mixtures
Criteria for purity
Drugs are manufactured to a very high degree of
purity (Figure C2.5). To ensure that the highest possible
purity is obtained, the drugs are dissolved in a suitable
solvent and subjected to fractional crystallisation.
It is illegal to put anything harmful into food. Also,
government legislation requires that a lot of testing
takes place before a new pharmaceutical is marketed.
Figure C2.6 These pharmaceuticals must have been through a lot of
testing before they can be sold in a chemist’s shop.
Figure C2.5 Drugs are manufactured to a high degree of purity by Question
fractional crystallisation.
3 Describe how you would use chromatography to show
whether blue ink contains a single pure dye or a mixture
of dyes.
Throughout the chemical, pharmaceutical and food ●● Mixtures
industries it is essential that the substances used
are pure. Any impurities may interfere chemically, Many everyday things are not pure substances, they
physically or biologically with the behaviour of the are mixtures. A mixture contains more than one
product. The purity of a substance can be gauged by: substance (elements and/or compounds). An example
●● its melting point – if it is a pure solid it will have a of a common mixture is sea water (Figure C2.7).
sharp melting point. If an impurity is present then
melting takes place over a range of temperatures.
●● its boiling point – if it is a pure liquid the
temperature will remain steady at its boiling point.
If the substance is impure then the mixture will
boil over a temperature range.
●● chromatography – if it is a pure substance it Figure C2.7 Sea water is a common mixture.
will produce only one well-defined spot on a
chromatogram. If impurities are present then
several spots will be seen on the chromatogram
(see Figure C2.8 on the next page).
●● Solubility – if it is a pure substance then all of the
substance will dissolve in a suitable solvent at a
given temperature. Any undissolved solid would be
impurity and would need to be filtered off.
187
C2 exPeRimental teChniques
Other mixtures include the air, which is a mixture
of elements such as oxygen, nitrogen and neon and
compounds such as carbon dioxide (see Chapter C11,
p.295), and alloys such as brass, which is a mixture of
copper and zinc (for a further discussion of alloys see
Chapter C10, p.289).
Question a chromatographic separation of black ink
4 What criteria can be used to test the purity of a substance?
●● Separating mixtures chromatography
paper
Many mixtures contain useful substances mixed
with unwanted material. In order to obtain these black ink
useful substances, chemists often have to separate spot
them from the impurities. Chemists have developed before
many different methods of separation, particularly
for separating compounds from complex mixtures. solvent soaks beaker
Which separation method they use depends on up the paper solvent
what is in the mixture and the properties of the
substances present. It also depends on whether during
the substances to be separated are solids, liquids
or gases. dyeR
Separating solid/solid mixtures dyeQ
dyeP
Chromatography
What happens if you have to separate two or more after
solids that are soluble? This type of problem is b the black ink separates into three dyes: P, Q and R
encountered when you have mixtures of coloured Figure C2.8
materials such as inks and dyes. A technique called
chromatography is widely used to separate these Numerical measurements (retardation factors)
materials so that they can be identified. known as Rf values can be obtained from
chromatograms. An Rf value is defined as the
There are several types of chromatography; ratio of the distance travelled by the solute (for
however, they all follow the same basic principles. example P, Q or R) to the distance travelled by
The simplest kind is paper chromatography. To the solvent.
separate the different-coloured dyes in a sample of
black ink, a spot of the ink is put on to a piece of
chromatography paper. This paper is then set in a
suitable solvent as shown in Figure C2.8.
As the solvent moves up the paper, the dyes are
carried with it and begin to separate. They separate
because the substances have different solubilities in
the solvent and are absorbed to different degrees
by the chromatography paper. As a result, they are
separated gradually as the solvent moves up the
paper. The chromatogram in Figure C2.8b shows
how the ink contains three dyes, P, Q and R.
188
Separating mixtures
Separating solid/liquid mixtures The filter paper contains holes that, although
too small to be seen, are large enough to allow
If a solid substance is added to a liquid it may the molecules of water through but not the sand
dissolve to form a solution. In this case the solid is particles. It acts like a sieve. The sand gets trapped
said to be soluble and is called the solute. The liquid in the filter paper and the water passes through it.
it has dissolved in is called the solvent. An example of The sand is called the residue and the water is called
this type of process is when sugar is added to tea or the filtrate.
coffee. What other examples can you think of where
this type of process takes place? Crystallisation
In many parts of the world, salt is obtained from
Sometimes the solid does not dissolve in the liquid. sea water on a vast scale. This is done by using
This solid is said to be insoluble. For example, tea the heat of the sun to evaporate the water to leave
leaves themselves do not dissolve in boiling water a saturated solution of salt known as brine. A
when tea is made from them, although the soluble saturated solution is defined as one that contains
materials from which tea is made are seen to dissolve as much solute as can be dissolved at a particular
from them. temperature. When the solution is saturated the salt
begins to crystallise, and it is removed using large
Filtration scoops (Figure C2.10).
When a cup of tea is poured through a tea strainer
you are carrying out a filtering process. Filtration
is a common separation technique used in chemistry
laboratories throughout the world. It is used when a
solid needs to be separated from a liquid. For example,
sand can be separated from a mixture with water by
filtering through filter paper as shown in Figure C2.9.
Figure C2.9 I t is important when filtering not to overfill the Figure C2.10 S alt is obtained in north-eastern Brazil by evaporation
filter paper. of sea water.
Simple distillation
If we want to obtain the solvent from a solution,
then the process of distillation can be carried out.
The apparatus used in this process is shown in
Figure C2.11 on the next page.
189
C2 Experimental techniques
thermometer Figure C2.12 T his plant produces large quantities of drinking water in
Saudi Arabia.
water Liebig
out condenser Separating liquid/liquid mixtures
flask cooling Liquids that mix with each other can be separated
solution water in by fractional distillation. The apparatus used for
this process is shown in the photo and diagram
heat in Figure C2.13, and could be used to separate a
mixture of ethanol and water.
distillate
Figure C2.11 Water can be obtained from salt water by distillation. Fractional distillation relies upon the liquids
having different boiling points. When an ethanol and
Water can be obtained from salt water using this water mixture is heated the vapours of ethanol and
method. The solution is heated in the flask until it water boil off at different temperatures and can be
boils. The steam rises into the Liebig condenser, condensed and collected separately.
where it condenses back into water. The salt is left
behind in the flask. In hot and arid countries such Ethanol boils at 78 °C whereas water boils at
as Saudi Arabia this sort of technique is used on a 100 °C. When the mixture is heated the vapour
much larger scale to obtain pure water for drinking produced is mainly ethanol with some steam.
(Figure C2.12). This process is carried out in a Because water has the higher boiling point of the
desalination plant. two, it condenses out from the mixture with ethanol.
This is what takes place in the fractionating column.
The water condenses and drips back into the flask
while the ethanol vapour moves up the column and
into the condenser, where it condenses into liquid
ethanol and is collected in the receiving flask as the
distillate. When all the ethanol has distilled over,
the temperature reading on the thermometer rises
steadily to 100 °C, showing that the steam is now
entering the condenser. At this point the receiver can
be changed and the condensing water can now be
collected.
190
thermometer separating mixtures
condenser cooling Questions
water in
5 What is the difference between simple distillation and
cooling fractional distillation?
water out
fractionating column 6 Devise a method for obtaining salt (sodium chloride) from
with short lengths sea water in the school laboratory.
of glass rod inside
(increases surface area) 7 Which method would you use to separate:
a coffee grains from coffee solution
flask liquid flask b petrol (b.pt. 30 °C) from a crude oil
heat mixture distillate c pieces of steel from engine oil
support d oxygen (b.pt. −183 °C) from liquid air?
Figure C2.13 Typical fractional distillation apparatus. Checklist
After studying Chapter C2 you should know and understand
the following.
● The accuracy of each measurement in experimental work
depends on the quality of the measuring apparatus (e.g. the
thermometer or electronic balance) and on the skill of the
scientists taking the measurement.
● Chromatography is a technique employed for the separation
of mixtures of dissolved substances.
● Crystallisation is the process of forming crystals from a
liquid.
● Distillate is the condensed vapour produced from a mixture
of liquids on distillation.
● Distillation is the process of boiling a liquid and then
condensing the vapour produced back into a liquid. It is
used to purify liquids and to separate mixtures of liquids.
● Evaporation occurs when a solution is heated the solvent
evaporates and leaves the solute behind.
● Filtrate is the liquid which passes through the filter paper
during filtration.
● Filtration is the process of separating a solid from a liquid
using a fine filter paper which does not allow the solid to
pass through.
● Fractional distillation is a distillation technique used to
separate a mixture of liquids that have different boiling
points.
● If the solute does not dissolve in the solvent it is said to be
insoluble.
● A mixture is a system of two or more substances that can be
separated by physical means.
● The residue is the solid left behind in the filter paper after
filtration has taken place.
● The Rf value is the ratio of the distance travelled by
the solute to the distance travelled by the solvent in
chromatography.
● A saturated solution is a solution which contains as much
dissolved solute as it can at a particular temperature.
● If the solute dissolves in the solvent it is said to be soluble.
● Solute is the substance that dissolves (disappears) in(to) the
solvent.
● Solution is the liquid formed when a substance (solute)
disappears (dissolves) into another substance (solvent).
● Solvent is the liquid that the solute has dissolved in.
191
C3 Atoms, elements and compounds
Combined Co-ordinated
● Identify physical and chemical changes, and understand the ● Identify physical and chemical changes, and understand the
differences between them differences between them
● Describe the differences between elements, mixtures and ● U nderstand that some chemical reactions can be reversed
compounds, and between metals and non-metals by changing the reaction conditions (limited to the effects
of heat and water on hydrated and anhydrous copper(ii)
● Define the terms solvent, solute, solution and concentration sulfate and cobalt(ii) chloride – concept of equilibrium is
● Describe the structure of an atom in terms of a central not required)
nucleus, containing protons and neutrons, and ‘shells’ ● Describe the differences between elements, mixtures and
of electrons compounds, and between metals and non-metals
● Describe the build-up of electrons in ‘shells’ and understand
the significance of the noble gas electronic structures and of ● Define the terms solvent, solute, solution and concentration
the outer electrons (the ideas of the distribution of electrons ● Describe the structure of an atom in terms of a central
in s and p orbitals and in d block elements are not required)
● State the charges and approximate relative masses of nucleus, containing protons and neutrons, and ‘shells’
protons, neutrons and electrons of electrons
● Define and use proton number (atomic number) as the ● Describe the build-up of electrons in ‘shells’ and understand
number of protons in the nucleus of an atom the significance of the noble gas electronic structures and of
● Define and use nucleon number (mass number) as the total outer electrons (the ideas of the distribution of electrons in s
number of protons and neutrons in the nucleus of an atom and p orbitals and in d block elements are not required)
● State the charges and approximate relative masses of
● Use proton number and the simple structure of atoms protons, neutrons and electrons
to explain the basis of the Periodic Table, with special ● Define and use proton number (atomic number) as the
reference to the elements of proton number 1 to 20 number of protons in the nucleus of an atom
● Define and use nucleon number (mass number) as the total
● Describe the formation of ions by electron loss or gain number of protons and neutrons in the nucleus of an atom
● Use dot-and-cross diagrams to describe the formation of ionic ● Use proton number and the simple structure of atoms to
explain the basis of the Periodic Table, with special reference
bonds between Group I and Group VII to the elements of proton number 1 to 20
● Define isotopes as atoms of the same element which have
● D escribe the formation of ionic bonds between metallic the same proton number but a different nucleon number
and non-metallic elements, to include the strong
attraction between ions because of their opposite ● U nderstand that isotopes have the same properties
electrical charges because they have the same number of electrons in their
outer shell
● Describe the lattice structure of ionic compounds as a
regular arrangement of alternating positive and negative ● Describe the formation of ions by electron loss or gain
ions, exemplified by the sodium chloride structure ● Use dot-and-cross diagrams to describe the formation of ionic
● State that non-metallic elements form simple molecules with bonds between Group I and Group VII
covalent bonds between atoms
● D escribe the formation of ionic bonds between metallic
● Describe the formation of single covalent bonds in H2, Cl2, and non-metallic elements, to include the strong
H2O, CH2, NH3 and HCl as the sharing of pairs of electrons attraction between ions because of their opposite
leading to the noble gas configuration, including the use of electrical charges
dot-and-cross diagrams
● Describe the lattice structure of ionic compounds as a
● U se and draw dot-and-cross diagrams to represent the regular arrangement of alternating positive and negative
bonding in the more complex covalent molecules such as ions, exemplified by the sodium chloride structure
N2, C2H4, CH3OH and CO2
● State that non-metallic elements form simple molecules with
● Describe the differences in volatility, solubility and electrical covalent bonds between atoms
conductivity between ionic and covalent compounds
● Describe the formation of single covalent bonds in H2, Cl2,
● E xplain the differences in melting point and boiling point H2O, CH4, NH3 and HCl as the sharing of pairs of electrons
of ionic and covalent compounds in terms of attractive leading to the noble gas configuration, including the use of
forces dot-and-cross diagrams
● Describe the differences in volatility, solubility and electrical
conductivity between ionic and covalent compounds
192
Physical and chemical changes
Co-ordinated (continued...) ● Describe the giant covalent structures of graphite and
diamond
● ●Use and draw dot-and-cross diagrams to represent the
bonding in the more complex covalent molecules such as ● R● elate the structures of diamond and graphite to their
N2, C2H4, CH3OH and CO2 uses, e.g. graphite as a lubricant and a conductor and
diamond in cutting tools
● ●Explain the differences in melting point and boiling point
of ionic and covalent compounds in terms of attractive ● ●Describe the macromolecular structure of silicon(IV) oxide
forces (silicon dioxide, SiO2)
● State that there are several different forms of carbon,
including diamond and graphite
●● Physical and chemical
changes
In Chapter C1 (p.181) we noted that when a
substance changes its physical state from, say, a solid
to a liquid then this is called a physical change. A
good example of this process takes place when water,
as ice, melts to form liquid water. In these changes no
new substances are produced.
Figure C3.2 This child is getting energy from the chemical changes that
take place in her body from eating the apple.
Reversible reactions
Many of these chemical reactions cannot be reversed
easily. However, there are some that can be reversed
easily. For example, when metal hydrates such as
copper(ii) sulfate pentahydrate (CuSO4.5H2O) are
heated the water of crystallisation is driven away.
Anhydrous copper(ii) sulfate remains. Also in this
case a colour change takes place from blue to white.
copper(ii) sulfate → anhydrous copper(ii) + water
pentahydrate sulfate
Figure C3.1 Magnesium burns brightly CuSO4.5H2O(s) → CuSO4(s) + 5H2O(g)
in oxygen to produce magnesium oxide. (blue) (white)
When a new substance is formed during a chemical When water is added to anhydrous copper(ii)
reaction, a chemical change has taken place. sulfate the reverse process occurs. It turns blue and
the pentahydrate is produced. This is an extremely
magnesium + oxygen → magnesium oxide exothermic process.
2Mg(s) + O2(g) → 2MgO(s) CuSO4(s) + 5H2O(g) → CuSO4.5H2O(s)
However, chemical changes do not just happen (white) (blue)
with reactions in test tubes. All living things need
chemical changes to stay alive. For example we These processes give a simple example of a reversible
eat lots of food (Figure C3.2) so that we can stay reaction:
healthy. The food has to be digested and this
involves chemical changes. CuSO4(s) + 5H2O(l) → CuSO4.5H2O(s)
193
C3 atOms, elements anD COmPOunDs
Questions classify them as metals or non-metals (Figures C3.3
and C3.4). Table C3.1 shows the physical data for
1 Define the following terms using specific examples to help some common metallic and non-metallic elements.
with your explanation:
a element You will notice that many metals have high densities,
b metal high melting points and high boiling points, and that
c non-metal. most non-metals have low densities, low melting points
and low boiling points. Table C3.2 summarises the
2 Which of the substances listed below are: different properties of metals and non-metals.
a metallic elements
b non-metallic elements? A discussion of the chemical properties of metals
silicon, chromium, chlorine, iron, copper, sulfur, nitrogen, zinc is given in Chapters C9 and C10. The chemical
properties of certain non-metals are discussed in
●● Elements Chapters C9, C11 and C12.
Robert Boyle used the name element for any Table C3.1 Physical data for some metallic and non-metallic
substance that cannot be broken down further, into a elements at room temperature and pressure
simpler substance. This definition can be extended to
include the fact that each element is made up of only Element Metal or Density/ Melting Boiling
one kind of atom. The word atom comes from the non-metal point/°C
Greek word atomos meaning ‘unsplittable’. g/cm3 point/°C
For example, aluminium is an element which is Aluminium Metal 2.70 660 2580
made up of only aluminium atoms. It is not possible
to obtain a simpler substance chemically from Copper Metal 8.92 1083 2567
the aluminium atoms. You can only make more
complicated substances from it, such as aluminium Gold Metal 19.29 1065 2807
oxide, aluminium nitrate or aluminium sulfate.
Iron Metal 7.87 1535 2750
There are 118 elements which have now been
identified. Twenty-five of these do not occur in Lead Metal 11.34 328 1740
nature and have been made artificially by scientists.
They include elements such as curium and Magnesium Metal 1.74 649 1107
unnilpentium. Ninety-one of the elements occur
naturally and range from some very reactive gases, Nickel Metal 8.90 1453 2732
such as fluorine and chlorine, to gold and platinum,
which are unreactive elements. Silver Metal 10.50 962 2212
All elements can be classified according to their Zinc Metal 7.14 420 907
various properties. A simple way to do this is to
Carbon Non-metal 2.25 Sublimes at 3642
Hydrogen Non-metal 0.07a −259 −253
Nitrogen Non-metal 0.88b −210 −196
Oxygen Non-metal 1.15c −218 −183
Sulfur Non-metal 2.07 113 445
Source: Earl B., Wilford L.D.R. Chemistry data book. Nelson Blackie,
1991 a At −254 °C b At −197 °C c At −184 °C.
b aluminium has many uses in the aerospace industry
a gold is very decorative c these coins contain nickel
Figure C3.3 Some metals.
194
Compounds
Table C3.2 How the properties of metals and non-metals compare
Property Metal Non-metal
Usually solid Solid, liquid or gas
Physical state at room (occasionally liquid)
temperature Good Poor – usually soft or
Good brittle
Malleability Shiny (lustrous) Dull
Usually high Usually low
Ductility Usually high Usually low
Usually high Usually low
Appearance (solids) Good Very poor
Melting point
Boiling point
Density
Conductivity (thermal
and electrical)
c neon is used in advertising signs
Figure C3.4 S ome non-metals.
●● Compounds
Compounds are pure substances which are formed when
two or more elements chemically combine together.
Water is a simple compound formed from the elements
hydrogen and oxygen (Figure C3.5). This combining of
the elements can be represented by a word equation:
hydrogen + oxygen → water
a a premature baby needs oxygen hydrogen oxygen hydrogen and water
b artists often use charcoal (carbon) to produce an initial sketch a pure a pure oxygen mixed a pure
element element together compound
formed from
hydrogen
burning in
oxygen
O H O
H
HH O HH O H
O
ϩ O
H O H
H O
HH O HH HH
H O O
H H H H
H H
O
H
Figure C3.5 The element hydrogen reacts with the element oxygen to
produce the compound water.
Water molecules contain two atoms of hydrogen
and one atom of oxygen, and hence water has
the chemical formula H2O. Elements other than
hydrogen will also react with oxygen to form
compounds called oxides. For example, magnesium
195
C3 atOms, elements anD COmPOunDs
reacts violently with oxygen gas to form the white During the reaction heat energy is given out as
powder magnesium oxide (Figure C3.1). This new chemical bonds are formed. This is called an
reaction is accompanied by a release of energy as new exothermic reaction and accompanies a chemical
chemical bonds are formed. change (Chapter C6, p.243). The iron(ii) sulfide
formed has totally different properties to the
What is the difference between mixture of iron and sulfur (Table C3.3). Iron(ii)
mixtures and compounds? sulfide, for example, would not be attracted
towards a magnet.
There are differences between compounds and
mixtures. This can be shown by considering the Table C3.3 The iron(II) sulfide formed has totally different
reaction between iron filings and sulfur. A mixture properties from the mixture of iron and sulfur
of iron filings and sulfur looks different from the
individual elements (Figure C3.6). This mixture has Substance Appearance Effect of a Effect of dilute
the properties of both iron and sulfur; for example, a magnet hydrochloric acid
magnet can be used to separate the iron filings from Iron Dark grey Attracted
the sulfur (Figure C3.7). powder to it Very little action when
Sulfur cold. When warm, a gas
Substances in a mixture have not undergone Yellow None is produced with a lot of
a chemical reaction and it is possible to separate Iron/sulfur powder bubbling (effervescence)
them provided that there is a suitable difference mixture Dirty yellow Iron powder
in their physical properties. If the mixture of iron Iron(ii) powder attracted to it No effect when hot
and sulfur is heated a chemical reaction occurs and sulfide Black solid No effect or cold
a new substance is formed called iron(ii) sulfide
(Figure C3.6). The word equation for this reaction is: Iron powder reacts
as above
A foul-smelling gas is
produced with some
effervescence
iron + sulfur heat iron(ii) sulfide Table C3.4 The major differences between mixtures and compounds
Mixture Compound
It contains two or more substances It is a single substance
The composition can vary The composition is always the same
No chemical change takes place When the new substance is formed
when a mixture is formed it involves chemical change
The properties are those of the The properties are very different to
individual elements/compounds those of the component elements
The components may be The components can only be
separated quite easily by physical separated by one or more chemical
means reactions
Figure C3.6 The elements sulfur and iron at the top of the photograph, In iron(ii) sulfide, FeS, one atom of iron has
and (below) black iron(ii) sulfide on the left and a mixture of the two combined with one atom of sulfur. No such ratio
elements on the right. exists in a mixture of iron and sulfur, because the
atoms have not chemically combined. Table C3.4
summarises how mixtures and compounds compare.
Questions
3 Make a list of some other common mixtures, stating what
they are mixtures of.
4 Make a list of some other common compounds, stating the
elements they contain and giving the chemical formulae
for them.
Figure C3.7 A magnet will separate the iron from the mixture.
196
Atomic structure
●● Solutions This concentration effect can be seen quite clearly
if we use copper(ii) sulfate solution. There is less
Some people cannot swallow tablets. To make copper(ii) sulfate in the solution present in Figure
medicines easier to swallow, many manufacturers now C3.11a (it’s a dilute solution) compared with that
produce tablets which can be dissolved in water, shown in Figure C3.11b (it’s more concentrated).
such as the soluble aspirin shown in Figure C3.8.
ab
Figure C3.8 T he soluble aspirin dissolves in the water Figure C3.11 a Left b eaker containing dilute copper(II) sulfate solution;
b right beaker containing more concentrated solution of copper(II) sulfate
When stirred in water the tablet dissolves, leaving no solution.
trace of solid. The solid aspirin has dissolved in the
water to form a solution. The solid chemicals in the Other solvents
tablets are the solute and the water is the solvent.
The solute dissolves in the solvent to form a solution. Water is not the only solvent which can be used
The process of dissolving involves the separation to dissolve substances. There are many solutions
of the added solute particles as the liquid particles which we use every day which use other solvents
move about and collide with them (Figures C3.9 (Figure C3.12).
and C3.10).
a Initially, the aspirin particles hold b Soon the aspirin particles are
one another together. Then the scattered in a random way
water particles begin to pull the through the water solvent. We
particles in the aspirin tablet apart. now have a solution. The aspirin
has dissolved in the water to form
an aspirin solution.
Figure C3.9 Particles in an aspirin Figure C3.10 A solution of ab
tablet surrounded by water molecules. aspirin in water.
Figure C3.12 a Correction fluid is a solution of a plastic dissolved in a
If more aspirin tablets were dissolved in the solution non-toxic, non-volatile hydrocarbon. b Aftershave lotion is a solution of
then the solution would become more concentrated. perfume dissolved in ethanol.
There would be more particles of aspirin present in
the same volume of water. ●● Atomic structure
We have already seen in Chapter C1 that everything
you see around you is made out of tiny particles
called atoms (Figure C3.13 on the next page). When
John Dalton developed his atomic theory, about
200 years ago (1807/1808), he stated that the atoms
of any one element were identical and that each atom
was ‘indivisible’. Scientists in those days believed that
atoms were solid particles like marbles.
197
C3 Atoms, elements and compounds
About 1837 electrons are equal in mass to the mass
of one proton or one neutron. A summary of each
type of particle, its mass and relative charge is shown
in Table C3.5. You will notice that the masses of all
these particles are measured in atomic mass units
(amu). This is because they are so light that their
masses cannot be measured usefully in grams.
Table C3.5 Characteristics of a proton, a neutron and an electron
Particle Symbol Relative mass/amu Relative charge
Proton p 1 +1
Neutron n 1 0
Electron e 1/1837 −1
Figure C3.13 A micrograph of atoms. Although atoms contain electrically charged particles,
the atoms themselves are electrically neutral (they
However, in the last hundred years or so it has been have no overall electric charge). This is because atoms
proved by great scientists, such as Niels Bohr, Albert contain equal numbers of electrons and protons. For
Einstein, Henry Moseley, Joseph Thomson, Ernest example, the diagram in Figure C3.15 represents the
Rutherford and James Chadwick, that atoms are in atom of the non-metallic element helium. The atom
fact made up of even smaller ‘sub-atomic’ particles. of helium possesses two protons, two neutrons and
The most important of these are electrons, protons two electrons. The electrical charge of the protons
and neutrons, although 70 sub-atomic particles have in the nucleus is, therefore, balanced by the opposite
now been discovered. charge of the two electrons.
●● Inside atoms nucleus which holds
the 2 protons and
The three sub-atomic particles are found in distinct 2 neutrons
and separate regions. The protons and neutrons are
found in the centre of the atom, which is called the Ϫ
nucleus. The neutrons have no charge and protons
are positively charged. The nucleus occupies only a ϩ
very small volume of the atom but is very dense. ϩ
The rest of the atom surrounding the nucleus is Ϫ 2 electrons circling
where electrons are most likely to be found. The the nucleus
electrons are negatively charged and move around
very quickly in electron shells or energy levels. Figure C3.15 A n atom of helium has two protons, two electrons and
The electrons are held within the atom by an two neutrons.
electrostatic force of attraction between themselves
and the positive charge of protons in the nucleus Proton number and nucleon number
(Figure C3.14).
The number of protons in the nucleus of an atom
nucleus, containing is called the proton number (or atomic number)
neutrons and protons and is given the symbol Z. Hence in the diagram
shown in Figure C3.15, the helium atom has a
region where electrons proton number of 2, since it has two protons in its
are found nucleus. Each element has its own proton number
and no two different elements have the same proton
Figure C3.14 D iagram of an atom.
198
inside atoms
number. For example, a different element, lithium, The number of neutrons present can be calculated by
has a proton number of 3, since it has three protons rearranging the relationship between the proton number,
in its nucleus. nucleon number and number of neutrons to give:
Neutrons and protons have a similar mass. number of = nucleon number − proton number
Electrons possess very little mass. So the mass of neutrons (A) (Z )
any atom depends on the number of protons and
neutrons in its nucleus. The total number of protons For example, the number of neutrons in one atom of
and neutrons found in the nucleus of an atom is 2142Mg is:
called the nucleon number (or mass number) and is 24 − 12 = 12
given the symbol A. (A) (Z )
nucleon number = proton number + number of and the number of neutrons in one atom of 20872Pb is:
(A) (Z ) neutrons 207 − 82 = 125
(A) (Z)
Hence, in the example shown in Figure C3.15 the
helium atom has a nucleon number of 4, since it has Table C3.6 shows the number of protons, neutrons
two protons and two neutrons in its nucleus. If we and electrons in the atoms of some common elements.
consider the metallic element lithium, it has three
protons and four neutrons in its nucleus. It therefore Isotopes
has a nucleon number of 7.
Not all of the atoms in a sample of chlorine, for
The proton number and nucleon number of example, will be identical. Some atoms of the same
an element are usually written in the following element can contain different numbers of neutrons
shorthand way: and so have different nucleon numbers. Atoms of the
same element which have the same proton number but
nucleon number ((AZ ))←←42He ← symbol of the element different neutron numbers are called isotopes. The
proton number two isotopes of chlorine are shown in Figure C3.16 on
the next page.
Table C3.6 Number of protons, neutrons and electrons in some elements
Element Symbol Proton number Number of electrons Number of protons Number of neutrons Nucleon number
Hydrogen H 1 1 1 0 1
Helium He 2 2 2 2 4
Carbon C 6 6 6 6 12
Nitrogen N 7 7 7 7 14
Oxygen O 8 8 8 8 16
Fluorine F 9 9 9 10 19
Neon Ne 10 10 10 10 20
Sodium Na 11 11 11 12 23
Magnesium Mg 12 12 12 12 24
Sulfur S 16 16 16 16 32
Potassium K 19 19 19 20 39
Calcium Ca 20 20 20 20 40
Iron Fe 26 26 26 30 56
Zinc Zn 30 30 30 35 65
199
C3 atOms, elements anD COmPOunDs
3157 Cl 37 Cl
17 protons 17
18 neutrons 17 protons
Figure C3.16 The two isotopes of chlorine.
20 neutrons
Generally, isotopes behave in the same way during Figure C3.17 Cobalt-60 is used in radiotherapy treatment.
chemical reactions because they have the same
number of electrons on their outer shell (see p.201). Questions
5 Calculate the number of neutrons in the following atoms:
The only effect of the extra neutron is to alter the a 1237Al b 3115P c 120672Uus d 19706Os
mass of the atom and properties which depend on it,
such as density. Some other examples of atoms with 6 a Define the terms: proton, neutron and electron.
isotopes are shown in Table C3.7. b An atom X has a proton number of 19 and relative
There are two types of isotopes: those which are atomic mass of 39.
stable and those which are unstable. The isotopes i How many electrons, protons and neutrons are
which are unstable, as a result of the extra neutrons
in their nuclei, are radioactive and are called there in an atom of X?
radioisotopes. For example, uranium-235, which ii How many electrons will there be in the outer
is used as a source of power in nuclear reactors, and
cobalt-60, which is used in radiotherapy treatment energy level (shell) of an atom of X?
(Figure C3.17), are both radioisotopes. iii What is the electronic configuration of X?
7 60 Ga and 71 Ga are isotopes of gallium. With reference
31 31
to this example, explain what you understand by the
term isotope.
Table C3.7 Some atoms and their isotopes ●● The arrangement of
electrons in atoms
Element Symbol Particles present
Hydrogen The nucleus of an atom contains the heavier sub-
(Deuterium) 11H 1 e, 1 p, 0 n atomic particles – the protons and the neutrons. The
(Tritium) 21H 1 e, 1 p, 1 n electrons, the lightest of the sub-atomic particles,
Carbon 31H 1 e, 1 p, 2 n move around the nucleus at great distances from the
126C 6 e, 6 p, 6 n nucleus relative to their size. They move very fast in
Oxygen 163C 6 e, 6 p, 7 n electron energy levels very much as the planets orbit
146C 6 e, 6 p, 8 n the Sun.
Strontium 168O 8 e, 8 p, 8 n
178O 8 e, 8 p, 9 n It is not possible to give the exact position of an
Uranium 188O 8 e, 8 p, 10 n electron in an energy level. However, we can state
8368Sr 38 e, 38 p, 48 n that electrons can only occupy certain, definite energy
3888Sr 38 e, 38 p, 50 n levels and that they cannot exist between them. Each
9308Sr 38 e, 38 p, 52 n
23952U 92 e, 92 p, 143 n
23982U 92 e, 92 p, 146 n
200
of the electron energy levels can hold only a certain The arrangement of electrons in atoms
number of electrons.
l First energy level holds up to two electrons. HH 1p
l Second energy level holds up to eight electrons. 1p
l Third energy level holds up to 18 electrons.
There are further energy levels which contain hhyyddrrooggeenn
increasing numbers of electrons. LLii 3p
The third energy level can be occupied by a 43np
maximum of 18 electrons. However, when eight 4n
electrons have occupied this level a certain stability is
given to the atom and the next two electrons go into lliitthhiiuumm
the fourth energy level, and then the remaining ten
electrons complete the third energy level. NNaa11p
121np
The electrons fill the energy levels starting from 12n
the energy level nearest to the nucleus, which has the
lowest energy. When this is full (with two electrons) ssooddiiuumm
the next electron goes into the second energy level.
When this energy level is full with eight electrons, AArr18p
then the electrons begin to fill the third and fourth 2128np
energy levels as stated above. 22n
For example, a 168O atom has a proton number aarrggoonn
of 8 and therefore has eight electrons. Two of the
eight electrons enter the first energy level, leaving
six to occupy the second energy level, as shown in
Figure C3.18. The electron configuration for oxygen
can be written in a shorthand way as 2,6.
8p
8n
Figure C3.18 Arrangement of electrons in an oxygen atom. KK19p
2109np
There are 118 elements, and Table C3.8 on the 20n
next page shows the way in which the electrons are
arranged in the first 20 of these elements. The way ppoottaassssiiuumm
in which the electrons are distributed is called the Figure C3.19 Electron arrangements of hydrogen, lithium, sodium,
electronic structure (or electron configuration). argon and potassium.
Figure C3.19 shows the electronic structure of a
selection of atoms.
201
C3 atOms, elements anD COmPOunDs
Table C3.8 Electronic structure of the first 20 elements The outer electrons are mainly responsible for the
chemical properties of any element, and, therefore,
Element Symbol Proton Number of Electron elements in the same group have similar chemical
number electrons structure properties (Tables C.3.9, C.3.10 and C3.11).
Hydrogen H 1
Helium He 1 1 2 Table C3.9 Electronic structure of the first three elements of Group I
Lithium Li 2 2 2,1
Beryllium Be 3 3 2,2 Element Symbol Proton number Electronic structure
Boron B 4 4 2,3 Lithium Li 3 2,1
Carbon C 5 5 2,4 Sodium Na 2,8,1
Nitrogen N 6 6 2,5 Potassium K 11 2,8,8,1
Oxygen O 7 7 2,6 19
Fluorine F 8 8 2,7
Neon Ne 9 9 2,8 Table C3.10 Electronic structure of the first three elements of Group II
Sodium Na 10 10 2,8,1
Magnesium Mg 11 11 2,8,2 Element Symbol Proton number Electronic structure
Aluminium Al 12 12 2,8,3 Beryllium Be 4 2,2
Silicon Si 13 13 2,8,4 Magnesium Mg 12 2,8,2
Phosphorus P 14 14 2,8,5 Calcium Ca 20 2,8,8,2
Sulfur S 15 15 2,8,6
Chlorine Cl 16 16 2,8,7 Table C3.11 Electronic structure of the first three elements in Group VII
Argon Ar 17 17 2,8,8
Potassium K 18 18 2,8,8,1 Element Symbol Proton number Electronic structure
Calcium Ca 19 19 2,8,8,2 Fluorine F 9 2,7
20 20 Chlorine Cl 17 2,8,7
Bromine Br 2,8,18,7
35
You will notice from Table C3.8 that the elements The metallic character of the elements in a group
helium, neon and argon have completely full outer increases as you move down the group. This is
shells. In Chapter C9 you will see that these elements because electrons become easier to lose as the outer
are known as the noble or inert gases and that they shell electrons become further from the nucleus.
are generally very stable and unreactive (p.274–275). There is less attraction between the nucleus and
This is linked to the full outer shells that they possess. the outer shell electrons because of the increased
It would seem that when elements react to form distance between them.
compounds they do so to achieve full electron energy
levels. This idea forms the basis of the electronic Questions
theory of chemical bonding.
8 How many electrons may be accommodated in the first
●● Electronic structure and three energy levels?
the Periodic Table
9 What is the same about the electron structures of:
a lithium, sodium and potassium
b beryllium, magnesium and calcium?
The number of electrons in the outer energy ●● Ions and ionic bonding
level was discussed above. It can be seen that it
corresponds with the number of the group in the Ions
Periodic Table in which the element is found. For
example, the elements shown in Table C3.9 have An ion is an electrically charged particle. When
one electron in their outer energy level and they an atom loses one or more electrons it becomes
are all found in Group I. The elements in Group 0, a positively charged ion. For example, during the
however, are an exception to this rule, as they have chemical reactions of potassium, each atom loses an
two or eight electrons in their outer energy level. electron to form a positive ion, K+.
202
ions and ionic bonding
19 protons = 19+ sodium + chlorine → sodium chloride
19K+ 18 electrons = 18−
Overall charge = 1+
When an atom gains one or more electrons it
becomes a negatively charged ion. For example,
during some of the chemical reactions of oxygen it
gains two electrons to form a negative ion, O2−.
8 protons = 8+
8O2− 10 electrons = 10−
Overall charge = 2−
Table C3.12 shows some common ions. You will
notice from Table C3.12 that:
● some ions contain more than one type of atom, for
example NO3−
● an ion may possess more than one unit of charge
(either negative or positive), for example Al3+, O2−
or SO42−.
Table C3.12 Some common ions
Name Formula Figure C3.20 The properties of salt are very different from those of the
Lithium ion Li+ sodium and chlorine it was made from. To get your salt you would not
Sodium ion Na+ eat sodium or inhale chlorine!
Potassium ion K+
Magnesium ion Mg2+ Sodium has just one electron in its outer energy level
Calcium ion Ca2+ (11Na 2,8,1). Chlorine has seven electrons in its outer
Aluminium ion Al3+ energy level (17Cl 2,8,7). When these two elements react,
Zinc ion Zn2+ the outer electron of each sodium atom is transferred to
Ammonium ion NH4+ the outer energy level of a chlorine atom (Figure C3.21
Fluoride ion F− on the next page).
Chloride ion Cl−
Bromide ion Br− In this way both the atoms obtain full outer energy
Hydroxide ion OH− levels and become ‘like’ the nearest noble gas. The
Oxide ion O2− sodium atom has become a sodium ion and the
Sulfide ion S2− process is known as ionisation. This sodium ion has
Carbonate ion CO32− an electron configuration like neon.
Nitrate ion NO3−
Sulfate ion SO42−
Ionic bonding sodium atom → sodium ion + electron
Ionic bonds are usually found in compounds that Na(g) → Na+(g) + e−
contain metals combined with non-metals. When this
type of bond is formed, electrons are transferred from The chlorine atom has become a chloride ion with an
the metal atoms to the non-metal atoms during the electron configuration like argon.
chemical reaction. In doing this, the atoms become
more stable by getting full outer energy levels. For chlorine atom + electron → chloride ion
example, consider what happens when sodium and
chlorine react together and combine to make sodium Cl(g) + e− → Cl−(g)
chloride (Figure C3.20).
Only the outer electrons are important in bonding,
so we can simplify the diagrams by missing out the
inner energy levels (Figure C3.22 on the next page).
203
C3 Atoms, elements and compounds
The charges on the sodium and chloride ions are This type of bonding is called ionic bonding. The
equal but opposite. They balance each other and the alternative name, electrovalent bonding, is derived from
resulting formula for sodium chloride is NaCl. These the fact that there are electrical charges on the atoms
oppositely charged ions attract each other and are pulled, involved in the bonding.
or bonded, to one another by strong electrostatic forces.
Ϫ
ϩ
sodium atom chlorine atom sodium ion chloride ion
(Na) (Cl) (Naϩ) (ClϪ)
11 protons ϭ 11ϩ 17 protons ϭ 17ϩ
10 electrons ϭ 10Ϫ 18 electrons ϭ 18Ϫ
Overall charge ϭ ϩ1 Overall charge ϭ Ϫ1
Figure C3.21 Ionic bonding in sodium chloride.
Ϫ
Na Cl [Na] ϩ Cl
Figure C3.22 S implified diagram of ionic bonding in sodium chloride. sodium chloride (NaCl)
Mg O 2Ϫ
[Mg]2ϩ O
Figure C3.23 S implified diagram of ionic bonding in magnesium oxide. magnesium oxide (MgO)
Figure C3.23 shows the electron transfers that take Magnesium obtains a full outer energy level by
place between a magnesium atom and an oxygen losing two electrons. These are transferred to the
atom during the formation of magnesium oxide. oxygen atom. In magnesium oxide, the Mg2+ and
204
Ions and ionic bonding
O2− are oppositely charged and are attracted to one gives each of the two chlorine atoms one electron.
another. The formula for magnesium oxide is MgO. In this case, a compound is formed containing two
chloride ions (Cl−) for each calcium ion (Ca2+). The
Figure C3.24 shows the electron transfers that chemical formula is CaCl2.
take place during the formation of calcium chloride.
When these two elements react, the calcium atom Ϫ
Cl Cl
Ca [Ca]2ϩ
Ϫ
Cl Cl
calcium chloride (CaCl2 )
Figure C3.24 The transfer of electrons that takes place during the formation of calcium chloride.
Ionic structures Ionic structures are solids at room temperature
and have high melting and boiling points. The
Scientists, using a technique called X-ray diffraction, ions are packed together in a regular arrangement
have obtained photographs that indicate the called a lattice. Within the lattice, oppositely
way in which the ions are arranged in crystals of charged ions attract one another strongly.
ionic substance such as sodium chloride. Figure
C3.25 shows the structure of sodium chloride as Figure C3.25 shows only a tiny part of a small
determined by the X-ray diffraction technique. crystal of sodium chloride. Many millions of
sodium ions and chloride ions would be arranged
ClϪ in this way in a crystal of sodium chloride to make
Naϩ up the giant ionic structure. Each sodium ion in
the lattice is surrounded by six chloride ions, and
Naϩ surrounded ClϪ surrounded each chloride ion is surrounded by six sodium ions.
by 6 C lϪ ions by 6 Na ϩ ions
Not all ionic substances form the same structures.
Figure C3.25 The structure of sodium chloride. Caesium chloride (CsCl), for example, forms a different
structure due to the larger size of the caesium ion
compared with that of the sodium ion. This gives rise to
the structure shown in Figure C3.26 on the next page,
which is called a body-centred cubic structure. Each
caesium ion is surrounded by eight chloride ions and, in
turn, each chloride ion is surrounded by eight caesium ions.
205
C3 atOms, elements anD COmPOunDs
Csϩ Csϩ surrounded by 8 ClϪ ions
CIϪ In the same way, each ClϪ
a the arrangement of ions
Figure C3.26 The structure of caesium chloride. is surrounded by 8 Csϩ ions
b the unit cell
Properties of ionic compounds Cl– Na+
Na+ O
Ionic compounds have the following properties.
HH
● They are usually solids at room temperature, with Cl–
high melting points. This is due to the strong
electrostatic forces holding the crystal lattice Figure C3.27 Salt (an ionic compound) dissolving in water.
together. A lot of energy is therefore needed to
separate the ions and melt the substance. Question
● They are usually hard substances. 10 Draw diagrams to represent the bonding in each of the
● They usually cannot conduct electricity when solid, following ionic compounds:
a potassium fluoride (KF)
because the ions are not free to move. b lithium chloride (LiCl)
● They mainly dissolve in water. This is because
c magnesium fluoride (MgF2)
water molecules are able to bond with both the d calcium oxide (CaO).
positive and the negative ions, which breaks up
the lattice and keeps the ions apart. Figure C3.27
shows the interaction between water molecules
(the solvent) and sodium and chloride ions from
sodium chloride (the solute).
● They usually conduct electricity when in the
molten state or in aqueous solution. The forces
of attraction between the ions are weakened
and the ions are free to move to the appropriate
electrode. This allows an electric current to
be passed through the molten compound (see
Chapter C5, pp.230–231).
206
Covalent bonding
●● Covalent bonding 0.175 0.175
Another way in which atoms can gain the stability 0.200 0.200
of the noble gas electron configuration is by sharing
the electrons in their outer energy levels. This occurs H 0.150 H
between non-metal atoms, and the bond formed is 0.125
called a covalent bond. The simplest example of 0.100
this type of bonding can be seen by considering the 0.075
hydrogen molecule, H2. 0.050
0.025
Each hydrogen atom in the molecule has one
electron. In order to obtain a full outer energy level 0.074 nm
and gain the electron configuration of the noble gas
helium, each of the hydrogen atoms must have two electron densities in electrons per cubic atomic unit of length
electrons. To do this, the two hydrogen atoms allow c E lectron density map of a hydrogen molecule
their outer energy levels to overlap (Figure C3.28a). Figure C3.28
A molecule of hydrogen is formed, with two
hydrogen atoms sharing a pair of electrons
(Figure C3.28a, b, c). This shared pair of electrons is
known as a single covalent bond and is represented
by a single line as in hydrogen:
H—H
A similar example exists in the diatomic halogen
molecule chlorine, Cl2 (Figure C3.29).
Cl Cl
HH HH
2 hydrogen atoms hydrogen molecule (H2)
a The electron sharing to form the single covalent bond in H2 molecules 2 chlorine atoms
b Model of hydrogen molecule Cl Cl
chlorine molecule (Cl2)
Figure C3.29 The electron sharing to form the single covalent bond in
Cl2 molecules (sharing outer electron shells only).
207
C3 Atoms, elements and compounds
Other covalent compounds outer electron energy levels are shown. Figure C3.31
shows the shape of the methane molecule.
Methane (natural gas), CH4, is a gas whose
molecules contain atoms of carbon and hydrogen. Ammonia, NH3, is a gas containing the elements
The electron structures are: nitrogen and hydrogen. It is used in large amounts
to make fertilisers. The electron configurations of the
6C 2,4 1H 1 two elements are:
The carbon atom needs four more electrons to attain 7N 2,5 1H 1
the electron configuration of the noble gas neon. Each The nitrogen atom needs three more electrons
hydrogen atom needs only one electron to form the to obtain the noble gas structure of neon. Each
electron configuration of helium. Figure C3.30 shows hydrogen requires only one electron to form the
how the atoms gain these electron configurations by noble gas structure of helium. The nitrogen and
the sharing of electrons. You will note that only the
H
H HH
C C
H HH
H
carbon atom 4 hydrogen atoms methane molecule (CH4)
Figure C3.30 F ormation of methane.
H
HCH
H b M odel of the methane molecule. The tetrahedral shape is caused by the
the methane molecule is tetrahedral repulsion of the C–H bonding pairs of electrons.
H
C 109.5°
HH
H
a M ethane molecule.
Figure C3.31
208
Covalent bonding
hydrogen atoms share electrons, forming three single The oxygen atom needs two electrons to gain the
covalent bonds (Figure C3.32). Unlike methane electron configuration of neon. Each hydrogen
the shape of an ammonia molecule is pyramidal requires one more electron to gain the electron
(Figure C3.33). configuration of helium. Again, the oxygen and
hydrogen atoms share electrons, forming a water
Water, H2O, is a liquid containing the elements molecule with two single covalent bonds as shown in
hydrogen and oxygen. The electronic structures of Figure C3.34 on the next page. A water molecule is
the two elements are: V-shaped (Figure C3.35 on the next page).
8O 2,6 1H 1
H HN H
NH
H
H
nitrogen atom 3 hydrogen atoms ammonia molecule (NH3)
Figure C3.32 F ormation of ammonia.
N
HH
H 107°
a A mmonia molecule b M odel of the ammonia molecule. The pyramidal shape is caused by the
Figure C3.33 repulsion between the bonding pairs of electrons as well as the lone pair
(or non-bonding pair) of electrons.
209
C3 Atoms, elements and compounds
H O
O HH
water molecule (H2O)
H
oxygen atom 2 hydrogen atoms
Figure C3.34 F ormation of water.
O
H 104.5° H
a Water molecule b M odel of a V-shaped water molecule
Figure C3.35
OCO
O carbon dioxide molecule (CO2)
C
O
2 oxygen atoms carbon atom
Figure C3.36 Formation of carbon dioxide.
210
Covalent bonding
OCO Carbon dioxide, CO2, is a gaseous compound
a carbon dioxide molecule, note the double covalent bond is found in the air. It contains the elements carbon
represented by a double line and oxygen. The electronic structures of the two
elements are:
b m odel of the linear carbon dioxide molecule
Figure C3.37 6C 2,4 8O 2,6
In this case each carbon atom needs to share four
electrons to gain the electron configuration of neon.
Each oxygen needs to share two electrons to gain the
electron configuration of neon. This is achieved by
forming two double covalent bonds in which two
pairs of electrons are shared in each case, as shown
in Figure C3.36. Carbon dioxide is a linear molecule
(Figure C3.37).
NN H H H H
C C C C
H H H H
NN
ethene
Figure C3.38 Nitrogen. HH
Nitrogen, N2, is a gaseous element found in the C C
air and it contains two atoms of nitrogen. The H H
electronic structures of the nitrogen atoms is: propene Figure C3.39 Ethene.
6N 2,5
H HCbtbwyetoCfwocerHHamernbinoegnacaahtdocomaurbsboalesncwaotevolalmltewnaont dbsiontwngdolebhceyotdvwraoeleegnnetntbhaoetnodmss
In itthsieslceacsteroenaschtongitariongtehneaetloemctrnoenedcos ntofigsuhraarteiothnHroeef C
of
neon. This is achieved by forming a triple covalent H
bond in which three pairs of electrons are shared in (Figure C3.39).
each case, as shown in Figure C3.38. butene
Ethene, C2H4, is a gaseous compound H usHeMdeitnhHtahneoml, aCnHufa3OctuHre, is a volatile liquid compound
used in tHhe of biodiesel. It contains
manufacture of polyethene. It contains the elemH entCs C thCe elCements carbon, oxygen and hydrogen. The
carbon and hydrogen. The electronic structures of H electronHic structures of the three elements are:
the two elements present are: H 6C 2,4 8O 2,6 1H 1
6C 2,4 1H 1 In this case each carbon atom needs to share four
In this case each carbon atom needs to share its four electrons to gain the electron configuration of neon.
electrons to gain the electron configuration of neon. Each oxygen needs to share two electrons to gain
Each hydrogen atom needs only one electron to gain the electron configuration of neon. Each hydrogen
the electron configuration of helium. This is achieved
211
C3 atOms, elements anD COmPOunDs
H strong covalent bonds. Giant molecular substances
H C OH have higher melting points, because the whole
structure is held together by strong covalent bonds.
H It should be noted that in ionic compounds the
interionic forces are much stronger than the
intermolecular forces in simple covalent substances and
so the melting and boiling points are generally higher.
H ● Generally, they do not conduct electricity when
molten or dissolved in water. This is because they
H COH do not contain ions. However, some molecules
actually react with water to form ions. For example,
H hydrogen chloride gas produces aqueous hydrogen
Figure C3.40 Methanol. ions and chloride ions when it dissolves in water:
atom needs to share one electron to gain the electron HCl(g) water H+(aq) + Cl−(aq)
configuration of helium. This is achieved by forming
three carbon–hydrogen, one carbon–oxygen and one ● Generally, they do not dissolve in water. However,
oxygen–hydrogen single covalent bonds as shown in water is an excellent solvent and can interact with and
Figure C3.40. dissolve some covalent molecules better than others.
Covalent substances are generally soluble in organic
Questions solvents. For a further discussion of solubility of
substances in organic solvents see Chapter C14.
11 Draw diagrams to represent the bonding in each of the
following covalent compounds: Allotropy
a tetrachloromethane (CCl4) d hydrogen chloride (HCl) When an element can exist in more than one physical
b oxygen gas (O2) e ethane (C2H6) form in the same state it is said to exhibit allotropy
c hydrogen sulfide (H2S) (or polymorphism). Each of the different physical
12 The diagram shows the arrangement of the outer forms is called an allotrope. Allotropy is actually
quite a common feature of the elements the Periodic
electrons only in a molecule of ethanoic acid. Table chapter C9, p.268. Some examples of elements
which show allotropy are sulfur, tin, iron and carbon.
H O
HC C Allotropes of carbon
H OH Carbon is a non-metallic element which exists in more
than one solid structural form. Its allotropes are called
a Name the different elements found in this compound. graphite and diamond. Each of the allotropes has a
b What is the total number of atoms present in this molecule? different structure (Figures C3.42 and C3.43) and
c Between which two atoms is there a double covalent bond? so the allotropes exhibit different physical properties
d How many single covalent bonds does each carbon (Table C3.13). The different physical properties
that they exhibit lead to the allotropes being used in
atom have? different ways (Table C3.14 and Figure C3.41).
e Write a paragraph explaining the sorts of properties you
would expect this substance to have.
Properties of covalent compounds Table C3.13 Physical properties of graphite and diamond
Covalent compounds have the following properties. Property Graphite Diamond
● As simple molecular substances, they are usually gases, Appearance
A dark grey, shiny A colourless transparent
liquids or solids with low melting and boiling points. Electrical solid crystal which sparkles in light
The melting points are low because of the weak conductivity
intermolecular forces of attraction which exist between Hardness Conducts electricity Does not conduct electricity
simple molecules. These are weaker compared to the
Density/g/cm3 A soft material with A very hard substance
212 a slippery feel 3.51
2.25
Covalent bonding
Figure C3.41 Uses of graphite (as a pencil ‘lead’ and in a squash racket)
and diamond (as a toothed saw to cut marble and on a dentist’s drill).
Table C3.14 Uses of graphite and diamond Graphite
Figure C3.42a (on the next page) shows the structure
Graphite Diamond of graphite. This is a layer structure. Within each
Pencils Jewellery layer each carbon atom is bonded to three others by
Electrodes Glass cutters strong covalent bonds. Each layer is therefore like a
Lubricant Diamond-studded saws giant molecule. Between these layers there are weak
Drill bits forces of attraction (van der Waals’ forces) and so the
Polishers layers will pass over each other easily.
213
C3 Atoms, elements and compounds With only three covalent bonds formed between
carbon atoms within the layers, an unbonded
one layer electron is present on each carbon atom. These
‘spare’ (or delocalised) electrons form electron
clouds between the layers and it is because of these
spare electrons that graphite conducts electricity.
Diamond
In diamond the carbon atoms form a tetrahedral
arrangement (Figure C3.43a (top)). This bonding
scheme gives rise to a very rigid, three-dimensional
structure and accounts for the extreme hardness of
the substance. All the outer energy level electrons of
the carbon atoms are used to form covalent bonds, so
there are no electrons available to enable diamond to
conduct electricity.
a small part of the structure
showing how the layers fit together
a a portion of the graphite structure
a view of a much larger part of the structure
a the structure of diamond
b piece of graphite as imaged through a scanning tunnelling microscope b the Regent Diamond has been worn by Queen Elizabeth II
Figure C3.42 Figure C3.43
214
Covalent bonding
In diamond the carbon atoms form a tetrahedral properties. Since the structure is like that of diamond
arrangement similar to that found in silicon(iv) you might expect them to be similar. However,
oxide. This bonding scheme gives rise to a very rigid, silicon(iv) oxide is not as hard as diamond, which is
three-dimensional structure and accounts for the the hardest naturally occurring substance.
extreme hardness of the substances silicon(iv) oxide
and diamond. All the outer energy level electrons of O
the carbon atoms are used to form covalent bonds, Si
so there are no electrons available to enable diamond
to conduct electricity. a the silicon(iV) oxide structure in quartz
It is possible to manufacture the different
allotropes of carbon. Diamond is made by heating
graphite to about 300 °C at very high pressures.
Diamond made by this method is known as industrial
diamond. Graphite can be made by heating a mixture
of coke and sand at a very high temperature in an
electric arc furnace for about 24 hours.
Silicon(IV) oxide b quartz is a hard solid at room temperature, it has a melting point of
1610 °C and a boiling point of 2230 °C
Silicon is also a Group 4 non-metal element but with
an electronic structure of 2,8,4. Crystalline silicon Figure C3.44
has the same structure as diamond. To change it
into a similar macromolecular structure shown by
silicon(iv) oxide, all you need to do is to modify the
silicon structure by including some oxygen atoms.
You will notice that each silicon atom is bridged
to its neighbours by an oxygen atom forming the
macromolecular structure shown in Figure C3.44.
As before, the structure shown is just a tiny part of
a giant structure extending into three dimensions.
Again the structure of this substance gives rise to its
Questions 14 From your knowledge of the structure of silicon dioxide
and diamond predict the following properties of silicon
13 Explain why graphite conducts electricity while diamond dioxide: hardness, melting and boiling point, electrical
does not. conductivity and solubility.
215
C3 Atoms, elements and compounds
Checklist ● Ion is an atom or group of atoms which has either lost one or
more electrons, making it positively charged, or gained one
After studying Chapter C3 you should know and understand the or more electrons, making it negatively charged.
following.
● Ionic (electrovalent) bond is a strong electrostatic force of
● Allotropy is the existence of an element in two or more attraction between oppositely charged ions.
different forms in the same physical state.
● Ionisation is the process whereby an atom gains or loses an
● Atomic mass unit is exactly 1/12th of the mass of one atom electron(s) to become an ion.
of the most abundant isotope of carbon-12.
● Isotopes are atoms of the same element which possess
● Concentration is a measure of the amount of solute that different numbers of neutrons. They differ in nucleon
has been dissolved in the solvent. The more that has been number (mass number).
dissolved then the higher the concentration and vice versa.
● Lattice is a regular three-dimensional arrangement of
● Covalent bond is a chemical bond formed by the sharing of atoms/ions in a crystalline solid.
one or more pairs of electrons between two atoms.
● Neutron is a fundamental, uncharged sub-atomic particle
● Delocalised refers to spreading out of electrons within the present in the nuclei of atoms.
metal structure.
● Nucleon number (mass number) Symbol A, is the total
● Electron is a fundamental sub-atomic particle with a negative number of protons and neutrons found in the nucleus of an
charge present in all atoms within energy levels around the atom.
nucleus.
● Nucleus is found at the centre of the atom, and contains the
● Electron energy levels (shells) are the allowed energies of protons and neutrons.
electrons in atoms.
● Proton is a fundamental sub-atomic particle which has a
● Electronic structure (configuration) is a shorthand method of positive charge equal in magnitude to that of an electron.
describing the arrangement of electrons within the energy Protons occur in all nuclei.
levels of an atom.
● Proton number (atomic number) Symbol Z, is the number of
● Electrostatic force of attraction is a strong force of attraction protons in the nucleus of an atom. The number of electrons
between opposite charges. present in an atom. It gives the order of the element within
the Periodic Table.
● Giant ionic structure is a lattice held together by the
electrostatic forces of attraction between ions. ● Radioactivity is a property of unstable isotopes. They
disintegrate spontaneously to give off one or more types of
● Giant molecular or macromolecular substance is a substance radiation.
containing thousands of atoms per molecule.
● Radioisotope is a radioactive isotope.
216
C4 Stoichiometry
Combined Co-ordinated
● Use the symbols of the elements and write the formulae of ● Use the symbols of the elements and write the formulae of
simple compounds simple compounds
● Deduce the formula of a simple compound from the relative ● Deduce the formula of a simple compound from the relative
numbers of atoms present numbers of atoms present
● Deduce the formula of a simple compound from a model or a ● Deduce the formula of a simple compound from a model or a
diagrammatic representation diagrammatic representation
● Determine the formula of an ionic compound from the ● Determine the formula of an ionic compound from the
charges on the ions present charges on the ions present
● Construct and use word equations ● Construct and use word equations
● Interpret and balance simple symbol equations ● Interpret and balance simple symbol equations
● Construct and use symbol equations, with state symbols, ● Construct and use symbol equations, with state symbols,
including ionic equations including ionic equations
● Deduce the balanced equation for a chemical reaction,
given relevant information
● Define relative atomic mass, Ar, as the average mass of
naturally occurring atoms of an element on a scale where
the 12C atom has a mass of exactly 12 units
● Define relative molecular mass, Mr, as the sum of the
relative atomic masses (relative formula mass or Mr will be
used for ionic compounds)
● Define the mole in terms of a specific number of particles
called Avogadro’s constant
● Use the molar gas volume, taken as 24 dm3 at room
temperature and pressure
● Calculate stoichiometric reacting masses, volumes of gases
and solutions, and concentrations of solutions expressed
in g/dm3 and mol/dm3 (Calculations involving the idea of
limiting reactants may be set.)
●● Chemical symbols and elements and their symbols is given in Table C4.1 on
formulae the next page.
Symbols Chemical formulae
Chemists use shorthand symbols to label the elements The formula of a compound is made up from the
and their atoms. The symbol consists of one, two symbols of the elements present and numbers to
or three letters, the first of which must be a capital. show the ratio in which the different atoms are
Where several elements have the same initial letter, present. Carbon dioxide has the formula CO2. This
a second letter of the name or subsequent letter is tells you that it contains one carbon atom for every
added. For example, C is used for carbon, Ca for two oxygen atoms. The 2 in the formula tells you
calcium and Cl for chlorine. Some symbols seem that there are two oxygen atoms present in each
to have no relationship to the name of the element, molecule of carbon dioxide. For further discussion
for example Na for sodium and Pb for lead. These see p.210.
symbols come from their Latin names, natrium for
sodium and plumbum for lead. A list of some common Table C4.2 (on the next page) shows the names
and formulae of some common compounds which
you will meet in your study of chemistry.
217
C4 stOiChiOmetRy
Table C4.1 Some common elements and their symbols, the Latin Petrol is a mixture of many different chemicals
names of some of the elements are given in brackets designed to burn effectively in the car engine to
provide power to drive the car forward. One of the
Element Symbol Physical state at room chemicals in petrol is heptane. Heptane is a molecule
temperature and pressure made from 7 carbon atoms and 16 hydrogen atoms
Aluminium Al Solid bonded together. Knowing the relative number
Argon Ar Gas of each type of atom in a molecule allows us to
Barium Ba Solid write down its molecular formula. Heptane has the
Boron B Solid molecular formula C7H16.
Bromine Br Liquid The molecular formula can also be obtained from a
Calcium Ca Solid structural or displayed formula. The molecule ethyl
Carbon C Solid ethanoate can be shown as the structural formula
Chlorine Cl Gas below.
Chromium Cr Solid
Copper (Cuprum) Cu Solid O
Fluorine F Gas
Germanium Ge Solid CH3 C O CH2CH3
Gold (Aurum) Au Solid
Helium He Gas The molecular formula of ethyl ethanoate is found
Hydrogen H Gas by simply counting the number of each type of atom
Iodine I Solid in the molecule. The molecular formula of ethyl
Iron (Ferrum) Fe Solid ethanoate is C4H8O2.
Lead (Plumbum) Pb Solid
Magnesium Mg Solid Another way of showing the molecule ethyl
Mercury (Hydragyrum) Hg Liquid ethanoate is as a model. Here the carbon atoms are
Neon Ne Gas shown by the black spheres, the hydrogen atoms as
Nitrogen N Gas white spheres and oxygen atoms as red spheres.
Oxygen O Gas
Phosphorus P Solid Questions
Potassium (Kalium) K Solid
Silicon Si Solid 1 Write down the molecular formulae of the compounds with
Silver (Argentum) Ag Solid the following numbers and types of atom.
Sodium (Natrium) Na Solid a Ethanoic acid; two carbon atoms, four hydrogen atoms
Sulfur S Solid and two oxygen atoms.
Tin (Stannum) Sn Solid b Methyl ethanoate; three carbon atoms, six hydrogen
Zinc Zn Solid atoms and two oxygen atoms.
Table C4.2 Names and formulae of some common compounds 2 Write down the molecular formulae of these molecules.
Compound Formula O
Ammonia
Calcium hydroxide NH3 H3C CH2 CH2 C
Carbon dioxide Ca(OH)2 CH3
Copper sulfate CO2 O
Ethanol (alcohol) CuSO4
Glucose C2H5OH
Hydrochloric acid C6H12O6
Nitric acid HCl
Sodium carbonate
Sodium hydroxide HNO3
Sulfuric acid Na2CO3
NaOH
H2SO4
218
Chemical symbols and formulae
Formulae of ionic substances and CuCl2 is copper(ii) chloride. Similarly, iron forms
the Fe2+ and Fe3+ ions and so can also form two
On p.203 we saw that ionic compounds contain different compounds with, for example, chlorine:
positive and negative ions, whose charges balance. FeCl2 (iron(ii) chloride) and FeCl3 (iron(iii)
For example, sodium chloride contains one Na+ ion chloride).
for every Cl− ion, giving rise to the formula NaCl.
This method can be used to write down formulae Table C4.3 shows the valencies of a series of ions
which show the ratio of the number of ions present you will normally meet in your study of chemistry.
in any ionic compound.
You will notice that Table C4.3 includes groups
The formula of magnesium chloride is MgCl2. of atoms which have net charges. For example,
This formula is arrived at by each Mg2+ ion the nitrate ion is a single unit composed of one
combining with two Cl− ions, and once again the nitrogen atom and three oxygen atoms and has one
charges balance. The size of the charge on an ion is single negative charge. The formula, therefore, of
a measure of its valency or combining power. Na+ magnesium nitrate would be Mg(NO3)2. You will
has a valency of 1, but Mg2+ has a valency of 2. Na+ notice that the NO3 has been placed in brackets with
can bond (combine) with only one Cl− ion, whereas a 2 outside the bracket. This indicates that there are
Mg2+ can bond with two Cl− ions. two nitrate ions present for every magnesium ion.
The ratio of the atoms present is therefore:
Some elements, such as copper and iron, possess
two ions with different valencies. Copper can form Mg (N O3)2
the Cu+ ion and the Cu2+ ion, with valencies 1 and 1Mg : 2N : 6O
2 respectively. Therefore it can form two different The charge on the element ion is often referred to as
compounds with chlorine, CuCl and CuCl2. We its oxidation state.
can also distinguish the difference by using Roman
numerals in their names: CuCl is copper(i) chloride
Table C4.3 Valencies (and oxidation states) of some elements (ions) and groups of atoms
Valency (oxidation state)
1 2 3
Metals Lithium (Li+) Magnesium (Mg2+) Aluminium (Al3+)
Sodium (Na+) Iron(iii) (Fe3+)
Potassium (K+) Calcium (Ca2+)
Silver (Ag+) Phosphate (PO43−)
Copper(i) (Cu+) Copper(ii) (Cu2+)
(F−) Zinc (Zn2+)
(Cl−)
(Br−) Iron(ii) (Fe2+)
(H+)
(OH−) Lead (Pb2+)
(NO3−)
(NH4+) Barium (Ba2+)
(HCO3−)
Non-metals Fluoride (MnO4−) Oxide (O2−)
Groups of atoms Chloride
Bromide Sulfide (S2−)
Hydrogen
Hydroxide Carbonate (CO32−)
Nitrate Sulfate (SO42−)
Ammonium Dichromate(Vi) (Cr2O72−)
Hydrogencarbonate
Manganate(Vii)
219
C4 stOiChiOmetRy
The ‘cross-over’ method For example, in aluminium sulfide the valencies of
A simple method to work out the formula of the elements are:
compounds is called the ‘cross-over’ method. In this
method it is only necessary to ‘swap’ the valencies of the Al = 3 and S = 2
elements or groups of atoms (or radicals) concerned. Hence the chemical formula is Al2S3.
This is easily done by ‘crossing over’ their valencies, Similarly, in sodium sulfate the valencies are:
by placing the valency of the first element after the
symbol of the second and placing the valency of the Na = 1 and SO4 = 2
second element or radical after the symbol of the first. Hence the chemical formula is Na2SO4.
Questions When we replace the words with symbols for the
reactants and the products and include their physical
3 Using the information in Table C4.3, write the formulae for: state symbols, we obtain:
a copper(i) oxide d lead bromide Fe(s) + S(s) heat FeS(s)
b zinc phosphate e potassium manganate(Vii) Since there is the same number of each type of atom
on both sides of the equation this is a balanced
c iron(iii) chloride f sodium dichromate(Vi). chemical equation.
4 Using the formulae in your answer to question 3, write In the case of magnesium reacting with oxygen,
the word equation was:
down the ratio of atoms present for each of the compounds.
magnesium + oxygen heat magnesium oxide
●● Chemical equations
When we replace the words with symbols for the
Word equations reactants and the products and include their physical
state symbols, it is important to remember that
A word equation describes, in words, a chemical reaction. oxygen is a diatomic molecule:
The reactants are named on the left-hand side of the
arrow, which represents the reaction occurring, and Mg(s) + O2(g) heat MgO(s)
the products appear on the right-hand side. The word
equation for the reaction between zinc and hydrochloric In the equation there are two oxygen atoms on the
acid to produce zinc chloride and hydrogen gas is: left-hand side (O2) but only one on the right (MgO).
We cannot change the formula of magnesium oxide,
zinc + hydrochloric acid → zinc chloride + hydrogen so to produce the necessary two oxygen atoms on
reactants products the right-hand side we will need 2MgO – this means
2 × MgO. The equation now becomes:
Word equations are a useful way of representing
chemical reactions but a better and more useful Mg(s) + O2(g) heat 2MgO(s)
method is to produce a balanced chemical equation.
This type of equation gives the formulae of the
reactants and the products as well as showing the
relative numbers of each particle involved.
Balanced equations often include the physical state There are now two atoms of magnesium on the
symbols: right-hand side and only one on the left. By placing a
(s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous solution 2 in front of the magnesium, we obtain the following
balanced chemical equation:
The word equation to represent the reaction between
iron and sulfur is: 2Mg(s) + O2(g) heat 2MgO(s)
iron + sulfur heat iron(ii) sulfide This balanced chemical equation now shows us that
two atoms of magnesium react with one molecule
of oxygen gas when heated to produce two units of
magnesium oxide.
220
Reacting masses
Ionic equations atomic mass, Ar, of 12.00. All other atoms of the
other elements are given a relative atomic mass
An ionic equation differs from a chemical equation compared to that of carbon.
in that it shows only the chemical species which take
part in the chemical reaction. For example, when An H atom is 1 An Mg atom is twice
sodium chloride solution and silver nitrate solution 12 the mass of a C atom
react together the products are solid silver chloride
and sodium nitrate solution. the mass of a C atom
The chemical equation for the reaction is: H C Mg S Ca
NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
The ions which are present in aqueous solution 1 12 24 32 40
act independently of one another. In this reaction fixed
the sodium ions and nitrate ions start off as ions in
solution and end the reaction unchanged. These Figure C4.1 The relative atomic masses of the elements H, C, Mg, S and Ca.
ions are called spectator ions and they do not appear
in the ionic equation. Without the spectator ions ●● Reacting masses
the ionic equation for the reaction is:
Chemists often need to be able to show the relative
Ag+(aq) + Cl–(aq) → AgCl(s) masses of the atoms involved in a chemical process.
An ionic equation has not only to balance in For example, what mass of carbon dioxide would
terms the number of each type of atom on each side be produced if 6 g of carbon was completely
of the equation but also each side of the equation combusted?
must have the same overall charge.
C + O2 → CO2
Question Instead of using the actual masses of atoms we use
the relative atomic mass to help us answer this type
5 Write the word and balanced chemical equations for the of question.
reactions which take place between: In the example above we can work out the
relative formula mass (RFM) of molecules such
a calcium and oxygen b copper and oxygen. as O2 and CO2 using the relative atomic masses of
the atoms they are made from. The RFM is the sum
of the relative atomic masses of all those elements
shown in the formula of the substance.
●● Relative atomic mass O2 has a relative formula mass of 2 × 16 = 32
CO2 has a relative formula mass of 12 + (2 × 16) = 44
There are at present 118 different elements known. So we can now use the equation to answer the
The atoms of these elements differ in mass because question asked earlier.
of the different numbers of protons, neutrons and
electrons they contain. The actual mass of one atom C + O2 → CO2
is very small. For example, the mass of a single atom 12 32 44
of sulfur is around:
Converting these relative masses to actual masses by
0.000 000 000 000 000 000 000 053 16 g adding mass units, g, would give:
This small quantity is not easy to work with so a
scale called the relative atomic mass scale is used. C + O2 → CO2
In this scale an atom of carbon is given a relative 12 32 44
12 g 32 g 44 g
The above calculation shows that if 12 g of carbon
were burned completely then 44 g of carbon dioxide
221
C4 stOiChiOmetRy
gas would be formed. So 6 g of carbon burning The mole
would result in the formation of 22 g of carbon
dioxide gas. Chemists often need to know how much of a
substance has been formed or used up during a
Let us look at another example. What mass of chemical reaction (Figure C4.2). This is particularly
hydrogen gas would be produced if 46 g of sodium important in the chemical industry, where the
was reacted with water? First of all write down the substances being reacted (the reactants) and the
balanced chemical equation: substances being produced (the products) are
worth thousands of pounds. Waste costs money!
2Na + 2H2O → 2NaOH + H2
Next find the relative atomic mass of sodium To solve this problem they need a way of
from the Periodic Table and work out the relative counting atoms, ions or molecules. Atoms, ions
formula masses of water, sodium hydroxide and and molecules are very tiny particles and it is
hydrogen gas. impossible to measure out a dozen or even a
Relative atomic mass of sodium is 23. hundred of them. Instead, chemists weigh out a
Relative formula mass of water, H2O, is very large number of particles.
(2 × 1) + 16 = 18
Relative formula mass of sodium hydroxide is This number is 6 × 1023 atoms, ions or
23 + 16 + 1 = 40 molecules and is called Avogadro’s constant
Relative formula mass of hydrogen gas, H2, after the famous Italian scientist Amedeo
is 2 × 1 = 2 Avogadro (1776–1856). An amount of substance
Now you can write these masses under the balanced containing 6 × 1023 particles is called a mole
chemical equation taking into account the numbers (often abbreviated to mol). So, a mole of the
used to balance the equation. element magnesium is 6 × 1023 atoms of magnesium
and a mole of the element carbon is 6 × 1023 atoms
of carbon (Figure C4.3).
2Na + 2H2O → 2NaOH + H2
2 × 23 = 46 2 × 18 = 36 2 × 40 = 80 2
These relative masses can now be converted into
actual or reacting masses by putting in mass units,
for example, grams.
2Na + 2H2O → 2NaOH + H2
2 × 23 = 46 2 × 18 = 36 2 × 40 = 80 2
46 g 36 g 80 g 2 g
So the answer to the question of what mass of
hydrogen would be produced if 46 g of sodium was
reacted with water is 2 g.
Questions Figure C4.2 The chemists at paint manufacturers need to know
how much pigment is going to be produced.
6 What mass of carbon dioxide gas would be produced
if 10 g of calcium carbonate reacted with an excess of
hydrochloric acid?
7 What mass of sulfur dioxide would be produced if 64 tonnes
of sulfur was completely reacted with oxygen gas?
222
Calculating moles
The mass of a substance present in any number of
moles can be calculated using the relationship:
mass = number × mass of 1 mole
(in grams) of moles of the element
a A mole of magnesium Example 1
Calculate the mass of a 2 moles and b 0.25 mole of
iron. (Ar: Fe = 56)
a mass of 2 moles of iron
= number of moles × relative atomic mass (Ar)
= 2 × 56
= 112 g
b mass of 0.25 mole of iron
= number of moles × relative atomic mass (Ar)
= 0.25 × 56
= 14 g
If we know the mass of the element then it is
possible to calculate the number of moles of that
element using:
number of moles = mass of the element
mass of 1 mole of that element
Example 2
b A mole of carbon Calculate the number of moles of aluminium present
Figure C4.3 in a 108 g and b 13.5 g of the element. (Ar: Al = 27)
a number of moles of aluminium
●● Calculating moles
= mass of aluminium
We can compare the masses of all the other atoms mass of 1 mole of aluminium
with the mass of carbon atoms. This is the basis of 108
the relative atomic mass scale. Chemists have found = 27
by experiment that if you take the relative atomic
mass of an element in grams, it always contains 6 × = 4 moles
1023 or one mole of its atoms.
b number of moles of aluminium
Moles and elements
= mass of aluminium
For example, the relative atomic mass (Ar) of iron is mass of 1 mole of aluminium
56, so one mole of iron is 56 g. Therefore, 56 g of
iron contains 6 × 1023 atoms. = 13.5
27
The Ar for aluminium is 27. In 27 g of aluminium = 0.5 mole
it is found that there are 6 × 1023 atoms. Therefore,
27 g of aluminium is one mole of aluminium atoms. Moles and compounds
The idea of the mole has been used so far only with
elements and atoms. However, it can also be used
with compounds (Figure C4.4 on the next page).
What is the mass of 1 mole of water (H2O)
molecules? (Ar: H = 1; O = 16)
223
C4 Stoichiometry
a One mole of C2H5OH contains 2 moles of carbon
atoms, 6 moles of hydrogen atoms and 1 mole of
oxygen atoms. Therefore:
mass of 1 mole of ethanol
= (2 × 12) + (6 × 1) + (1 × 16)
= 46 g
b The RFM of ethanol is 46.
Example 2
What is a the mass of 1 mole and b the RFM of
nitrogen gas, N2? (Ar: N = 14)
a Nitrogen is a diatomic gas. Each nitrogen
molecule contains two atoms of nitrogen.
Therefore:
mass of 1 mole of N2
= 2 × 14
= 28 g
Figure C4.4 One mole of water (H2O) (left) and one mole of ethanol b The RFM of N2 is 28.
(C2H5OH) (right) in separate measuring cylinders. The mass of a compound found in any number of
moles can be calculated using the relationship:
mass of = number of moles × mass of 1 mole
compound of the compound of the compound
From the formula of water, H2O, you will see that Example 3
1 mole of water molecules contains 2 moles of Calculate the mass of a 3 moles and b 0.2 moles of
hydrogen (H) atoms and 1 mole of oxygen (O) carbon dioxide gas, CO2. (Ar: C = 12; O = 16)
atoms. The mass of 1 mole of water molecules is a One mole of CO2 contains 1 mole of carbon
therefore:
(2 × 1) + (1 × 16) = 18 g atoms and 2 moles of oxygen atoms. Therefore:
The mass of 1 mole of a compound is called its
molar mass. If you write the molar mass of a mass of 1 mole of CO2
compound without any units then it is the relative = (1 × 12) + (2 × 16)
formula mass, often called the relative molecular = 44 g
mass (Mr). So the relative formula mass of water
is 18. mass of 3 moles of CO2
= number of moles × mass of 1 mole of CO2
Now follow these examples to help you learn and = 3 × 44
understand more about moles and compounds. = 132 g
Example 1 b mass of 0.2 mole of CO2
What is a the mass of 1 mole and b the relative = number of moles × mass of 1 mole of CO2
formula mass (RFM) of ethanol, C2H5OH? = 0.2 × 44
(Ar: H = 1; C = 12; O = 16) = 8.8 g
If we know the mass of the compound then we can
calculate the number of moles of the compound
using the relationship:
224
Calculating moles
number of moles = mass mass of compound Example 1
of compound of 1 mole of the compound
Calculate the number of moles of ammonia gas,
NH3, in a volume of 72 dm3 of the gas measured
Example 4 at rtp.
Calculate the number of moles of magnesium oxide, number of moles volume of ammonia in dm3
MgO, in a 80 g and b 10 g of the compound. of ammonia = 24 dm3
(Ar: O = 16; Mg = 24)
= 72
a One mole of MgO contains 1 mole of magnesium 24
atoms and 1 mole of oxygen atoms. Therefore:
= 3 moles
mass of 1 mole of MgO
= (1 × 24) + (1 × 16) Example 2
= 40 g Calculate the volume of carbon dioxide gas, CO2,
occupied by a 5 moles and b 0.5 mole of the gas
number of moles of MgO in 80 g measured at rtp.
= mass mass of MgO a volume of CO2
of 1 mole of MgO = number of moles of CO2 × 24 dm3
80 = 5 × 24
= 40 = 120 dm3
= 2 moles b volume of CO2
= number of moles of CO2 × 24 dm3
b number of moles of MgO in 10 g = 0.5 × 24
= 12 dm3
= mass mass of MgO
of 1 mole of MgO The volume occupied by one mole of any gas must
10 contain 6 × 1023 molecules. Therefore, it follows
= 40 that equal volumes of all gases measured at the
same temperature and pressure must contain the
= 0.25 moles same number of molecules. This idea was also first
put forward by Amedeo Avogadro and is called
Moles and gases Avogadro’s Law.
Many substances exist as gases. If we want to find Moles and solutions
the number of moles of a gas we can do this by
measuring the volume rather than the mass. Chemists often need to know the concentration
of a solution. Sometimes it is measured in grams
Chemists have shown by experiment that: per cubic decimetre (g/dm3) but more often
concentration is measured in moles per cubic
One mole of any gas occupies a volume of decimetre (mol/dm3). When 1 mole of a substance
approximately 24 dm3 (24 litres) at room temperature is dissolved in water and the solution is made up to
and pressure (rtp). This quantity is also known as the 1 dm3 (1000 cm3), a 1 molar (1 mol/dm3) solution
molar gas volume, Vm. is produced. Chemists do not always need to make
Therefore, it is relatively easy to convert volumes of up such large volumes of solution. A simple method
gases into moles and moles of gases into volumes of calculating the concentration is by using the
using the following relationship: relationship:
number of moles = volume of the gas (in dm3 at rtp)
of a gas 24 dm3
or
number of moles
volume of = number of moles × 24 dm3 concentration (in mol/dm3) = volume (in dm3)
a gas of gas (in dm3 at rtp)
225
C4 stOiChiOmetRy
Example 1 1 mole of KOH contains 1 mole of potassium,
Calculate the concentration (in mol/dm3) of a 1 mole of oxygen and 1 mole of hydrogen.
solution of sodium hydroxide, NaOH, which was Therefore:
made by dissolving 10 g of solid sodium hydroxide
in water and making up to 250 cm3. (Ar: H = 1; mass of 1 mole of KOH
O = 16; Na = 23) = (1 × 39) + (1 × 16) + (1 × 1)
= 56 g
1 mole of NaOH contains 1 mole of sodium,
1 mole of oxygen and 1 mole of hydrogen. Therefore:
Therefore:
mass of KOH in 1 mole
mass of 1 mole of NaOH = number of moles × mass of 1 mole
= (1 × 23) + (1 × 16) + (1 × 1) = 1 × 56
= 40 g = 56 g
number of moles of NaOH in 10 g Questions
mass of NaOH
= mass of 1 mole of NaOH Use the values of Ar which follow to answer the questions
below.
10 H = 1; C = 12; N = 14; O = 16; Ne = 20; Na = 23;
= 40 Mg = 24; S = 32; K = 39; Fe = 56; Cu = 63.5; Zn = 65.
= 0.25 moles One mole of any gas at rtp occupies 24 dm3.
(250 cm3 = 1205000 dm3 = 0.25 dm3) 8 Calculate the number of moles in:
a 2 g of neon atoms
concentration of the NaOH solution b 4 g of magnesium atoms
number of moles of NaOH c 24 g of carbon atoms.
= volume of solution (dm3)
9 Calculate the mass of:
= 0.25 a 0.1 mole of oxygen molecules
0.25 b 5 moles of sulfur atoms
c 0.25 mole of sodium atoms.
= 1 mol/dm3
Sometimes chemists need to know the mass of a 10 Calculate the number of moles in:
substance that has to be dissolved to prepare a known a 9.8 g of sulfuric acid (H2SO4)
volume of solution at a given concentration. A simple b 40 g of sodium hydroxide (NaOH)
method of calculating the number of moles and so the c 720 g of iron(ii) oxide (FeO).
mass of substance needed is by using the relationship:
11 Calculate the mass of:
number of = concentration × volume in solution a 2 moles of zinc oxide (ZnO)
moles (in mol/dm3) (in dm3) b 0.25 mole of hydrogen sulfide (H2S)
c 0.35 mole of copper(ii) sulfate (CuSO4).
Example 2
Calculate the mass of potassium hydroxide, KOH, 12 Calculate the number of moles at rtp in:
that needs to be used to prepare 500 cm3 of a a 2 dm3 of carbon dioxide (CO2)
2 mol/dm3 solution in water. (Ar: H = 1; O = 16; b 240 dm3 of sulfur dioxide (SO2)
K = 39) c 20 cm3 of carbon monoxide (CO).
number of moles of KOH 13 Calculate the volume of:
a 0.3 mole of hydrogen chloride (HCl)
= concentration of solution × volume of solution b 4.4 g of carbon dioxide
(mol/dm3) (dm3) c 34 g of ammonia (NH3).
= 2 × 500 14 Calculate the concentration of solutions containing:
1000 a 0.2 mole of sodium hydroxide dissolved in water and
made up to 100 cm3
b 9.8 g of sulfuric acid dissolved in water and made up
to 500 cm3.
15 Calculate the mass of:
a copper(ii) sulfate (CuSO4) which needs to be used to
prepare 500 cm3 of a 0.1 mol/dm3 solution
b potassium nitrate (KNO3) which needs to be used to
prepare 200 cm3 of a 2 mol/dm3 solution.
= 1 mole
226
moles and chemical equations
●● Moles and chemical equations
When we write a balanced chemical equation we are
indicating the numbers of moles of reactants and
products involved in the chemical reaction. Consider
the reaction between magnesium and oxygen.
magnesium + oxygen → magnesium oxide
2Mg(s) + O2(g) → 2MgO(s)
This shows that 2 moles of magnesium react with
1 mole of oxygen to give 2 moles of magnesium
oxide.
Using the ideas of moles and masses we can use
this information to calculate the quantities of the
different chemicals involved.
2Mg(s) + O2(g) → 2MgO(s)
2 moles 1 mole 2 moles
2 × 24 1 × (16 × 2) 2 × (24 + 16)
= 48 g = 32 g = 80 g
You will notice that the total mass of reactants is Figure C4.5 A rotary kiln for burning (calcining) limestone into lime,
equal to the total mass of product. This is true for located in Moha, Belgium.
any chemical reaction and it is known as the Law
of conservation of mass. This law was understood CaCO3
by the Greeks but was first clearly formulated by
Antoine Lavoisier in 1774. Chemists can use this
idea to calculate masses of products formed and
reactants used in chemical processes before they are
carried out.
Example 1 – using a solid 10 tonnes of limestone
Lime (calcium oxide, CaO) is used in the heat
manufacture of mortar. It is manufactured in large
quantities in Europe (see Figure C4.5) by heating
limestone (calcium carbonate, CaCO3).
The equation for the process is:
CaCO3(s) → CaO(s) + CO2(g)
1 mole 1 mole 1 mole
(40 + 12 + (3 × 16)) 40 + 16 12 + (2 × 16)
= 100 g = 56 g = 44 g
Calculate the amount of lime produced when CaO
10 tonnes of limestone are heated (Figure C4.6).
(Ar: C = 12; O = 16; Ca = 40)
1 tonne (t) = 1000 kg
1 kg = 1000 g
From this relationship between grams and tonnes we ? tonnes of lime
can replace the masses in grams by masses in tonnes. Figure C4.6 How much lime is produced?
227
C4 Stoichiometry
Hence CaCO3(s) → CaO(s) + CO2(g) number of moles of magnesium
100 t 56 t 44 t
10 t 5.6 t 4.4 t mass of magnesium
= mass of 1 mole of magnesium
The equation now shows that 100 t of limestone will
produce 56 t of lime. Therefore, 10 t of limestone = 6
will produce 5.6 t of lime. 24
= 0.25 moles
Example 2 – using a gas
Many chemical processes involve gases. The volume Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g)
of a gas is measured more easily than its mass. This 1 mole 1 mole 1 mole 1 mole
example shows how chemists work out the volumes 0.25 mol 0.25 mol 0.25 mol 0.25 mol
of gaseous reactants and products needed using
Avogadro’s Law and the idea of moles. So 0.25 mol of H2SO4(aq) is required. Using:
Some rockets use hydrogen gas as a fuel. When volume of H2SO4(aq) (dm3)
hydrogen burns in oxygen it forms water vapour.
Calculate the volumes of a O2(g) used and b water, = moles of H2SO4
H2O(g), produced if 960 dm3 of hydrogen gas, concentration of H2SO4 (mol/dm3)
H2(g), were burned in oxygen. (Ar: H = 1; O = 16) 0.25
Assume 1 mole of any gas occupies a volume of = 1
24 dm3.
= 0.25 dm3 or 250 cm3
2H2(g) + O2(g) → 2H2O(g) Example 4 – using a solution
2 moles 1 mole 2 moles What is the concentration of sodium hydroxide
2 × 24 1 × 24 2 × 24 solution used in the following neutralisation
= 48 dm3 = 24 dm3 = 48 dm3 reaction? 40 cm3 of 0.2 mol/dm3 solution of
hydrochloric acid just neutralised 20 cm3 of sodium
Therefore: hydroxide solution.
(× 2) 96 dm3 48 dm3 96 dm3 number of moles of HCl used
(× 10) 960 dm3 480 dm3 960 dm3 = concentration (mol/dm3) × volume (dm3)
= 0.2 × 0.04
When 960 dm3 of hydrogen are burned in oxygen: = 0.008 moles
a 480 dm3 of oxygen are required and HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
b 960 dm3 of H2O(g) are produced. 1 mole 1 mole 1 mole 1 mole
Example 3 – using a solution 0.008 mol 0.008 mol 0.008 mol 0.008 mol
Chemists usually carry out reactions using solutions.
If they know the concentration of the solution(s) You will see that 0.008 mole of NaOH was present.
they are using they can find out the quantities The concentration of the NaOH(aq) is given by:
reacting.
concentration of NaOH (mol/dm3)
Calculate the volume of 1 mol/dm3 solution of
H2SO4 required to react completely with 6 g of = number of moles of NaOH
magnesium. (Ar: Mg = 24) volume of NaOH (dm3)
(volume of NaOH in dm3 = 20 = 0.02)
1000
0.008
= 0.02
= 0.4 mol/dm3
228
limiting reactant
●● Limiting reactant So, we can see that the reaction starts with the
same amount of magnesium and hydrochloric acid.
A chemical reaction stops when one of the reactants From the balanced chemical equation, you can see
has been completely used up. This reactant is known there is a 1:2 mole ratio between magnesium and
as the limiting reactant. It is this amount of this hydrochloric acid, so 0.1 moles of magnesium will
reactant which was used up in the reaction which react with 0.2 moles of hydrochloric acid. Since we
determines the amount of the products which are have only used 0.1 moles of hydrochloric acid this
made. It is, therefore, important that the limiting will all be used up in the reaction. The hydrochloric
reactant is determined before trying to determine acid is the limiting reactant. The magnesium is in
how much of a product is made during at reaction. excess, 0.05 moles of it will be left at the end of the
reaction.
For example, when 2.4 g of magnesium reacts with
50 cm3 of 2 mol/dm3 hydrochloric acid what volume So when we determine the volume of hydrogen gas
of hydrogen gas will be produced? produced it is the amount of hydrochloric acid that
we must use in the calculation.
The chemical equation for this reaction is:
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) There is a 2:1 mole ratio between hydrochloric
Moles of magnesium reacting = 2.4/24 acid and hydrogen so 0.1 moles of hydrochloric acid
= 0.1 moles will give 0.05 moles of hydrogen gas.
Moles of hydrochloric acid reacting = 2 × 50/1000
= 0.1 moles Volume of hydrogen gas = 0.05 × 24 = 1.2 dm3.
Questions a Write a word and balanced symbol equation for the
16 This question concerns the reaction of copper(ii) carbonate reaction taking place.
with dilute hydrochloric acid. The equation for the b Draw a diagram of an apparatus which could be used to
reaction is:
CuCO3(s) + 2HCl(aq) → CuCl2(aq) + CO2(aq) + H2O(l) carry out this experiment and collect the hydrogen gas.
Calculate the mass of copper(ii) carbonate used to produce c How many moles of magnesium were used?
60 cm3 of carbon dioxide. One mole of a gas occupies d Using the equation you have written in your answer to a,
24 dm3 at room temperature and pressure (rtp).)
calculate the number of moles of hydrogen and hence the
17 0.048 g of magnesium was reacted with excess dilute
hydrochloric acid at room temperature and pressure. volume of this gas produced.
The hydrogen gas given off was collected. e Calculate the volume of a solution containing 0.1mol/dm3
hydrochloric acid which would be needed to react exactly
with 0.048g of magnesium.
Checklist
After studying Chapter C4 you should know and understand the following.
● Avogadro’s Law states that equal volumes of all gases ● To calculate moles of solutions
concentration of a solution (in mol/dm3) = number of
measured under the same conditions of temperature and moles of solute volume (in dm3)
number of moles = concentration (mol/dm3) × volume
pressure contain equal numbers of molecules. of solution (in dm3)
● To calculate moles of compounds ● One mole is the amount of substance which contains
6 × 1023 atoms, ions or molecules. This number is called
mass of compound (in grams) = Avogadro’s constant.
number of moles × mass of 1 mole of compound ● Atoms – 1 mole of atoms has a mass equal to the relative
atomic mass (Ar) in grams.
number of moles = mass of compound
mass of 1 mole of compound ● Molecules – 1 mole of molecules has a mass equal to the
relative molecular mass (Mr) in grams.
● To calculate moles of elements
● Molecular formula shows the actual number of atoms of
mass of compound (in grams) = each element present in one molecule.
number of moles × mass of 1 mole of compound ● Relative formula mass (RFM) is the sum of the relative
mass of the element atomic masses of all those elements shown in the formula
number of moles = mass of 1 mole of that element of the substance. This is often referred to as the relative
molecular mass (Mr).
● To calculate moles of gases
1 mole of any gas occupies 24 dm3 (litres) at room
temperature and pressure (rtp). gas (in dm3
volume of the 24 dm3 at rtp)
number of moles of gas =
229
C5 Electricity and chemistry
Combined Co-ordinated
● Define electrolysis as the breakdown of an ionic compound ● Define electrolysis as the breakdown of an ionic compound
when molten or in aqueous solution by the passage of when molten or in aqueous solution by the passage of
electricity electricity
● Use the terms inert electrode, electrolyte, anode and cathode ● Use the terms inert electrode, electrolyte, anode and cathode
● D escribe electrolysis in terms of the ions present and the ● Describe electrolysis in terms of the ions present and the
reactions at the electrodes, in terms of gain of electrons by reactions at the electrodes, in terms of gain of electrons by
cations and loss of electrons by anions to form atoms cations and loss of electrons by anions to form atoms
● Describe the electrode products and the observations made, ● Describe the electrode products, using inert electrodes
using inert electrodes (platinum or carbon), in the electrolysis (platinum or carbon), in the electrolysis of: molten lead(ii)
of: molten lead(ii) bromide; concentrated aqueous sodium bromide; concentrated aqueous sodium chloride; dilute
chloride; dilute sulfuric acid sulfuric acid
● P redict the products of the electrolysis of a specified ● S tate the general principle that metals or hydrogen are
molten binary compound formed at the negative electrode (cathode), and that non-
metals (other than hydrogen) are formed at the positive
electrode (anode)
● Relate the products of electrolysis to the electrolyte and
electrodes used, exemplified by the specific examples in
the Core together with aqueous copper(ii) sulfate using
carbon electrodes and using copper electrodes (as used in
the refining of copper)
● Construct simple ionic half-equations for the formation of
elements at the cathode
● Describe electroplating with copper
● Predict the products of the electrolysis of a specified
molten binary compound
● Describe, in outline, the chemistry of the manufacture
of: aluminium from pure aluminium oxide in molten
cryolite; chlorine, hydrogen and sodium hydroxide from
concentrated aqueous sodium chloride
(Starting materials and essential conditions should be
given but not technical details or diagrams.)
What do all the items in the photographs shown ●● Substances that conduct electricity to a small
in Figure C5.1 have in common? They all involve extent in the molten state or in solution are called
electricity through a process known as electrolysis. weak electrolytes.
Electrolysis is the breakdown of an ionic compound,
molten or in solution, by the passage of electricity The electric current enters and leaves the electrolyte
through it. The substance which is decomposed is through electrodes, which are usually made of
called the electrolyte (Figure C5.2). An electrolyte is a unreactive metals such as platinum or of the non-
substance that conducts electricity when in the molten metal carbon (graphite). These are said to be inert
state or in solution. electrodes because they do not react with the
The electricity is carried through the electrolyte by products of electrolysis. The names given to the
ions. In the molten state and in solution the ions are two electrodes are cathode, the negative electrode
free to move to the appropriate electrodes due to which attracts cations (positively charged ions), and
weakened forces of attraction between them. anode, the positive electrode which attracts anions
●● Substances that do not conduct electricity when (negatively charged ions).
in the molten state or in solution are called non-
electrolytes.
230
Electrolysis of lead(ii) bromide
a This watch has a thin coating of gold over steel;
the thin coating is produced by electrolysis
b This picture frame has been silver plated using an c A luminium is produced by electrolysis
electroplating process involving electrolysis
Figure C5.1
inert carbon inert carbon Note that the conduction which takes place in the
cathode (Ϫ) anode (ϩ) electrodes is due to the movement of delocalised
(an electrode) (an electrode) electrons (pp.214 and 278) whereas in the electrolyte,
as stated earlier, the charge carriers are ions.
Electrolysis is very important in industry. To help
you to understand what is happening in the process
shown in the photographs, we will first consider the
electrolysis of lead(ii) bromide.
electrolyte ●● Electrolysis of lead(II)
(e.g. lead bromide
bromide)
Consider Figure C5.3 on the next page, which shows
heat if necessary solid lead(ii) bromide (PbBr2) in a crucible with
Figure C5.2 T he important terms used in electrolysis. two carbon electrodes in contact with it. When the
electrodes are first connected, the bulb does not light,
The transfer of charge during electrolysis is by: because the solid compound does not allow electricity
● the movement of electrons in the metallic or to pass through it. However, when the compound
is heated until it is molten, the bulb does light. The
graphite electrodes lead(ii) bromide is now behaving as an electrolyte.
● the removal or addition of electrons from the When this happens an orange-red vapour is seen at
the anode and lead metal is produced at the cathode.
external circuit at the electrodes
● the movement of ions in the electrolyte. The break-up (decomposition) of lead(ii) bromide
into its constituent elements by the passage of an
electric current is called electrolysis.
molten lead(ii) bromide →● bromine + lead
PbBr2(l) → Br2 (g) + Pb(l)
231
C5 eleCtRiCity anD ChemistRy
cathode (Ϫ) anode (ϩ)
cathode (Ϫ) (ϩ) anode
Pb2ϩ BrϪ molten lead (II)
BrϪ bromide
graphite (PbBr2)
electrodes
solid lead (II) Figure C5.4 The lead ions (Pb2+) are attracted to the cathode and the
bromide heat bromide ions (Br−) are attracted to the anode.
strongly
This process of gaining electrons is called
bromine vapour reduction.
(orange) To form bromine molecules each bromide ion
must first of all move to the anode and lose its extra
molten lead molten lead (II) negative charge at the anode and so form a neutral
bromide bromine atom:
bromide ion → bromine atom + electron
Br−(l) → Br + e−
Two bromine atoms then combine to form a
bromine molecule:
bromine atoms → bromine molecule
2Br → Br2(g)
This process of losing electrons is called oxidation.
Figure C5.3 The electrolysis of molten lead(ii) bromide.
Questions
For lead metal to be formed, or deposited, at the 1 Explain the meaning of each of the following terms.
cathode, the lead ions must be attracted to and
move towards the electrode (Figure C5.4). To Use a suitable example, in each case, to help with your
produce lead metal atoms these lead ions must each
collect two electrons at the cathode: explanation:
lead ion + electrons → lead atom a anode d electrolyte
Pb2+(l) + 2e− → Pb(l)
b cathode e anion
c electrolysis f cation.
2 Predict the products of the electrolysis of:
a potassium chloride
b lead oxide.
●● Extraction of aluminium oxide
Aluminium is the most abundant metallic element in metal (see Chapter C10, p.279) and is difficult to
the Earth’s crust. It makes up 8% of the crust and is extract from its ore. Reactive metals hold on tightly
found mainly as bauxite (which is mainly aluminium to the element(s) they have combined with and,
oxide) (Figure C5.5), as well as in the minerals today, many, including aluminium, are therefore
cryolite and mica, and also in clay. It is a reactive extracted from their ores by electrolysis.
232
Extraction of aluminium oxide
Uses of aluminium are many and varied: ● The purified aluminium oxide is dissolved in
molten cryolite (Na3AlF6) in an electrolysis
Alloy or metal Use cell. Cryolite is used to reduce the working
element temperature of the Hall–Héroult cell from
2017 °C (the melting point of pure aluminium
alloy in aircraft construction and bicycle parts oxide) to between 800 and 1000 °C. Therefore,
the cryolite provides a considerable saving in the
element metal cans, cooking foils, car body parts, energy requirements of the process. In recent
saucepans, electricity cables years, it has become necessary to manufacture the
cryolite since the naturally occurring mineral is in
The process of extraction is often called the Hall– very short supply.
Héroult process after the two people who developed
it. The process involves the electrolysis of aluminium
oxide (alumina) in the following stages.
● Bauxite, an impure form of aluminium oxide, is
treated with sodium hydroxide to obtain pure
aluminium oxide by removing impurities such
as iron(iii) oxide and sand. This improves the
conductivity of the molten aluminium oxide.
Figure C5.5 Bauxite mining.
233
C5 eleCtRiCity anD ChemistRy
crust of solid electrolyte problems arise with other metals being deposited,
since the cryolite is largely ‘unaffected’ by the flow
graphite anodes of electricity. Problems do arise, however, with the
graphite anodes. At the working temperature of the
ϩ cell, the oxygen liberated reacts with the graphite
anodes, producing carbon dioxide.
aluminium Ϫ
out carbon + oxygen → carbon dioxide
4m C(s) + O2(g) → CO2(g)
siphon wide
The anodes burn away and have to be replaced on a
1m regular basis.
deep
The electrolysis of aluminium oxide is a
8m long continuous process in which vast amounts of
electricity are used. In order to make the process an
molten aluminium tank lined economic one, a cheap form of electricity is required.
collecting on with graphite Hydroelectric power (HEP) is usually used for this
floor of cell cathode process. The plant shown in Figure C5.7 uses an
HEP scheme to provide some of the electrical energy
molten electrolyte required for this process. Further details about HEP
are given in Chapter P2, p.367.
Figure C5.6 The Hall–Héroult cell is used in industry to extract
aluminium by electrolysis. Figure C5.7 An aluminium smelting plant.
● The molten mixture is then electrolysed in a cell
similar to that shown in Figure C5.6.
The anodes of this process are blocks of graphite
which are lowered into the molten mixture from
above. The cathode is the graphite lining of the steel
vessel containing the cell.
Aluminium oxide is an ionic compound. When
it is melted the ions become mobile, as the strong
electrostatic forces of attraction between them
are broken by the input of heat energy. During
electrolysis the negatively charged oxide ions are
attracted to the anode (the positive electrode),
where they lose electrons (oxidation).
oxide ions → oxygen molecules + electrons
2O2−(l) → O2(g) + 4e−
The positive aluminium ions are attracted to the
cathode (the negative electrode). They gain electrons
to form molten aluminium metal (reduction).
aluminium ions + electrons → aluminium metal
Al3+(l) + 3e− → Al(l)
A handy way of remembering it is OIL RIG
(Oxidation Is Loss, Reduction Is Gain of electrons).
The overall reaction which takes place in the cell is:
aluminium oxide electrolysis aluminium + oxygen
2Al2O3(l) 4Al(l) 3O2(g)
The molten aluminium collects at the bottom of the
cell and it is siphoned out at regular intervals. No
234
electrolysis of aqueous solutions
Questions oxygen hydrogen
collected collected
3 Produce a flow chart to summarise the processes involved here here
in the extraction of aluminium metal from bauxite.
4 Explain why the mixture of gases formed at the anode
contains oxygen, carbon dioxide and traces of fluorine.
●● Electrolysis of aqueous water (with a little dilute
solutions sulfuric acid added to increase
its conductivity)
Other industrial processes involve the electrolysis of
aqueous solutions. To help you to understand what platinum
is happening in these processes, we will first consider electrodes (inert)
the electrolysis of water.
anode (ϩ) cathode (Ϫ)
Pure water is a very poor conductor of electricity
because there are so few ions in it. However, it can power
be made to decompose if an electric current is passed supply
through it in a Hofmann voltameter, shown below in
Figure C5.8. Figure C5.9 A Hofmann voltameter used to electrolyse water.
To enable water to conduct electricity better because
it only contain a few H+ and OH− ions, some dilute
sulfuric acid (or sodium hydroxide solution) is added.
When the power is turned on and an electric current
flows through this solution, gases can be seen to be
produced at the two electrodes and they are collected
in the side arms of the apparatus. After about
20 minutes, roughly twice as much gas is produced at
the cathode as at the anode.
The gas collected at the cathode burns with a squeaky
pop, showing it to be hydrogen gas. For hydrogen to
be collected in this way, the positively charged hydrogen
ions must have moved to the cathode.
hydrogen ions + electrons → hydrogen molecules
4H+(aq) + 4e− → 2H2(g)
The OH−(aq) ions are attracted to the anode and
lose electrons. Water and oxygen are formed. The
oxygen collects at the anode. It relights a glowing
splint, showing it to be oxygen.
Figure C5.8 A Hofmann voltameter. This gas is produced in the following way.
hydroxide → water + oxygen + electrons
ions molecules molecules
4OH−(aq) → 2H2O(l) + O2(g) + 4e−
This experiment was first carried out by Sir Humphry
Davy. It confirmed that the formula for water was H2O.
235
C5 eleCtRiCity anD ChemistRy
It should be noted that in the electrolysis of dilute Everything in Figure C5.10 is made using substances
sulfuric acid, platinum (an inert electrode) may be obtained from salt. Salt is one of the most important
replaced by carbon (graphite). The only difference to raw materials. Salt is found in the sea and also in
occur is that as well as oxygen being produced at the underground deposits left by ancient seas.
anode, a little carbon dioxide will also be formed.
Sodium chloride is an ionic compound of sodium
Question B and chlorine. If we wish to obtain the constituent
elements from sodium chloride then we can
5A electrolyse either the molten sodium chloride or
aqueous sodium chloride.
XY
dilute sulfuric acid The chlor-alkali industry
–+ The electrolysis of saturated sodium chloride
solution (brine) is the basis of a major industry.
This is a diagram of an experiment in which electricity was In countries where rock salt (sodium chloride) is
passed through a mixture of distilled water containing a found underground it is mined. In other countries
little dilute sulfuric acid. it can be obtained by evaporation of sea water in
a Name the gas that collects at A. large shallow lakes. Three very important substances
b Name the gas that collects at B. are produced in this electrolysis process – chlorine,
c If 100 cm3 of gas collects in A how much would there sodium hydroxide and hydrogen. The electrolytic
process is a very expensive one, requiring vast
be in B? amounts of electricity. The process is economical only
d Name the metal usually used for X and Y. because all three products have a large number of uses
e X is called the ___________ . (Figure C5.11).
f Y is called the ___________ .
g Write down the formulae of the three ions present in There are two well-established methods for
electrolysing brine, the diaphragm cell and the
the solution. mercury cell. However, recent developments in
h Write down the equations for the reactions that take electrolysis technology, by chemical engineers, have
produced the membrane cell (Figure C5.12 on
place at both X and Y (or describe the changes that take p.238). This method is now preferred to the other
place if you cannot write the equations). two because it produces a purer product, it causes
less pollution and it is cheaper to run.
The electrolysis of aqueous sodium
chloride The brine is first purified to remove calcium,
strontium and magnesium compounds by a process
of ion exchange.
The membrane cell is used continuously, with
fresh brine flowing into the cell as the process
breaks up the brine. The cell has been designed to
ensure that the products do not mix. The ions in
this concentrated sodium chloride solution are:
from the water: H+(aq) OH−(aq)
from the sodium chloride: Na+(aq) Cl−(aq)
When the current flows, the chloride ions, Cl−(aq),
are attracted to the anode. Chlorine gas is produced
by the electrode process.
Figure C5.10 The PVC pipes, hydrochloric acid and fertiliser are made chloride ions oxidation chlorine molecules + electrons
from substances obtained by the electrolysis of saturated salt solution. 2Cl−(aq) Cl2(g) + 2e−
236
The words you are searching are inside this book. To get more targeted content, please make full-text search by clicking here.
Cambridge IGCSE Combined and Co-ordinated Sciences
Discover the best professional documents and content resources in AnyFlip Document Base.
Search
Cambridge IGCSE Combined and Co-ordinated Sciences
- 1 - 50
- 51 - 100
- 101 - 150
- 151 - 200
- 201 - 250
- 251 - 300
- 301 - 350
- 351 - 400
- 401 - 450
- 451 - 500
- 501 - 546
Pages: