HKDSE Chemistry Bridging Programe 1A

© 2009 Aristo Educational Press Ltd. 14/F Lok's Industrial Building, 204 Tsat Tsz Mui Road, North Point, Hong Kong. Tel.: 2811 2908 Fax: 2565 6626 Website: http://www.aristo.com.hk

All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photo-copying, recording or otherwise, without the prior permission of Aristo Educational Press Ltd.

First published July, 2009

Chapter 1

Fundamentals of Chemistry

1

1.1

What is chemistry about?

1

1.2

Chemistry in our lives today

1

1.3

Classification of matter

2

1.4

Properties of substances

8

1.5

Physical and chemical changes

10

1.6

Working in chemistry laboratory

11

Key terms

17

Summary

17

Part I

Planet Earth

Chapter 2

The atmosphere

19

2.1

Getting to know our planet Earth

19

2.2

The atmosphere

20

2.3

Separation of oxygen and nitrogen from air

22

2.4

Properties of oxygen

23

Key terms

24

Summary

25

Chapter 3

Oceans

26

3.1

Introducing oceans and seas

26

3.2

Composition of sea water

26

3.3

Extraction of common salt from sea water

27

3.4

Tests for sodium and chloride ions in common salt

31

3.5

Tests for the presence of water in a sample

33

3.6

Electrolysis of sea water and uses of products

33

Key terms

35

Summary

36

Chapter 4

Rocks and minerals

37

4.1

Rocks

37

4.2

Extraction of metals from their ores

38

4.3

Limestone, chalk and marble

40

4.4

Weathering and erosion of rocks

40

4.5

Chemical changes involving calcium carbonate

42

4.6

Tests for calcium carbonate in a sample of limestone/chalk/marble

45

/

/

Key terms

47

Summary

48

Chapter 1 Fundamentals of chemistry

1.1

What is chemistry about?

1.1

Chemistry is the study of substances, about their compositions, structures, properties and the changes among

N1

Note 1 Refer to Chapter 8 and Chapter 9 for structures of substances.

them.

Chemistry is a branch of science. ‘Science’ means the N2 knowledge gathered systematically from observations and experiments.

1.2

Note 2 Refer to ‘Supplementary information: The scientific method’ in the Teacher’s Guide.

Chemistry in our lives today

1.2

Clothing, food, housing, transport and medical care are the basic necessities of life. Chemistry plays a major role in each of these parts of our modern lives.

Class practice 1.1

1.1

The photos below are some commodities or facilities in our daily lives which are affected by or related to chemistry. Write down the names of chemicals in them. The first one has been done for you as an example. A1.1

(a) Clothing polyester, nylon, dyes

(d) Transport metals, alloys, fuels, glass, plastics

(b) Food fertilizers, insecticides, food additives

(e) Medicines drugs, antibiotics, artificial hormones

(c) Housing metals, alloys, cement, glass, plastics

(f) Amusement park facilities metals, alloys, cement, glass, plastics, semi-conductors

1


1.3


Elements

N3

1.3

Note 3 Refer to ‘Supplementary information: Kinetic theory of matter’ in the Teacher’s Guide.

An element is a pure substance that cannot be broken down into anything simpler by chemical methods.

Oxygen, hydrogen and carbon are elements. Until January 2008, scientists have discovered or reported

2008

118 elements. (You can find the names of the elements in the

(

Periodic Table on the front inside cover of the book.)

)

Percentage by mass of elements in nature The percentage by mass of elements in nature is shown in

1.1

Figure 1.1.

silicon ( 27.7%

oxygen 46.6%

iron

aluminium 8.1%

5.0%

all other elements 1.5%

)

magnesium 2.1%

potassium 2.6%

sodium 2.8%

calcium 3.6%

Figure 1.1 Percentage by mass of elements in nature.

2

1

118


Oxygen is the most abundant element in nature — it alone takes up almost 50% by mass of all elements. See Figure 1.2.

50% 1.2

air

Figure 1.2 Oxygen is present as a free element in air, and in combined forms in water and sand.

water sand

1.2

Class practice 1.2 By referring to the Periodic Table (on the front inside cover), state which of the following substances are elements:

(

)

(a) Phosphorus

(b) Sodium chloride

(a)

(b)

(c) Ammonia

(d) Glucose

(c)

(d)

(e) Sulphuric acid

(f)

(e)

(f)

Compounds

Mercury

A1.2 Phosphorus and mercury are elements. The others are not. (Note: A substance with a name consisting of two words (e.g. sodium chloride) is not an element. A substance with a name of only one word (e.g. ammonia) may or may not be an element. The only sure way is to check the name against the Periodic Table.)

A compound is a pure substance made up of two or more elements chemically combined together.

Many common substances are compounds, such as water, common salt and sugar. Another example of a compound is copper(II) chloride. It is

(II)

made up of the elements copper and chlorine chemically combined. In this case, the word equation is: copper + chlorine

copper(II) chloride

reactants

products

+ (reacting substances)

(II)

(substance produced) (

)

(

)

3


Example 1.1 Statements about elements and compounds

1.1

This question consists of two separate statements. Decide whether each of the two statements is true or false; if both are true, then decide whether or not the second statement is a correct explanation of the first statement. ‘When two or more elements are mixed and heated, a compound is always formed.’ ‘A compound is a pure substance made up of two or more elements that are chemically combined.’ Solution The first statement is false. A compound may or may not form, depending on which elements are mixed and heated together. The second statement is true.

Decomposition of compounds We may also decompose (break down) a compound into its

(

)

constituent elements (or simpler substances) using electricity

(

(electrolysis) or heat. However, we can never decompose an

)

element chemically.

Properties of compounds compared with those of constituent elements Once formed, a compound has its own physical and chemical properties. The properties are entirely different from those of the constituent elements.

Mixture A mixture consists of two or more pure substances (elements or compounds) which have not chemically combined together.

4

(

)


A pure substance is either an element or a compound. An impure substance is always a mixture. There are three kinds of mixtures:

• element/element mixture

•

/

• element/compound mixture

•

/

• compound/compound mixture

•

/

Example 1.2 Distinguishing between elements, compounds and mixtures

1.2 (a)

(a) Give an example of (i) a pure substance which is an element. (ii) a pure substance which is a compound. (iii) an impure substance. (b) Explain why the example given in (a) (iii) is a mixture.

(i) (ii) (iii) (b)

(a) (iii)

Solution (a) (i)

Copper wire (used as electrical wire) is pure copper, an element. (ii) Distilled water is pure water, a compound. (iii) Sea water is an impure substance (impure water). (b) Sea water consists of water (a compound), sodium chloride (a compound) and other substances (compounds and elements), which have not chemically combined together.

(a) (i)

(

)

(ii)

( )

(iii)

( )

(b)

( (

➲ Try Chapter Exercise Q33

) (

) )

➲

33

Differences between mixtures and compounds If we just mix iron filings and sulphur powder, there is no heat change. We get a mixture of the two elements. Iron and sulphur still retain their original properties in the mixture.

5


Properties of iron, sulphur, iron/sulphur mixture and

1.1

iron(II) sulphide (the compound formed from iron and

(II) (

)

sulphur) are compared in Table 1.1. Note 4 Another example of element, compound and mixture that can be used to further elaborate the point (students can be involved in giving the differences between element and compound):

Property or test

Appearance

Iron

Sulphur

black solid

yellow solid

Iron/sulphur mixture

yellowish grey

Iron(II) sulphide

black solid

solid

Action of

attracted by

not attracted by

only iron attracted

not attracted by

magnet

magnet

magnet

by magnet

magnet

Action of water

sinks

most sulphur

all iron and most

sinks

sinks, while a

sulphur sink,

little floats

while a little sulphur floats

Action of dilute

liberates

no reaction

hydrochloric

hydrogen gas

acid

only iron reacts to

gives toxic

liberate hydrogen

hydrogen

gas

sulphide gas (with smell of bad eggs) ( )

Structure

Table 1.1 Comparison of properties of iron, sulphur, iron/sulphur mixture and iron(II) sulphide. N4 (II) Property or test

6

Appearance When in a fire

Hydrogen

Oxygen

Colourless gas Colourless gas Burns with a pop sound

Supports burning

Hydrogen/oxygen mixture

water

Colourless gas

Colourless liquid

Burns explosively or burns smoothly with a very hot flame

Puts out fire

Placed into a Rises in air balloon

Sinks in air

Rises or sinks depending on the composition of the mixture

Falls freely in air


Table 1.2 summarizes the main differences between

1.2

mixtures and compounds.

Mixture 1. Composition by mass

Compound

variable (the substances in the mixture

fixed

can be mixed together in any

(e.g. in water, the ratio by mass of

proportion)

hydrogen to oxygen is always 1 : 8)

( )

( 1 8)

2. Changes in formation

3. Melting point (m.p.) and boiling point (b.p.)

4. General properties

5. Separation of constituents

no chemical reaction takes place;

a chemical reaction takes place;

usually no heat change in making a

heat is usually given out or absorbed

mixture

when a compound is made

melts or boils over a wide range of

melts or boils at a definite

temperatures

temperature

(i.e. does not have a sharp m.p. or b.p.)

(i.e. has a sharp m.p. and b.p.)

(

(

)

)

each constituent substance retains its

properties are entirely different from

own properties

those of its constituent elements

constituents can be separated by

constituent elements can only be

physical methods, based on differences

separated by chemical methods, not

in physical properties

by physical methods

Table 1.2 Main differences between mixtures and compounds.

1.3

Class practice 1.3 List (a) five elements (b) five compounds and (c) five mixtures. A1.3 (a) Hydrogen, oxygen, nitrogen, iron, sulphur (b) Water, carbon dioxide, carbon monoxide, sodium chloride, iron(II) sulphide (c) Air, sea water, town gas, sodium chloride solution, wine (Other answers may be given.)

(a)

(b)

(c)

7


Classification of matter Based on what we have discussed in this chapter, we can classify matter as shown in Figure 1.3.

1.3

Matter

separation by physical methods

Mixtures

Pure substances

direct mixing

chemical decomposition

Compounds

Elements chemical combination

Figure 1.3 Classification of matter.

1.4

Properties of substances

The properties of any substance can be classified into its physical properties and chemical properties, as explained below.

Physical properties Physical properties of a substance are those properties that can be determined without the substance changing into another substance.

8

1.4


Typical physical properties include appearance (colour and

(

physical state), odour (smell), taste, hardness, density,

)

solubility (in various solvents), melting point, boilng point, N5

(

malleability (ability to be rolled into sheets), ductility (ability N6

(

to be drawn into wires), electrical conductivity and thermal conductivity.

Chemical properties

Note 5 The solubility of a solute (X) in a solvent (Y), at a given temperature, is the maximum mass (in g) of X that can dissolve in 100 g of Y at that temperature. The relationship between solubility behaviour and solubility (at 20°C) is roughly as follows: Very soluble: > 10 g Soluble: 1 – 10 g Slightly soluble: 0.01 – 1 g Insoluble: < 0.01 g

)

)

(

)

Chemical properties of a substance are the chemical reactions of the substance, and the respective conditions under which each reaction takes place.

(

For example, an effervescent tablet reacts with water quickly to release carbon dioxide. This is a chemical property of effervescent tablet.

)

Note 6 3 Gold is the most malleable element — 1 cm of gold can be rolled into a thin foil enough to cover up a football field! Gold is also the most ductile element — 1 g of gold can be drawn to 2400 m (or 1 ounce to 43 miles)!

Example 1.3 Distinguishing between physical and chemical properties of substances (a) Explain why ‘boiling point’ is regarded as a physical property. Illustrate your answer by using water as an example. (b) What is the boiling point of water at 1 atmospheric pressure? (c) State three other physical properties of water. (d) State one chemical property of water.

1.3 (a) (b) (c) (d)

(a)

Solution (a) We can determine the boiling point of water by heating water until it boils, and then measuring the temperature of the boiling water. During the measurement, liquid water changes to steam, but no new substance is formed. Note: Steam is water in gaseous state — it is still water. (b) 100°C (c) (i) Water is a colourless liquid at room conditions. (ii) Melting point of water = 0°C at 1 atmospheric pressure –3 (iii) Density of water = 1 g cm at room conditions (d) Water reacts with iron and air to form rust at room conditions.


( ) (b) 100°C (c) (i) (ii) = 0°C (iii) = 1 g cm

–3

(d)

➲

31 9


1.5


1.5

Changes can be classified as either a physical change or a chemical change.

Physical change

A physical change is a change in which no new substances are formed.

Change of state is a common example of physical change. See Figure 1.4.

1.4

heat absorbed heat given out

sublimation

solid

gas

ng ezi fre

g ltin me

Figure 1.4 Change of state is a common example of physical change.

co nd en sa tio n bo ilin g

deposition

liquid

Chemical change A chemical change is a change in which one or more new substances are formed.

Thus, the main difference between chemical and physical N7 changes is whether new substances are formed. Note 7 In some cases, classification into physical change or chemical change may not be easy (e.g. in dissolution process).

10


1.4

Class practice 1.4 State whether each of the following is a physical change or a chemical change. Give your reasons. (a) A magnesium ribbon burns in air. (b) Sugar dissolves in water. (c) Water changes to ice in a freezer. (d) Iron rusts.

1.6

A1.4 (a) Chemical change (b) Physical change (c) Physical change (d) Chemical change (b) and (c) are physical changes because no new substances are formed. (a) and (d) are chemical changes because new substances are formed.

Working in chemistry laboratory

(a) (b) (c) (d)

1.6

Observation in chemistry Observation in chemistry includes four activities:

• • • •

Seeing with eyes Feeling with hands Smelling with nose Hearing with ears

Note 8 Some students have the misconception that they need not mention the colour of a substance if it is colourless or white. Another misconception is that when no observable change is noticed, just write down ‘no observation’. Actually ‘no observable change’ is an important observation!

• • • •

Tasting with the tongue is also one way of observation, but it is not allowed in the laboratory.

Note 9 This question asks for observations. Thus it is wrong to put down something like: ‘Hydrogen is evolved from the magnesium surface.’ as it is impossible to tell whether the bubbles are hydrogen or not.

(

)

Example 1.4 Making observations in a reaction Add a small piece of magnesium ribbon to a test tube containing dilute sulphuric acid (Figure 1.5). What changes can you observe? Solution

1.4 gas bubbles

(

magnesium ribbon

dilute sulphuric acid

(1) There is effervescence — colourless gas bubbles N8 are evolved from the magnesium surface. Figure 1.5 Magnesium reacting with (2) A steamy fume is given dilute sulphuric acid.

1.5)

(1) (2) (3)

out. (3) The magnesium ribbon gradually becomes smaller in size; it eventually dissolves completely to form a Note 10 colourless solution.

N9

Another observation may be: A choking smell

(4) The test tube becomes hot. is detected. In fact hydrogen has no smell. The (5) A hissing sound is heard.

choking smell detected is due to other gases N10

(4) (5) ( (1) )

(2)

(3)

(Note: All the above are observable changes, but only (1), (2) formed by the reaction of dilute sulphuric acid and (3) are visible changes.) and the impurities present in magnesium.

11


Interpretation and prediction in chemistry For all the experiments you will do in this course, try your best to:

(1)

Observe carefully and fully when doing experiments.

(1)

(2)

Report experimental results clearly and accurately.

(2)

(3)

Analyse the results and try to interpret them. Then

(3)

draw conclusions and make predictions.

1.5

Class practice 1.5 Scientists make predictions. Which of the following predictions, do you think, are highly reliable? (a) Movement of the planets A1.5 (b) (c) (d) (e) (f)

Occurrence of a sun eclipse Weather forecast Occurrence of an earthquake Tidal movement Eruption of a volcano

(a), (b) and (e).

Laboratory safety To avoid accidents, always remember: ‘Laboratory safety is of first importance in any experimental work.’

12

(a) (b) (c) (d) (e) (f)


Basic laboratory safety rules

(1)

Do not work in the laboratory unless your teacher is

(1)

present. (2)

Follow strictly the instructions given by your teacher.

N11

(3)

Never run around or play in the laboratory. Do not

N12

leave your bench unless it is necessary. (4)

(2)

(3)

Dispose of solid waste (e.g. broken glass, filter paper, copper turnings, etc.) in the waste bin, never in the

(4)

(

sink. (5)

)

Clean up all the spillage (on the floor or bench) at once. Clean up the bench after experiment.

(6)

(5)

Report all accidents and breakages to your teacher at once.

(7)

In case any chemical gets into eyes, flush the eyes with

N13

(6)

running cold water immediately for at least three minutes. (8)

(7)

For chemical burns on skin, place the affected area 3

under slowly running cold water until the pain fades. (9)

Take all necessary safety precautions.

(8)

(9)

For more details on the safety precautions in a chemistry laboratory, refer to the ‘Laboratory Handbook’. Note 11 For example, never heat strongly when only gentle heating is required; never use concentrated acid/alkali when only a dilute one is needed; never use excessive amounts of chemicals; never do any experiment not allowed/instructed by the teacher.

Note 12 In case students are working in groups, it is a good practice to appoint a group leader, by rotation each time. He is the only person allowed to get and return apparatus and chemicals. In so doing, ‘movement’ in the laboratory can be minimized.

Note 13 Remind students that the only reliable immediate treatment for all chemical burns is washing with a lot of water.

13


Class practice 1.6

1.6

Study the following picture and point out all the improper actions that are against the rules of laboratory safety. A1.6

Hazardous chemicals Hazardous chemicals are substances which may cause injury to N14 people or damage to property. Chemicals can be classified according to their hazardous nature. A bottle containing a hazardous chemical should display the appropriate hazard warning label(s) to warn the users. Figure 1.6 shows examples of some common hazard warning labels.

Note 14 Some household chemicals have potential risks. Take the example of bleaching solution. It liberates toxic chlorine gas when mixed with an acidic substance.

Figure 1.6 Some common hazard warning labels. 14

1.6


Common laboratory apparatus Many different pieces of apparatus are required when we do experiments in the laboratory. Figure 1.7 shows some of the

1.7

common laboratory apparatus.

Flat-bottomed flask

Round-bottomed flask

Clamp and stand

Conical flask

Wire gauze

Evaporating dish (basin)

Tripod

Crucible

Pipeclay triangle

Bunsen burner

Spatula

Heat-resistant mat

Pestle

Mortar

N15

Desiccator

Test tube holder

Test tube rack

Test tube

Boiling tube

Reagent bottle

Note 15 Brown reagent bottles are for holding chemicals that decompose in the presence of light. Colourless reagent bottles are for most other chemical solutions.

Gas syringe

Measuring cylinder

Beaker

Funnel

Plastic washbottle

Safety spectacles

Figure 1.7 Common laboratory apparatus.

15


Class practice 1.7

1.7

Name the apparatus in the following figures. (c)

(b)

(i)

(h)

(q)

(j) (d)

(a)

(g)

(k) (p)

(f)

(e)

(l) (o)

(m) (n)

(t)

(u)

(v)

(y)

(z) (aa)

(s)

(bb)

(x)

(cc)

(r)

(dd)

(ee)

(ff)

(gg)

Flat-bottomed flask (a) _________________________

(l)

Crucible tongs _________________________

Reagent bottle (w) _________________________

Round-bottomed flask (b) _________________________

Spatula (m) _________________________

Gas syringe (x) _________________________

Clamp (c) _________________________

Heat-resistant mat (n) _________________________

Measuring cylinder (y) _________________________

Retort stand (d) _________________________

Pestle (o) _________________________

Beaker (z) _________________________

Conical flask (e) _________________________

Mortar (p) _________________________

Funnel (aa) _________________________

Wire gauze _________________________

Desiccator (q) _________________________

Plastic washbottle (bb) _________________________

(f)

16

(w)

Evaporating basin (g) _________________________

(r)

Test tube holder _________________________

Teat pipette (cc) _________________________

Tripod (h) _________________________

(s)

Test tube rack _________________________

Thermometer (dd) _________________________

(i)

Crucible _________________________

(t)

Test tube _________________________

Watch glass (ee) _________________________

(j)

Pipeclay triangle _________________________

Boiling tube (u) _________________________

Separating funnel (ff) _________________________

Bunsen burner (k) _________________________

Dropping bottle (v) _________________________

Glass rod (gg) _________________________


Key terms Page 1. chemical change

10

2. chemical property

9

3. chemistry

1

4. compound

3

5. element

2

6. hazard warning label

14

7. laboratory safety

12

8. mixture

4

9. observation

11

10. physical change

10

11. physical property

8

12. sublimation

10

13. word equation

3

Summary 1.1

What is chemistry about? Chemistry

1.

is a branch of science. It is the study of various substances, about their compositions, structures, properties and the changes among them.

1.2

Chemistry in our daily lives today Chemistry

2.

plays a major role in clothing, food, housing, transport and medical care.

1.3


3.

element An chemical methods.

4.

A

5.

mixture A consists of two or more pure substances (elements or compounds) which have not chemically combined.

compound

is a pure substance that cannot be broken down into anything simpler by is a pure substance made up of two or more elements chemically combined.

17


1.4 6.

Properties of substances Physical properties

of a substance are those properties that can be determined without the substance changing into another substance. Examples: appearance, colour, odour, taste, hardness, density, solubility, melting point, boiling point, malleability, ductility, electrical conductivity, thermal conductivity.

7.

Chemical properties

of a substance are the chemical reactions of the substance and the respective conditions under which each reaction occurs. For example, a chemical property of sodium is: sodium reacts with water to form sodium hydroxide.

1.5


8.

A

physical change

is a change in which no new substances are formed.

Examples: change of state, passing electricity through a light bulb. 9.

chemical change

A formed.

is a change in which one or more new substances are

Examples: burning of a candle, rusting of iron.

1.6 10.

Working in a chemistry laboratory Observation

in chemistry includes four activities:

• Seeing with eyes • Feeling with hands • Smelling with nose • Hearing with ears 11.

Laboratory safety

is of first importance in any experimental work. Refer to p.13

for basic laboratory safety rules.

18

Hazard warning labels

12.

are displayed on bottles containing hazardous chemicals, which must be handled with great care.

13.

Some common laboratory apparatus are shown in Figure 1.7 on p.15.

Chapter 2 The atmosphere

2.1


2.1

An introduction to the Earth Here are some data about the Earth: Age: 4.5 billion years Shape and size: roughly spherical, about 6400 km in radius Mass: 6 10

24

kg

6400

N1

6 10

Surface: 70% covered by water, 30% covered by land, surrounded by a gaseous layer (about 80 km thick) called the atmosphere

24

70%

30% 80

Note 1 99.99% of the total mass of the atmosphere is within 80 km of the surface of the Earth.

2.1

Class practice 2.1 (a) Why did most people in ancient times believe that the Earth’s shape was flat?

(a)

(b) Give one piece of evidence to support that the Earth is

(b)

spherical in shape. A2.1 (a) People in ancient times had little scientific knowledge. In fact, any visible portion of the Earth appeared more or less flat to the eyes. (b) Satellite photos clearly show that the Earth is roughly spherical. (Other answers may be given.)

Structure of the Earth The Earth has a layered structure The Earth consists of four layers, namely crust (5–70 km thick), mantle (about 2900 km thick), outer core and inner core. Note 2 Like the core (outer core: 2890–5150 km from surface, and inner core: 5150–6360 km from surface), the mantle can also be divided into the upper mantle (35–660 km from surface) and the lower mantle (660–2890 km from surface).

N2

(

5 – 70

2900

)

)

(

19

Part I Planet Earth

2.2

Class practice 2.2 (a) The diagram below is the structure of the Earth. Label the

(a)

different layers of the Earth. (b) Add approximately to scale, the atmosphere to the diagram and label it.

(b)

( )

A2.2 inner core outer core crust

atmosphere

(Hints: atmosphere, crust, mantle, inner core, outer core) mantle

The Earth’s crust The Earth’s crust is made up of rocks (see Chapter 4) and soils.

Planet Earth as a source of chemicals

(

)

N3

The Earth’s crust, the oceans and the atmosphere are the major sources of chemicals. Note 3 While the Earth is a very abundant source of chemicals for humans, some resources (like petroleum or some metals) are running out. Scientists are now looking into space for chemicals, including nearby planets (e.g. Mars) and satellites (e.g. the Moon). It is also hopeful that new chemicals that are not available on Earth may be found from other extra-terrestrial bodies.

2.2

The atmosphere

The atmosphere and air

Note 4 Notice that ‘atmosphere’ and ‘air’ are two similar but not identical terms. For example, we can say, ‘We breathe in air’, but we cannot say ‘We breathe in atmosphere’.

The atmosphere is a gaseous layer (about 80 km thick) surrounding the Earth. Note 5 The atmosphere has a layered structure similar to that of the Earth. From lower layer to outer layer: troposphere, stratosphere, mesosphere, thermosphere and magnetosphere.

Air is a gaseous mixture making up the atmosphere.

20

2.2 N4

(

N5

80

)

Chapter 2 Note 6 The atmosphere is like a greenhouse, making weather on Earth suitable for living things to live. The atmosphere keeps in a lot of the heat energy from the Sun The atmosphere to give the Earth a small temperature difference between day and night (unlike the Moon, where day and night temperature changes are very extreme). However, too much greenhouse gases, like CO2, in the atmosphere can cause the temperature of the atmosphere to rise, causing ice caps at the pole to melt and subsequent catastrophic consequences. CO2, if present in a suitable concentration in the atmosphere, is NOT a pollutant at all as it is needed for green

Example 2.1 Understanding the importance of the atmosphere on the Earth

2.1 N6

Explain why the atmosphere is important to life on the plants, which are the starting point of nearly all food chains. SO2 and NOx are present Earth. in the atmosphere long before there are humans on Earth. SO is produced in volcanic 2

Solution

eruptions and hot springs, and NOx is produced during rainstorm with lightning. However as volcanoes on Earth became less active, the SO2 level had dropped to a level suitable of living things to live. Human industrial activities raised these gas

The atmosphere is important because it

(1)

(1) contains a lot of free oxygen which supports life. (2) provides a moderate climate for living things to live.

(2)

(3) protects living things from the high-energy radiations from outer space. concentrations to such levels as to endanger humans and other

(3)

species on Earth, which we now call air pollution.

2.3

Class practice 2.3 1.

2.

There are 8 planets in the solar system. Is the Earth the only planet

1.

8

(a) that has an atmosphere?

(a)

(b) that has an atmosphere which can support life?

(b) 2.

Imagine you are on the Moon. Suggest why you could not find any living things there.

A2.3 1. (a) No. (7 planets have an atmosphere.) (b) Yes. 2. There is no air on the Moon.

Composition of air Air is mainly a mixture of two gases — nitrogen and oxygen (Figure 2.1).

( 2.1)

1% other gases (including argon 0.93%, carbon dioxide 0.03% and small amounts of water vapour) 1% ( 0.93% 0.03% )

oxygen 21%

Note 7 Refer to ‘Supplementary information: A simple experiment to determine the percentage of oxygen in air’ in the Teacher’s Guide.

Figure 2.1 Percentage composition by volume of clean air.

nitrogen 78% N7

21

Part I Planet Earth

We should note that air also contains small amounts of water vapour and other gases. One of them is argon. It is very unreactive and is called a noble gas.

Class practice 2.4

A2.4 Elements nitrogen, oxygen helium, neon argon, krypton xenon

Componds carbon dioxide water vapour

2.4

Air contains mainly nitrogen and oxygen. It also contains other gases such as water vapour, carbon dioxide, helium, neon, argon, krypton and xenon. Classify the constituents of air into elements and compounds.

2.3


2.3

The components of air can be separated according to their differences in boiling points. The air is first liquefied by repeated cooling and compression (Figure 2.2). The liquid air is

(

then warmed up bit by bit very slowly. Different gases in air

2.2)

boil at different temperatures, so we can collect them one by one. Nitrogen (boiling point –196°C) boils off as gas first. Argon follows (boiling point –186°C) and then oxygen (boiling point –183°C). This process is called fractional distillation of liquid air.

( –196°C) (

–186°C)

(

–183°C)

air in

liquefaction unit

nitrogen gas (b.p. –196°C) (

–196°C)

filter argon gas (b.p. –186°C) water vapour and carbon dioxide removed as solids at –80°C –80°C

(

oxygen gas (b.p. –183°C) (

air compressed and then cooled air allowed to expand — it gets very cold (–200°C) and some turns to liquid

Figure 2.2 Separation of oxygen and nitrogen from air by fractional distillation. 22

–186°C)

( –200°C )

liquid air at –200°C –200°C

–183°C)


2.5

Class practice 2.5 Boiling point (°C) (°C)

The table on the right

Gas

shows the boiling points of some of the gases

Argon

–186

found in air.

Nitrogen

–196

(a) Rearrange the gases

Neon

–246

into the order in which they would

Oxygen

–183

Xenon

–109

boil off during fractional distillation

Carbon dioxide

of liquid air.

Helium

–269

Krypton

–153

(b) List the gases which would still be gaseous at –200°C.

–78

(a)

(b)

A2.5 (a)

Boiling points of some gases.

–200°C

Helium Neon Nitrogen Argon Oxygen Krypton Xenon Carbon dioxide

–269 –246 –196 –186 –183 –153 –109 –78

(b) Neon and helium.

2.4


2.4

Physical properties of oxygen

• Oxygen is a colourless, odourless gas.

•

• It is slightly denser than air (1.1 times as dense as air).

•

(

1.1

)

• It is slightly soluble in water. •

Chemical properties of oxygen Oxygen is very reactive. It reacts with many substances to form oxides. In many reactions, so much heat is given out that the substances burn in oxygen (or air) with a flame. (

)

Test for oxygen Oxygen is a good supporter of burning (combustion), that is, it allows substances to burn in it. Put the glowing splint into a test tube containing the gas to be tested. If the gas is oxygen, the splint immediately relights N8 — that is, bursts into flame. Note 8 Sometimes a small ‘pop’ sound is heard when a glowing splint is put into a tube of oxygen. This is because when the splint relights, it gets hot quickly and the wooden splint is decomposed to give off combustible gases. The combustible gases then burn in oxygen to give a small ‘pop’ sound. This doesn’t mean the tube contains hydrogen.

(

)

23

Part I Planet Earth

Check your concept

✘ ✔

Oxygen gives a ‘pop’ sound with a burning splint. Oxygen relights a glowing splint, but does not give a ‘pop’ sound with a burning splint. Hydrogen gives a ‘pop’ sound with a burning splint, but does not relight a glowing splint.

✘

(

)

✔ (

) ( )

Oxygen relights a glowing splint. This can be used as a test for oxygen.

Key terms Page

24

1. argon

21

2. atmosphere

20

3. fractional distillation

22

4. glowing splint

23

5. noble gas

21


Summary 2.1


1.

The Earth is composed of the atmosphere surrounded by the

crust

mantle

,

,

and

.

gaseous

2.

The atmosphere is a

3.

The Earth’s crust, the oceans and the atmosphere are major sources of useful

2.2

The atmosphere

4.

The atmosphere is important because it

• contains a lot of free • provides a moderate living things • protects

core

layer surrounding the Earth.

oxygen

which supports life

climate

for living things to live

chemicals

.

from the high-energy radiations from outer space

5.

Air contains 78% by volume of nitrogen, 21% of oxygen, 0.93% of argon, 0.03% of carbon dioxide, noble gases trace amounts of other and water vapour.

2.3


6.

Nitrogen and oxygen can be obtained by

2.4


7.

Oxygen is a good supporter of

combustion

8.

We can test for oxygen with a

glowing

fractional distillation

(burning). It relights a

of liquid air.

glowing

splint.

splint.

25

Part I Planet Earth

3.1


3.1

The Earth is sometimes called a ‘water planet’. This is because

70%

70% of it is covered by water. 97% of this water occurs in oceans

97%

and seas.

3.2


3.2

Sea water is a solution containing about 3.5% by weight of dissolved substances. (In other words, there is 35 g of dissolved

3.5% (

substances in 1000 g of sea water.)

1000

35

)

A solution is a homogeneous (uniform) mixture of two or more substances.

(

)

Most of the dissolved substances are salts. The main one is ‘common salt’ — sodium chloride. The composition of sea

(

water is almost constant, although it may vary slightly

68%)

according to location and depth (Figure 3.1). ( magnesium chloride 14.6% sodium sulphate common salt (sodium chloride) ( ) 68%

11.4%

calcium chloride 3.1% other salts 2.9%

Figure 3.1 The salt composition by weight of a typical sea water sample.

26

3.1)

Chapter 3 Oceans

Check your concept

✘ ✔

3.3

All the dissolved substances in sea water are salts.

✘

There are dissolved substances in sea water other than salts. For example, gases like oxygen, carbon dioxide, organic substances like urea can also be found in sea water.

✔


3.3 (

Common salt (sodium chloride) is the most abundant resource

)

in sea water. It is an important substance, useful at home and in industry.

Evaporation of sea water Common salt can be separated from sea water by evaporation. Sea water (solution) is led into a special enclosure exposed to

(

direct sunlight. Water (solvent) evaporates and the sea water is

) (

)

becoming more and more concentrated. At some point of the process, the sea water becomes so concentrated that some salts (solute) can no longer dissolve in it. Crystals of salts appear.

(

)

The sea water at this stage is said to be a saturated solution. A saturated solution is a solution in which the solvent has dissolved the maximum amount of the solute it can at a particular temperature. Natural evaporation is a slow process. In the school laboratory, we can obtain common salt from sea water quickly by heating it to dryness. We may use either set-up, as shown in

3.2

Figure 3.2. sea water evaporating basin steam

sea water evaporating basin

water

wire gauze heat

Figure 3.2 Getting common salt from sea water in the laboratory.

heat tripod

(a) Direct heating

(b) Heating with a steam-bath 27

Part I Planet Earth

To obtain pure sodium chloride If sea water is heated to dryness as above, what is left would be a powder, not crystals. Moreover, other salts would be present besides sodium chloride. To obtain pure sodium chloride, we can use filtration followed by crystallization.

Filtration Firstly, any insoluble substances such as sand should be

(

)

removed from sea water by filtration. To filter, sea water is poured onto a piece of folded filter N1 paper in a filter funnel (Figure 3.3). A glass rod is used to guide

(

the flow (Figure 3.4). A piece of filter paper acts as a sieve in

3.3) (

3.4)

filtration. There are many tiny holes in it. These holes allow very small particles of solvent and dissolved solutes to pass through as filtrate. Larger insoluble particles remain on the filter paper as residue.

Note 1 Refer to ‘Supplementary information: Folding filter paper in fluted form’ in the Teacher’s Guide.

one layer

three layers

fold

N2

fold

filter paper filter funnel

Figure 3.3 A piece of filter paper is folded into a conical shape and placed in a funnel. Note 2 Some teachers may like to teach students to fold the filter paper in the fluted form. The fluted form of filter paper is more difficult to fold but more efficient in filtration than the conical form as shown in figure 3.7.

glass rod

sea water

folded filter paper residue filter funnel

filtrate

28

stand

Figure 3.4 Filtration of sea water.

Chapter 3 Oceans

Crystallization After the removal of insoluble impurities, pure crystals of common salt can be separated out by crystallization. Sea water is allowed to evaporate slowly at room temperature (Figure 3.5). The solution becomes more and more concentrated. Eventually, the solution becomes saturated (with N3

( 3.5)

respect to sodium chloride). Further evaporation of the solution will cause pure sodium chloride crystals to separate out. As evaporation continues, the solid crystals slowly grow in size.

Note 3 Other dissolved salts may be present in smaller amounts than sodium chloride. The solution is still not saturated as far as these salts are concerned.

sea water

more concentrated solution

solution saturated with respect to sodium chloride

sodium chloride crystals

water slowly evaporates at room temperature

Figure 3.5 Crystallization from sea water by slow evaporation.

The sodium chloride crystals can be filtered from solution and then dried by filter paper.

Isolation of pure water from sea water Figure 3.6 shows a simple set-up for distilling sea water to get

3.6

pure water. If we boil the sea water, water turns into vapour. The hot water vapour condenses back to a liquid in the cold receiver test tube.

Learning tip The solution to be distilled should contain only non-volatile solutes, those which will not easily vaporize on heating.

( )

29

Part I Planet Earth

thermometer clamp boiling tube delivery tube

receiver test tube ( )

sea water

end of delivery tube should be above the distillate

heat

Figure 3.6 Distillation of sea water using simple apparatus.

anti-bumping granule (to prevent ‘bumping’ of solution)

water (cooling agent) ( )

pure water (distillate) ( )

In distillation, the pure liquid that distils over is called the distillate. The solid left behind is called the residue. We can also carry out the above distillation using ‘Quickfit’ apparatus (Figure 3.7).

(

3.7)

Distillation involves boiling of a solution followed by condensation of the vapour formed.

thermometer

screw-cap adaptor water out (to sink) rubber tubing thermometer bulb pear-shaped flask

solution

anti-bumping granule

Liebig condenser cold water in (from tap)

heat

(a)

Figure 3.7 (a) A set of ‘Quickfit’ apparatus. (b) Distillation using ‘Quickfit’ apparatus. (a) (b) 30

receiver adaptor

test tube (as receiver)

cold water

(b)

distillate

Chapter 3 Oceans

3.4

Tests for sodium and chloride ions in common salt

3.4

Chemical analysis Chemical analysis is an important part of chemistry. It is a process to find the chemical identity or composition of a given sample.

Test for sodium ions Some metals and metal compounds, when burnt or heated strongly, produce a characteristic coloured light. A simple test — the flame test, is based on this principle. The procedure of the flame test is shown below: 1.

Moisten a clean platinum wire with concentrated N4

1.

hydrochloric acid (Figure 3.8a). 2.

(

Dip the wire into a crushed sample (or solution) of the salt

2.

to be tested (Figure 3.8b). 3.

( )

(

3.8b)

Heat the end of the wire strongly in a non-luminous flame (Figure 3.8c).

3.8a)

3.

Note 4 Often a nichrome wire is used in place of the platinum wire, which is very expensive.

non-luminous flame

(

3.8c)

flame colour due to metal ions

sample of the salt to be tested platinum wire

concentrated hydrochloric acid

(a)

(b)

(c)

Figure 3.8 Performing a flame test to identify metal ions in a sample.

31

Part I Planet Earth

By observing the colour of the flame at the wire, we can identify some types of metal ions.

The results of flame test of some metal compounds are as follows:

Compound containing

Flame colour

Potassium ion

Lilac

Sodium ion

Brilliant golden yellow

Calcium ion

Brick red

Copper(II) ion

Bluish green

(II)

Learning tip The flame colour due to potassium ion is lilac, but it is crimson when viewed through cobalt glass.

Test for chloride ions To show that chloride ions are present in sea water, we can use the following test: Silver nitrate solution is added to a sample of sea water, followed by excess dilute nitric acid. The appearance of a white precipitate (insoluble in acid) indicates the presence of chloride

(

)

ions.

Learning tip The use of excess dilute nitric acid prevents formation of other precipitates (e.g. silver carbonate, silver sulphite) which will be soluble in dilute nitric acid.

( )

The word equation for this reaction is: sodium chloride + silver nitrate

32

silver chloride + sodium (white precipitate) nitrate

+

+ (

)

Chapter 3 Oceans

3.5

Tests for the presence of water in a sample

3.5 (II)

Test by anhydrous copper(II) sulphate

Water turns white anhydrous copper(II) sulphate blue.

(II)

Test by dry cobalt chloride paper

Water turns blue dry cobalt chloride paper pink.

Learning tip Cobalt chloride test paper is also called cobalt(II) chloride test paper.

(II)

Check your concept

✘ ✔

3.6

All liquids contain water.

✘

Many liquids do not contain water. Some liquids, like oil and dry-cleaning liquid, do not mix with water.

✔


Sea water is an important source of common salt (sodium

3.6 (

)

chloride) which has many uses. Moreover, by the electrolysis of sea water, many useful products may be obtained.

33

Part I Planet Earth

Electrolysis means ‘decomposition by electricity’. It is usually carried out by passing a direct electric current through an aqueous salt solution. When sea water is electrolysed, the products are hydrogen, chlorine and sodium hydroxide. See Figure N5

3.9

3.9. electrolysis

Sea water

+

hydrogen gas + chlorine gas + sodium hydroxide solution

Note 5 During electrolysis of sea water or brine, the anodic gaseous product contains trace amount of oxygen, though the major product is chlorine.

hydrogen gas

+

chlorine gas

sea water graphite electrode (–)

graphite electrode (+)

(–)

Figure 3.9 Electrolysis of sea water.

direction of electron flow

Some uses of products from the electrolysis of brine.

Brine electrolysis

hydrogen

chlorine

Uses: as rocket fuel, make

Uses: water sanitation, make

Uses: aluminium extraction,

margarine,

bleach, plastics (e.g. PVC),

soap, paper industry, treatment

solvents, pesticides, etc.

of acidic/heavy metal effluents

ammonia

and

fertilizers, make hydrochloric acid, etc.

from factories. ( )

34

sodium hydroxide

Chapter 3 Oceans

Key terms Page 1. chemical analysis

31

2. crystallization

28

3. distillate

30

4. distillation

30

5. electrolysis

33

6. filtrate

28

7. filtration

28

8. flame test

31

9. residue

28

10. saturated solution

27

11. solute

27

12. solvent

27

35

Part I Planet Earth

Summary 3.1


1.

70% of the Earth is covered by and seas.

3.2


2.

solute Sea water contains about 3.5% by weight of dissolved substances. The main in sea water is common salt (sodium chloride). solution A is a homogeneous (uniform) mixture of two or more substances. saturated solution solvent A is a solution in which the has dissolved the maximum amount of the solute it can at a particular temperature.

3. 4.

water

. 97% of this

3.3


5.

filtration Pure common salt can be extracted from sea water by crystallization . distillation Pure water can be isolated from sea water by . condensation Distillation involves boiling of a solution followed by During distillation, the pure liquid that distils over is called the residue behind is called the .

6. 7. 8.

3.4 9. 10.

occurs in oceans

and then

of the vapour formed. ; the solid left

distillate

Tests for sodium and chloride ions in common salt Chemical analysis

is a process to find the chemical identity or composition of a given sample. flame test Some metal ions can be identified by the . Some characteristic flame colours are:

Compound containing

Flame colour Lilac

Potassium ion Sodium ion

Brilliant golden yellow

Calcium ion

Brick red

Copper(II) ion

36

water

Bluish green

11.

silver nitrate To test for chloride ions in sea water, we add solution, nitric acid followed by excess dilute , to a sample. A white silver chloride precipitate forms if chloride ions are present.

3.5

Tests for the presence of water in a sample Anhydrous copper(II) sulphate

dry cobalt chloride test paper

12.

and used to detect the presence of water in a given sample.

3.6


13.

hydrogen Electrolysis of sea water produces useful chemicals: sodium hydroxide . Refer to p.34 for their uses.

,

can be

chlorine

and

Chapter 4 Rocks and minerals

4.1

Rocks

4.1

What are rocks and minerals? In science, the word ‘rock’ has a more specific meaning:

A rock is a solid mass of a mineral or a mixture of minerals.

A mineral is a naturally occurring solid with a definite crystalline structure and chemical composition.

A mineral can be a solid element (e.g. graphite), or in most (

cases a solid compound (e.g. aluminium oxide). (

Example 4.1 Extracting minerals from rocks

(a)

(a) Is rock salt a mineral?

(b)

(

) ( )

Rock salt

Solution

(a)

(a) Yes.

(b) (i)

(b) (i)

)

4.1

Common salt (sodium chloride) is usually obtained from sea water. However, there are underground deposits of sodium chloride (called rock salt) in some countries, e.g. U.K. (b) Suggest two methods to extract rock salt from underground.

)

By mining. Use explosives to break up the rock salt. Then use trucks to carry away the rock salt.

(ii)

(ii) Pump water into the underground deposit. Rock salt dissolves in the water, leaving most other minerals behind. Then pump the salt solution up to the ground.

37

Part I Planet Earth

Uses of minerals There are more than 2200 minerals in the Earth’s crust. Most

2200

are useful for many purposes. Some uses of minerals:

• • • • •

Graphite is used to make ‘pencil lead’.

•

Rock salt is used in cooking.

•

Jade is a gemstone used for decoration and in jewellery.

•

Marble is used as floors in commercial buildings. Gold is used in jewellery.

• •

Minerals and ores Most minerals require treatment before they become useful. For example, many minerals are metal-containing compounds. Before we can use the metals, we have to extract them from their ores first.

An ore is a mineral from which a constituent (usually a metal) can be profitably extracted.

( )

Some common ores are shown below. Bauxite — the main ore of aluminium. It is mostly aluminium oxide. Copper pyrite — the main ore of copper. It is mostly copper iron sulphide.

(II)

Haematite — the main ore of iron. It is mostly iron(III) oxide. (III)

4.2


Extracting metals from ores An ore of a metal is typically a compound of the metal.

38

4.2


To obtain a pure metal from its ore, the following processes are usually involved: 1.

Mining of the ore (that is, digging the ore from the ground)

1.

2.

Concentrating the ore

2.

3.

Extraction of the metal from the concentrated ore

3.

4.

Purification of the impure metal

4.

(

)

Extraction of iron from haematite Haematite contains mainly iron(III) oxide. We can obtain iron

(III)

from the ore by heating it with coke (carbon) to a high temperature in a blast furnace. The overall reaction can be

(

)

represented by a word equation: heat

iron(III) oxide + carbon

iron + carbon dioxide

(III) +

N1

+

Extraction of aluminium from bauxite Bauxite is first treated to give pure aluminium oxide. The aluminium oxide is then electrolysed in molten state to produce aluminium metal. electrolysis

aluminium oxide

+

aluminium + oxygen

Extraction of silver from its ores Silver is an unreactive metal. It can be extracted from its ores (e.g. silver glance) by heating alone. Silver oxide decomposes

(

)

on heating to produce the metal and oxygen. A glowing splint can be used to test for any oxygen evolved. heat

silver oxide

silver + oxygen

+

Limited reserves of natural resources Natural resources such as ores are limited in amount and non- N2

(

)

renewable. It is obvious that we must use natural resources wisely, so that they can last longer. Note 1 The overall equation is a simplified one. The reactions involved are: carbon + oxygen carbon dioxide carbon dioxide + carbon carbon monoxide iron(III) oxide + carbon monoxide iron + carbon dioxide

Note 2 Ma On Shan used to be an iron ore mine until after World War II. This means that the iron ore mine in Ma On Shan was operative up to the 1950’s. As the amount of iron ores became limited, the mine closed down. This example may help students to appreciate that natural resources are non-renewable.

39

Part I Planet Earth

4.3

Limestone, chalk and marble

4.3

Rocks containing calcite Limestone, chalk and marble are common rocks. They have

(

)

one thing in common — they all contain the same mineral calcite (a crystalline form of calcium carbonate). These three

(

)

forms of naturally occurring calcium carbonate have different appearances. Their hardness also differs. Limestone is the most common form of calcium carbonate. It is hard and strong, yet inexpensive. It is therefore widely used in the building industry. Chalk is slightly softer and is also used in buildings. Marble is a crystalline solid and is very hard. It can be smoothly polished to give a beautiful appearance. It is often used for building statues, monuments, and as floors and walls in some buildings.

Uses of limestone Limestone is used as a building material. Blocks of limestone can be used to construct buildings and roads. Limestone has many other uses as well. •

Limestone is used to make footpaths.

•

•

Limestone is a raw material for making cement.

•

•

Limestone is used in neutralizing water and soil affected

•

by acid rain.

4.4


4.4

Weathering and erosion In fact, all rocks exposed on the Earth’s surface are slowly worn away by weathering and erosion.

Weathering of rocks is the slow process (usually over thousands of years) in which exposed rocks are broken down into smaller pieces.

40

( )


Weathering occurs through the actions of water, wind, air and changes in temperature. Erosion of rocks is the slow process in which weathered rock pieces are transported away by gravity, wind and water. Erosion can also have a broader meaning. It may also refer to the process which involves both weathering of rocks and transportation of weathered rock pieces to another place.

Types of weathering Rocks can be weathered in two ways:

• Physical weathering • Chemical weathering

• •

Physical weathering Weathering by temperature changes Changes in temperature can break rocks. This happens when rocks get hot in the daytime but cool down quickly at night. The effect is even much more common in deserts. Weathering by frost action Rainwater can fill cracks in rocks. When the temperature drops below 0°C, water freezes, and expands to form ice. This forces rocks to break apart. We call this frost action. See Figure 4.1.

0°C 4.1

rainwater gathers in crack

ice

eventually a piece of rock breaks off

water freezes and expands

rock the crack gets bigger

Figure 4.1 Expanding ice breaks rocks.

temperature falls below 0°C 0°C

41

Part I Planet Earth

Note 3 Rainwater containing dissolved carbon dioxide corrodes limestone areas and underground caves are resulted. Over a long time as more calcium hydrogencarbonate solution sips through cracks in cave ceilings, dissolved calcium carbonate slowly deposits and forms stalactites and stalagmites.

Chemical weathering Attack by acid

stalactite

stalagmite

Rainwater attacks rocks, especially those containing calcium N3 carbonate. It is because carbon dioxide in air dissolves slightly in rainwater, forming an acidic solution. carbon dioxide + water

carbonic acid

+ (

(acidic)

)

The carbonic acid formed reacts with calcium carbonate: calcium carbonate + carbonic acid

+

calcium hydrogencarbonate

Calcium hydrogencarbonate is soluble in water and thus the limestone is slowly worn away. The results of this natural chemical weathering process include:

•

• Formation of sinkholes in limestone areas

•

• Damage to limestone statues Attack by oxygen Oxygen in air can attack some rocks, especially those containing iron. This causes the rock to wear away slowly.

4.5

Chemical changes involving calcium carbonate

4.5

Thermal decomposition of calcium carbonate (

When limestone (mainly calcium carbonate) is gently warmed

)

with a small Bunsen flame, there seems to be no visible change. However, when it is heated strongly with a roaring non-

(

luminous flame (at about 900°C), it decomposes to give calcium oxide and carbon dioxide gas (Figure 4.2).

42

900°C) (

4.2)


test tube holder test tube holder

calcium carbonate

roaring nonluminous flame

calcium carbonate

roaring nonluminous flame

Figure 4.2 Heating calcium carbonate strongly to make quicklime (calcium oxide).

Bunsen burner

Bunsen burner

( )

strong heat

calcium carbonate

calcium oxide + carbon dioxide

+

Calcium oxide is commonly known as quicklime. When treated with water, it turns into calcium hydroxide (slaked

(

)

lime), producing a lot of heat at the same time. calcium oxide + water

+

calcium hydroxide

(

(slaked lime)

)

On stirring calcium hydroxide with water, a white suspension is formed. (Calcium hydroxide is only slightly soluble in water.) If the suspension is filtered, a clear solution called limewater is produced.

Limewater test for carbon dioxide Limewater is a saturated solution of calcium hydroxide in water. (

It is a clear colourless solution, which turns milky when carbon

4.3)

dioxide is passed through it for a few seconds (Figure 4.3). This is because the white insoluble solid particles of calcium carbonate formed are suspended throughout the mixture. calcium hydroxide + carbon dioxide (colourless solution)

calcium carbonate + water

+ (

)

+ (

)

(white solid)

Learning tip Although limewater is a saturated solution of calcium hydroxide in water, it is only a dilute solution. This is because of the low solubility of calcium hydroxide in water. 43

Part I Planet Earth

limewater (colourless solution) ( )

limewater turned milky

(a)

(b) Figure 4.3 Carbon dioxide turns limewater milky.

Carbon dioxide is a colourless gas. It turns limewater milky.

The above changes can be summarized as shown in Figure

4.4

4.4: step 1

calcium carbonate

step 4 pass in carbon dioxide (limewater test)

4

strong heat

1 limestone

(

) carbon dioxide given off

calcium hydroxide solution

calcium oxide

limewater

quicklime

step 3

step 2

add more water, stir well and then filter

add a little water

3

2

calcium hydroxide

A4.1 heat

slaked lime

Step 1: calcium carbonate calcium oxide + carbon dioxide Step 2: calcium oxide + water calcium hydroxide Figure 4.4 How some chemical changes involving calcium carbonate are related. Step 3: calcium hydroxide + water calcium hydroxide solution (limewater) Step 4: calcium hydroxide solution (limewater) + carbon dioxide calcium carbonate + water

4.1

Class practice 4.1 Write word equations for steps 1–4 in Figure 4.4.

44

4.4

1

4


4.6


4.6

/

/

Test for calcium ions Calcium compounds give a brick red flame in a flame test.

Test for carbonate ions Dilute hydrochloric acid is added to each sample (Figure 4.5). If the sample is a carbonate, carbon dioxide is produced, which turns limewater milky.

( 4.5)

calcium carbonate + hydrochloric acid calcium chloride + carbon dioxide + water +

Limestone, chalk and marble all give a positive limewater

+ +

test.

dilute hydrochloric acid

delivery tube

solid sample under test limewater

Figure 4.5 Test for a carbonate by limewater test.

Check your concept

✘ ✔

All rocks can react with dilute hydrochloric acid to give carbon dioxide.

✘

Most rocks are made up of silicates. Only those made up of carbonates can react with dilute acids to give carbon dioxide.

✔

A4.2 calcium carbonate + nitric acid

calcium nitrate + carbon dioxide + water

Class practice 4.2

4.2

Write a word equation for the reaction between chalk and dilute nitric acid. 45

Part I Planet Earth

Example 4.2 Action of dilute acid on some rocks

4.2

You are provided with distilled water, dilute hydrochloric acid (an acid stronger than carbonic acid) and small pieces of the following rocks: limestone, chalk, marble, granite (a) Which rock samples have a visible change in distilled water?

dropper

Which rock samples would give a visible change?

(a) (b)

4.6

rock sample

(b) Add dilute hydrochloric acid to the rock samples, as shown in Figure 4.6. (i)

( )

(i)

dilute hydrochloric acid watch glass

(ii) Figure 4.6 Adding dilute hydrochloric acid to a rock sample.

(iii)

(i) (i)

Write a word equation for any reaction that occurs. (ii) What do you observe in (i)? (iii) Based on the results in (i), which rocks would be attacked by carbonic acid in rainwater?

(a) (b) (i)

Solution (a) None of them. (b) (i)

(ii)

Limestone, chalk and marble. calcium carbonate + hydrochloric acid calcium chloride + carbon dioxide + water

(ii) There is effervescence. (Colourless gas bubbles are given out.) (iii) Limestone, chalk and marble.


46

( )

(iii)

➲

35


Key terms Page 1. bauxite

38

2. calcite

40

3. calcium carbonate

40

4. calcium hydroxide

43

5. chalk

40

6. erosion

41

7. haematite

38

8. limestone

40

9. limewater

43

10. marble

40

11. mineral

37

12. ore

38

13. quicklime

43

14. slaked lime

43

15. weathering

40

47

Part I Planet Earth

Summary 4.1

Rocks

1.

A

2.

mineral A is a naturally occurring solid with a definite crystalline structure and chemical composition.

3.

ore An is a mineral from which a constituent (usually a metal) can be profitably extracted. Some common ores include bauxite, copper pyrite and haematite.

4.2


4.

Extraction of a metal from its ore usually involves four steps:

• • • • 5.

rock

is a solid mass of a mineral or a mixture of minerals.

Mining

of the ore (that is, digging the ore from the ground)

Concentrating

the ore

Extraction

of the metal from the concentrated ore

Purification

of the impure metal

Two examples of extracting metals from ores: (a) Iron can be extracted from iron ores by heating

haematite

with carbon:

heat

iron(III) oxide + carbon

iron + carbon dioxide

(b) Aluminium can be extracted from

bauxite

by electrolysis:

electrolysis

aluminium oxide

4.3

Limestone, chalk and marble Limestone

6.

chalk , and same compound, calcium carbonate.

7.

Many minerals are very useful. An example is limestone. See p.40 for some of its uses.

4.4


8. 9.

48

aluminium + oxygen

marble

are different natural forms of the

Weathering

of rocks refers to the slow process (usually over thousands of years) in which exposed rocks are broken down into smaller pieces. Erosion

of rocks refers to the slow process in which weathered rock pieces are transported away by gravity, wind and water.


10.

Rocks are weathered in two ways:

• •

4.5 11.

Physical

weathering (e.g. by temperature changes, frost action)

Chemical

weathering (e.g. attack by acid, attack by oxygen in air)

Chemical changes involving calcium carbonate Rainwater

attacks rocks, especially those containing calcium carbonate:

calcium carbonate + carbonic acid 12.

calcium hydrogencarbonate

On strong heating, limestone (mainly calcium carbonate) releases calcium oxide and is changed into (quicklime).

carbon dioxide

heat

calcium carbonate 13. 14.

calcium oxide + carbon dioxide

When calcium oxide (quicklime) is treated with water, it turns into (slaked lime).

calcium hydroxide

Limewater

is a saturated solution of calcium hydroxide in water. It is a colourless solution, which turns milky when carbon dioxide is passed through it for a few seconds. This limewater test can test for carbon dioxide gas. calcium hydroxide + carbon dioxide (colourless solution)

calcium carbonate + water (

white

solid)

4.6


15.

Calcium compounds give a

16.

On treatment with dilute hydrochloric acid, calcium carbonate dissolves and releases carbon dioxide .

brick red

calcium carbonate + hydrochloric acid

colour in the flame test.

calcium chloride + carbon dioxide + water

49

HKDSE Chemistry Bridging Programe 1A

Recommend Documents