acid in test tube

number 1.3 Using a pipette, transfer 1 mL of this solution to the

measuring cylinder. Add 9 mL of distilled water andpour the mixture into test tube number 2.

4 Using a pipette, transfer 1 mL of the solution in tube 2 tothe measuring cylinder. Add 9 mL of water and pour themixture into test tube number 4. Continue this method upto tube number 6.

5 Add 10 mL of distilled water to tube number 7.6 Add 10 mL of 1.0 M sodium hydroxide to tube number

14. 7 Using a pipette, remove 1 mL of this solution, add 9 mL

of water and pour the mixture into tube 13. Continuethis method down to tube number 8.

8 Add 3 drops of indicator to each tube and sketch theresult. The number of the tube is approximately the sameas the pH.

Questions1 On your sketch, show which tubes contain strong acids,

weak acids, strong bases and weak bases.

57 Chapter 2 Chemical reactions

U N I T 2 . 42 What is the dilution factor between:

a tubes 2 and 4?b tubes 10 and 11?

Unit 2.4 Prac 6Testing householdsolutions

You will need2 test tubes, test-tube rack, Pasteur pipettes, 2 watch-glasses,blue and red litmus paper, liquid universal indicator, distilledwater, safety glasses, lab coat, a variety of householdsolutions including orange juice, soft drink, fresh and sourmilk, vinegar. Solids may be used if dissolved in water first.

What to do1 Place 2 cm of solution into a test tube using a Pasteur

pipette.2 If the colour of the solution is quite strong, add distilled

water until it is faint.3 Pipette a small amount of the solution onto each of two

watch-glasses. Add red litmus paper to one and bluelitmus paper to the other. Record your results.

4 Add 3 drops of universal indicator to the test tube andrecord the pH of the solution.

5 Clean the equipment and repeat the procedure for yourother solutions.

Questions1 Arrange your solutions in a list from most acidic to least

acidic.2 A brick cleaner is marked as highly corrosive. Where do

you think it would go on your list?3 Explain the difference in pH between fresh and sour milk.

test solution + blue litmus paper

test tubestest-tuberack

testsolution

test solution +universal indicator

Pasteurpipette

red bluewatch-glass

test solution +red litmus paper

Fig 2.4.13

Unit 2.1 What are chemical reactions and why do they happen?

Word Clue

1 r___ ___ ___ ___ ___ ___ts Substances present at the start of a reaction.

2 ___ ___ ___ ___ ucts Substances present at the end of a chemical reaction.

3 c___ ___ m___ ___ ___l r___ ___ ___ti___ ___ Results in at least one new substance.

4 ___ ___ ___si___ ___l ch___ ___ ___ ___ Does not result in new substances being formed.

1 What observations might you make if a chemical reactionoccurs?

2 Is rain a physical change or a chemical change? Discuss.3 Is the melting of ice endothermic or exothermic?4 In sea water, what is the main solute and what is the

solvent?5 Why are some metals more reactive than others?6 Write formulas for:

a lithium hydroxideb barium sulfatec aluminium bromide

7 What are the systematic names of:a H2S?b PF3?c SiO2?

8 Are the atoms in H2O covalently or ionically bonded?Explain.

9 What charge do the ions of these metals have?a Sodiumb Strontiumc Aluminium

10 What are the names of these ions?a HCO3

–

b I–

c S2–

d NH4+

11 What does ‘diatomic’ mean? Give an example to illustrateyour answer.

12 What types of reactions are these?

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a lithium + chlorine → lithium chlorideb sulfuric acid + barium carbonate → barium sulfate +

carbon dioxide + water13 What sort of reaction could be used to coat a metal?14 What chemical is found in antacids?15 Are the following substances acids or bases?

a NaOHb Li2CO3

c HCld MgOe HNO3

16 What would be the products of the followingneutralisation reactions?a sodium carbonate + hydrochloric acidb calcium hydroxide + nitric acid

17 A clumsy student spills sulfuric acid on the lab floor.What chemical could you add to neutralise the acid?

18 What colour would the following indicators be at pH 4?a Blue litmusb Red litmusc Universal indicatord Methyl orange

19 Name an acid and a base found in the household.20 You are given three colourless solutions. One is pure

water, one is a solution of hydrochloric acid, one is asolution of sodium hydroxide. You are also given someuniversal indicator which you can add to only one. Howcould you determine which is which?

Chapter review questions

Sci-words

59 Chapter 2 Chemical reactions

5 g___ ___ Sign of chemical change.

6 e___ ___ ___t___ ___ ___ ___ic Reaction that absorbs energy.

7 ___ ___ ___ ___tion Solute + solvent = __________

8 re___ ___ ___ ___ve Metals that easily lose electrons.

9 p___ ___ ___ ___ ___it___ ___ ___ Causes cloudiness in solutions.

Unit 2.2 Naming compounds

Word Clue

1 s___ ___ ___ Another name for an ionic compound.

2 ___ ___ ___ic com___ ___ ___ ___ ___ Made up of positive and negative ions.

3 el___ ___ ___ros___ ___ ___ ___c Attraction of opposite charges.

4 v___ ___ ___ ___ ___y Charge on ion.

5 n___ ___-m___ ___ ___ ___ ___ Always form negative ions.

6 ___ ___ ___a___ ___nt Bonding in which electrons are shared.

7 r ___ ___ ___ ___ ___ ___ Another name for a polyatomic ion.

Unit 2.3 Reaction types

Word Clue

1 c___ ___ ___in___ ___ ___on Reaction which always gives one product.

2 d___ ___ ___ ___mic Molecule made up of two atoms.

3 ___ ___ ___ ___m___ ___ ___ ___ ___ ___ ___n Reaction in which one substance breaks down.

4 p___ ___ ___ ___ ___ita___ ___ ___ ___ Insoluble salt forming in solution.

5 neu___ ___ ___ ___ ___sa___ ___ ___ ___ Water is always a product of this reaction.

6 c___ ___ ___ ___ ___ ___ ___ ___n Oxygen is always a reactant in this reaction.

7___ ___ ___plac___ ___ ___nt Causes one metal to deposit on another.

Unit 2.4 Acids and bases

Word Clue

1 a___ ___ ___ Produces hydrogen ions in solution.

2 ___ ___ ___e Neutralises an acid.

3 c___ ___ ___ ___ s___ ___ ___ Strong acids have this property.

4 ___ ___ ___ ___ ___ ic Strong bases have this property.

5 i___ ___ ___ca___ ___ ___ Shows the pH of a substance.

6 w___ ___ ___ ___ Always a product in neutralisation reactions.