Download 2 the phases of matter

A Science Fundamental for the Pulp and Paper Industry
INTRODUCTION TO CHEMISTRY AND
PHYSICS
Facilitator Guide
NQF Level 2
Credits: 4
Unit Standard 9122
Compiled by:
Hester Oosthuizen
Johan Els
for
FIETA
Sparrow Research and Industrial Consultants © July 2005
Introductory Principles of Chemistry
and Physics
Learning Outcomes
Upon studying this module, the learner will be able to

Explain what matter is

Identify the different phases in which substances are found and explain how the
phase of a substance can be changed

Distinguish between pure substances, elements, compounds and mixtures

Describe the properties of pure substances, elements, compounds and mixtures

Explain the nature of matter and its building blocks

List a range of substances and their chemical formulae

Explain the nature and properties of cations and anions

List a range of common ions and state their oxidation state

List a range of acids and basis and their chemical compositions

Explain the properties and behaviour of acids and basis

Indicate the implication of different pH levels

Explain the chemical composition of water

Discuss the phases changes of water

Discuss the solvent characteristics of water

Explain the thermal properties of water and associated density changes

Explain what osmosis and diffusion are and the impact on nature

Explain the nature of energy and identify different forms of it

Define energy, temperature and heat

Explain the principles of heat transfer

Describe the origins of SI units

Work with different SI decimal multipliers and associated SI units
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
Use conversion factors or general principles to convert from various units and unit
basis to SI units

Explain the meaning of various physical quantities, their SI units, symbols,
applications and formulas where applicable.

Explain the principles of a filtration process

Explain the principles of a decanting process

Explain the principles of a centrifugal separation process
Unit Standard Specific Outcomes
Unit Standard 9122: Demonstrate knowledge of introductory principles of chemistry
and physics

Demonstrate knowledge of the nature of matter

Demonstrate knowledge of the nature of water

Demonstrate knowledge of temperature, energy and heat

Demonstrate knowledge of introductory principles of physics
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Table of Contents
CHAPTER 1: THE NATURE OF MATTER ............................................................................ 6
1 INTRODUCTION ............................................................................................................ 6
2 THE PHASES OF MATTER ........................................................................................... 9
3 COMPOSITION OF MATTER .......................................................................................13
3.1 Definitions ..................................................................................................................14
3.2 Pure substances ........................................................................................................15
3.3 Elements....................................................................................................................16
3.4 Compounds ...............................................................................................................22
3.5 Mixtures .....................................................................................................................23
4 ATOMS, IONS AND MOLECULES ...............................................................................24
4.1 Why do scientists believe that matter is made of particles?........................................24
4.2 Atoms ........................................................................................................................25
4.3 Ions............................................................................................................................26
4.4 Molecules ..................................................................................................................30
4.5 Acids and bases ........................................................................................................31
4.6 pH ..............................................................................................................................33
CHAPTER 2: THE NATURE OF WATER .............................................................................34
1 INTRODUCTION ...........................................................................................................34
1.1 The chemical composition of water ............................................................................34
1.2 Phases of water .........................................................................................................35
1.3 Phase changes ..........................................................................................................36
1.4 The solvent characteristics of water ...........................................................................40
1.5 The unusual thermal properties of water ....................................................................40
1.6 The change in the density of water with a change in temperature ..............................41
1.7 Osmosis and diffusion................................................................................................42
CHAPTER 3: TEMPERATURE, ENERGY & HEAT ..............................................................43
1 ENERGY .......................................................................................................................43
2 TEMPERATURE AND HEAT ........................................................................................44
CHAPTER 4: INTRODUCTORY PRINCIPLES OF PHYSICS ..............................................47
1 THE IMPORTANCE OF MEASUREMENT ....................................................................47
1.1 SI units ......................................................................................................................48
1.2 Decimal multipliers .....................................................................................................49
1.3 Converting between units ..........................................................................................50
1.4 Temperature ..............................................................................................................53
2 PHYSICAL QUANTITIES ...........................................................................................55
2.1 Length .......................................................................................................................55
2.2 Time ..........................................................................................................................56
2.3 Mass ..........................................................................................................................56
2.4 Weight .......................................................................................................................57
2.5 Gravitational force......................................................................................................58
2.6 Volume ......................................................................................................................58
2.7 Density ......................................................................................................................58
2.8 Pressure ....................................................................................................................59
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2.9 Temperature ..............................................................................................................60
CHAPTER 5: MECHANICAL SEPARATION TECHNIQUES ................................................63
1 INTRODUCTION .......................................................................................................63
2 CENTRIFUGAL SEPARATION..................................................................................64
2.1 Pressure screening ....................................................................................................64
2.2 Cyclone separators ....................................................................................................65
2.3 Centrifugal separator .................................................................................................66
3 FILTRATION..............................................................................................................67
4 DECANTING .............................................................................................................68
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CHAPTER 1: THE NATURE OF MATTER
On completion of this chapter, you should be able to:

Explain what matter is

Identify the different phases in which substances are found and explain how the
phase of a substance can be changed

Distinguish between pure substances, elements, compounds and mixtures

Describe the properties of pure substances, elements, compounds and mixtures

Explain the nature of matter and its building blocks

List a range of substances and their chemical formulae

Explain the nature and properties of cations and anions

List a range of common ions and state their oxidation state

List a range of acids and basis and their chemical compositions

Explain the properties and behaviour of acids and basis

Indicate the implication of different pH levels
1
INTRODUCTION
What is matter? Matter is defined as anything that has mass and occupies space.
Exercise 1 – Matter
Which of the following would you define as matter? Encircle all those that you think is
matter.
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If we look at the definition and apply it to the list, then the jacket, water, air, oxygen, gas
and smoke are matter because they have mass and occupy space. Even those we can’t
see, like air and oxygen occupy space and have mass.
Light and heat are not matter because they are forms of energy with no mass.
Exercise 2 - Experiment
I’m having a debate with a friend. He says you cannot weigh gases. I say you can. Who is
right?
Hot air rises because it is lighter. Helium balloons rise for the same reason. If a gas is
denser molecularly than another, you must be able to weigh it. If I put a heavy gas in a
balloon and drop it on a sensitive scale it should register some weight (minus balloon
weight).
Aim
The aim of the following demonstration is to show that gases do have weight.
We know that there is a blanket of air around the Earth called the atmosphere, and that this
results in what we call air pressure. But how can we actually prove that air has weight? Isn’t
it just an invisible mixture of gases that we need to breathe?
The air in the atmosphere is kept close to the Earth by the pull of gravity, the force that pulls
everything - including you and me - down to the ground. Without gravity, we would be
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weightless and would float above the ground, as we see with astronauts in space. As well as
giving humans weight, gravity also does the same for air. We can illustrate this with a simple
experiment using two balloons.
Instructions1
We want to test whether the balloon is heavier or lighter (or the same weight) after you blow
air into it. To do this, tie a piece of string around the middle of a stick or piece of cane so that
it balances. Then tie an empty balloon to each end of the cane. What happens?
The two balloons should balance evenly at each end.
Now remove one balloon and blow air into it. When you have done that, tie it back onto the
end of the cane. Is there any change?
That’s right; the end with the blown-up balloon should dip downwards. This is because the
air inside the balloon is making it heavier.
1 From: http://www.rcn27.dial.pipex.com/cloudsrus/pressure.html
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Matter is thus not only what we can see and feel!
2
THE PHASES OF MATTER
On a macroscopic scale, i.e. where matter is large enough to be seen, measured and
handled, matter can be classified into three main groups namely solids (s), liquids () and
gases (g). These groups are known as the phases of matter.
Definitions of the three phases:
Solid
The phase of matter in which a substance has both definite shape
and definite volume
Liquid
The phase of matter in which a substance has no definite shape but
a definite volume
Gas
The phase of matter in which a substance has no definite shape
and a volume defined only by the size of the container
Exercise 3 – Solid, liquid or gas
Classify the following materials as solid, liquid or gas and motivate your answer:
Flour is a heterogeneous mixture of grains of wheat (solid) and air (gas). Although it looks
as if flour can flow, it is not a liquid but a heterogeneous mixture of a solid and a gas.
A gold bar is a solid, because it has a definite shape and a definite volume.
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Oil is a liquid because it has no definite shape but it has a definite volume.
Smoke is not a gas but is finely divided solid particles, each small particle has a definite
shape and volume.
Steel wool is a solid because it has a definite shape and a definite volume. Although steel
wool has holes and can be compressed, it doesn’t make it a gas. In daily language, we use
the word “solid” for something without holes but in science a solid is defined differently.
A sponge is a solid, because it has a definite shape and a definite volume. Although we can
compress a sponge it doesn’t make it a gas.
Air is a gas because it has no fixed shape and the volume is determined by the container.
On the macroscopic scale we can observe the three phases of matter.
But what does it mean in terms of the sub-microscopic scale? At this scale, the particles are
too small to be seen, even with a microscope.
Scientists think that there exist sub-
microscopic particles because of the way in which matter behaves. Typical examples of submicroscopic particles are atoms and molecules, which we will discuss in the next section.
Try to answer the following questions:
Is a molecule of water in the solid state different to a molecule of water in the liquid state?
A water molecule in a block of ice is identical to a water molecule in a glass of water and to a
water molecule in water vapour.
Are the molecules arranged in the same way in ice and in water?
The difference lies in the position of these molecules relative to each other, i.e. they are
arranged differently.
The kinetic theory of matter is a model that helps us to interpret the properties of matter in
the various phases. According to this model, all matter consists of sub-microscopic particles,
i.e. atoms, molecules and ions. These particles are in constant motion.
In solids the particles (atoms, ions or molecules) are packed closely together, usually
arranged in a regular pattern. These particles vibrate back and forth about their average
positions. In a solid, a particle seldom moves past its immediate neighbour to come into
contact with a new set of particles.
The particles of liquids and gases are not arranged in a regular pattern and are not confined
to a specific location like the particles in solids. In a liquid and a gas the particles move
randomly and can move past one another. The particles of a liquid, although not fixed in
relation to each other are still closely packed together.
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Gas particles move extremely rapidly because they are not constrained by their
neighbours. They move all over the place, collide with each other and with the sides of the
container. This random motion allows the gas particles to fill the container.
Scientists draw pictures to make it easier to visualise the abstract concepts. These pictures
are symbols of what happens on the sub-microscopic level. The following figure represents
the three phases of matter.
Another important aspect of the kinetic theory is that the speed of the particles is related to
temperature. The higher the temperature the faster the particles move. The particles of
kinetic energy (energy of motion) acts to overcome the forces of attraction between the
particles. Solids melt to become a liquid when the temperature of the solid is raised to the
point at which the particles vibrate fast enough and far enough to push one another out of the
way and move out of their regularly spaced positions. As the temperature is increased, the
particles have enough energy to overcome the forces of attraction between them and escape
from the surface of the liquid into the gaseous state.
According to the model, the particles are in constant motion.
When the temperature is
decreased, the particles will move slower and slower. The motion would only stop when the
temperature has reached absolute zero, i.e. 0 K or –273ºC.
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Exercise 4 – Solid, liquid, gas
Classify the following diagrams as matter in the solid, liquid or gas phase:
Mercury
Liquid
Drop of water
Liquid
Gold
Solid
Wind
Gas
Clouds
Gas
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Steam / smoke
Gas
Empty glass
Glass seems solid but claimed to be a liquid since it
continues flowing slowly under gravity even in cold
conditions.
3
COMPOSITION OF MATTER
On the basis of its composition, matter can be divided into pure substances and mixtures,
each of them are again divided into different categories.
One way of summarising the
classification of matter is given in the following diagram:
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3.1
Definitions
Make sure you understand the following list of definitions:
Pure substance
A form of matter that cannot be separated into two different
species by any physical technique and that has a unique set of
properties
Element
Matter that is composed of only one kind of atom
Compound
Matter that is composed of two or more kinds of atoms
chemically combined in definite proportions
Mixture
A combination of two or more substances in which each
substance retains its identity
Homogeneous mixture
A mixture in which the properties are the same throughout,
regardless of the optical resolution used to examine it
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Heterogeneous mixture
A mixture in which the properties in one region or sample are
different from those in another region or sample
3.2
Pure substances
Exercise 5 – Pure substances
Encircle the pure substance(s).
Have you classified tap water and orange juice as pure? The everyday meaning of pure as
“being of natural origin”, “untampered with” and “having nothing added to it”, differs from the
scientific meaning of pure. In everyday language we say that tap water and orange juice is
pure because it is of natural origin.
Scientists say that a pure substance can not be
separated into two different substances by any physical means.
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Exercise 6 - Experiment
Do the following experiment:
Aim
To investigate whether tap water is a pure substance.
Apparatus
Tap water, sparkling clean pot, stove and knife
Procedure
Put on your safety glasses. Add tap water to the pot. Boil the water
to dryness. (Tip: Turn off the stove plate before all the water is gone.
The heat from the pot will finish the evaporation.) Scrape the pot with
the knife.
Questions
Are there any solids at the bottom of the pot?
Does the presence of solids indicate that tap water is a pure
substance or not?
When all the water has evaporated, a solid remains at the bottom of the pot. This is due to
soluble salts that were present in the water. We have used a physical method to separate
the salts from the water. Therefore according to our definitions, tap water is not a pure
substance but a homogeneous mixture.
When we look at fresh orange juice, we do not see a clear solution but we can see particles
present. Therefore orange juice is a heterogeneous mixture.
Pure substances are divided in elements and compounds.
3.3
Elements
Elements are substances that cannot be decomposed into simpler materials by
chemical means. Elements that you may be familiar with are oxygen, gold, platinum,
copper, zinc, to mention a few. Up to now, scientists have discovered 90 elements in nature
and have made another 23, giving us currently a total of 113. Most of the elements are
solids. Two elements, namely bromine and mercury are liquids. Some elements are
gases at room temperature; the more common ones are hydrogen, helium, nitrogen,
oxygen, neon, chlorine and argon.
Scientists have arranged the elements in a table which they call the periodic table (see
addendum).
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It has been said that chemists carry out experiments at the macroscopic level, but they think
about chemistry at the particulate level. They then write down their observations as symbols,
for example:
N2(g) + O2(g) → 2NO(g); ΔH = 90,2 kJ
Each element has been assigned a symbol. You might be familiar with some of them. The
symbols we use today are in most cases formed from one or two letters of the English name
for the element. For example, the symbol for carbon is C, the symbol for hydrogen is H, the
symbol for Neon is Ne and the symbol for chlorine is Cl. However, for some elements the
symbols are derived from the Latin name given to those symbols long ago. For example,
iron has the symbol Fe derived from its original Latin name Ferrum and the Latin for
potassium is Kalium, hence the symbol K.
The following table summarises the name and symbol of the most common elements and
their characteristics:
Name
(English)
(Latin)
Aluminium
Symbol
of
element
Al
Phase and colour at room
temperature, classification
Silvery solid metal
Formula
of
element
Al
Ion formed
Al3+
It is light, non-toxic (as the metal), non-magnetic and non-sparking. It is somewhat
decorative. It is easily formed, machined, and cast. Pure aluminium is soft and lacks
strength, but alloys with small amounts of copper, magnesium, silicon, manganese, and
other elements have very useful properties.
Argon
Ar
Colourless gas. Non-metal
Ar
No ion formed
Argon is colourless and odourless. Argon is very inert and is not known to form true
chemical compounds. It makes a good atmosphere for working with air-sensitive materials
since it is heavier than air and less reactive than N2.
Barium
Ba
Silvery white solid metal
Ba
Ba2+
The metal oxidises very easily and it reacts with water or alcohol. Small amounts of barium
compounds are used in paints and glasses.
Beryllium
Be
Lead grey solid metal
Be
Be2+
At ordinary temperatures, beryllium resists oxidation in air.
Boron
B
Black solid. Metalloid.
B
Boron has properties which are borderline between metals and non-metals. It is a semiconductor.
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Bromine
Br
Red-brown, liquid. Non-metal.
Br2
Br–
Bromine is the only liquid non-metallic element. It is a heavy, volatile, mobile, dangerous
reddish-brown liquid. The red vapour has a strong unpleasant odour and the vapour irritates
the eyes and throat. When spilled on the skin it produces painful sores. It is a serious health
hazard, and maximum safety precautions should be taken when handling it.
Calcium
Ca
Silvery white solid. Metal.
Ca
Ca2+
Calcium is an essential constituent of leaves, bones, teeth, and shells. Calcium does not
occur free in nature, it is found mostly as limestone, gypsum and fluorite. Stalagmites and
stalactites contain calcium carbonate.
Carbon
C
Graphite is a black solid, while
diamond is a colourless solid.
Non-metal.
C
Carbon is found free in nature in three allotropic forms: amorphous, graphite, and diamond.
Graphite is one of the softest known materials while diamond is one of the hardest. Carbon
is present as carbon dioxide in the atmosphere and dissolved in all natural waters. It is a
component of rocks as carbonates of calcium (limestone), magnesium, and iron. Coal,
petroleum, and natural gas are chiefly hydrocarbons. Carbon is unique among the elements
in the vast number of variety of compounds it can form.
Chlorine
Cl
Yellowish green gas. Non-metal.
Cl2
Cl–
Chlorine combines directly with nearly all the elements. Chlorine is a respiratory irritant. The
gas irritates the mucous membranes and the liquid burns the skin. As little as 3.5 ppm can
be detected as an odour, and 1000 ppm is likely to be fatal after a few deep breaths. It is not
found in a free state in nature, but is found commonly as NaCl (solid or seawater).
Copper
Cu
Reddish, metallic solid. Metal
Cu
Cu+: Copper (I)
Cu2+: Copper (II)
Copper is one of the most important metals. Copper is reddish with a bright metallic lustre.
It is malleable, ductile, and a good conductor of heat and electricity (second only to silver in
electrical conductivity). Its alloys, brass and bronze, are very important.
Fluorine
F
Pale yellow gas. Non-metal
F2
F–
Fluorine is the most electronegative and reactive of all elements. It is a pale yellow,
corrosive gas, which reacts with practically all organic and inorganic substances. Elemental
fluorine and the fluoride ion (in quantity) are highly toxic.
Helium
He
Un-reactive, colourless, odourless
gas. Non-metal.
He
No ion forms
The second most abundant element in the universe, after hydrogen
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Hydrogen
H
Colourless gas. Non-metal.
H2
With other nonmetals H+
With metals H–
It is by far the most abundant element in the universe
Iodine
I
Violet-dark grey, lustrous solid.
Non-metal
I2
I–
Iodine compounds are important in organic chemistry and very useful in medicine and
photography. Lack of iodine is the cause of goitre (Derbyshire neck).
Iron
Fe
(Ferrum)*
Lustrous, metallic, greyish tinged
solid. Metal.
Fe
Fe2+: Iron(II)
Fe3+: Iron(III)
Iron is a relatively abundant element in the universe. Iron is a vital constituent of plant and
animal life, and is the key component of haemoglobin.
Lead
Pb
Bluish white solid. Metal.
Pb
Pb2+
Lead is very soft, highly malleable, ductile, and a relatively poor conductor of electricity. It is
very resistant to corrosion but tarnishes upon exposure to air. Lead pipes bearing the
insignia of Roman emperors, used as drains from the baths, are still in service. Alloys
include pewter and solder. Tetraethyl lead (PbEt4) is still used in some grades of petrol
(gasoline) but is being phased out on environmental grounds.
Lithium
Li
Silvery white/grey solid. Metal
Li
Li+
A freshly cut chunk of lithium is silvery, but tarnishes in a minute or so in air to give a grey
surface. Lithium is mixed (alloyed) with aluminium and magnesium for light-weight alloys,
and is also used in batteries, some greases, some glasses, and in medicine.
Magnesium
Mg
Silvery white solid. Metal.
Mg
Mg2+
Magnesium tarnishes slightly in air, and finely divided magnesium readily ignites upon
heating in air and burns with a dazzling white flame. Normally magnesium is coated with a
layer of oxide, MgO, that protects magnesium from air and water. Magnesium is an
important element for plant and animal life. Chlorophylls are porphyrins based upon
magnesium. The adult human daily requirement of magnesium is about 0.3 g day-1.
Manganese
Mn
Silvery metallic solid. Metal.
Mn
Mn2+
The metal is gray-white, resembling iron, but is harder and very brittle. The metal is reactive
chemically, and decomposes cold water slowly. Manganese is widely distributed throughout
the animal kingdom. It is an important trace element and may be essential for utilisation of
vitamin B. It is an important component of steel.
* Latin word
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Mercury
Hg
Silvery white liquid. Metal.
Hg
Hg+: mercury(I)
Hg2+: mercury(II)
Mercury is the only common liquid metal at ordinary temperatures. Mercury is sometimes
called quicksilver. It rarely occurs free in nature. It is a rather poor conductor of heat as
compared with other metals but is a fair conductor of electricity. It alloys easily with many
metals, such as gold, silver, and tin. These alloys are called amalgams. Organic mercury
compounds are important – and dangerous. As mercury is a very volatile element,
dangerous levels are readily attained in air. It is therefore important that mercury be handled
with care. Containers of mercury should be securely covered and spillage should be
avoided. Mercury should only be handled in a well-ventilated area.
Neon
Ne
Colourless gas. Non-metal.
Ne
No ion forms
It is a very inert element. In a vacuum discharge tube, neon glows reddish orange. Liquid
neon has over 40 times more refrigerating capacity than liquid helium, and more than 3 times
that of liquid hydrogen.
Nitrogen
N
Un-reactive, colourless, odourless
gas. Non-metal.
N2
N3–
Nitrogen makes up about 78% of the atmosphere by volume. Nitrogen gas is generally inert,
however, its compounds are vital components of foods, fertilizers, and explosives.
Oxygen
O
Colourless, odourless, and
tasteless gas. Non-metal.
O2
O2–
One fifth of the atmosphere is oxygen gas. Oxygen is very reactive and oxides of most
elements are known. It is essential for respiration of all plants and animals and for most
types of combustion.
Phosphorus
P
Colourless/red/silvery white solid.
Non-metal.
P
It is an essential component of living systems and is found in nervous tissue, bones and cell
protoplasm. Phosphorus exists in several allotropic forms including white (or yellow), red,
and black (or violet). It catches fire spontaneously in air, burning to P4O10. White
phosphorous can be converted red phosphorus, which is a little less dangerous than white
phosphorus.
Potassium
(Kalium)*
K
Silvery white solid.
K
K+
Metal
Potassium is an essential constituent for plant growth and it is found in most soils. It is also a
vital element in the human diet. Potassium is never found free in nature.
* Latin word
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Silicon
Si
Dark grey solid with a bluish tinge
Si
Metalloid
Silicon makes up 25.7% of the earth's crust by weight, and is the second most abundant
element, exceeded only by oxygen. It is found largely as silicon oxides such as sand (silica),
quartz, etc. Silicon is important in plant and animal life. Diatoms in both fresh and salt water
extract silica from the water to use as a component of their cell walls. Silicon is an important
ingredient in steel. Silicon carbide is one of the most important abrasives. Workers in
environments where silicaceous dust is breathed may develop a serious lung disease known
as silicosis. Hydrolysis and condensation of substituted chlorosilanes can be used to
produce a very great number of polymeric products, or silicones. These range from liquids to
hard, glasslike solids with many useful properties. Elemental silicon has been used in lasers
to produce coherent light at 456 nm.
Silver
Ag
Silver solid.
Ag
Ag+
Metal
Silver is somewhat rare and expensive, although not as expensive as gold. Slag dumps in
Asia Minor and on islands in the Aegean Sea indicate that man learned to separate silver
from lead as early as 3000 B.C. Pure silver has a brilliant white metallic lustre. It is a little
harder than gold and is very ductile and malleable. Pure silver has the highest electrical and
thermal conductivity of all metals, and possesses the lowest contact resistance. Silver
iodide, AgI, is (or was?) used for causing clouds to produce rain.
Sodium
Na
(natrium)*
Silvery white solid.
Na
Na+
Metal
The most common compound is sodium chloride, (table salt). Soap is generally a sodium
salt of fatty acids.
Sulphur
S
A pale yellow, odourless, brittle
solid.
S
S2–
Non-metal
Sulphur is found in meteorites, volcanoes, hot springs, and as galena, gypsum, Epsom salts,
and barite. Jupiter's moon Io owes its colours to various forms of sulphur. Sulphur is
insoluble in water but soluble in carbon disulphide. Sulphur is essential to life. It is a minor
constituent of fats, body fluids, and skeletal minerals.
Zinc
Zn
A bluish-white, lustrous metal.
Zn
Zn2+
It is brittle at ambient temperatures but is malleable at 100 to 150°C. It is a reasonable
conductor of electricity, and burns in air at high red heat with evolution of white clouds of the
oxide. Zinc-deficient animals require 50% more food to gain the same weight of an animal
supplied with adequate amounts of zinc. Zinc is not particularly toxic and is an essential
element in the growth of all animals and plants. Plating thin layers of zinc on to iron or steel
is known as galvanising and helps to protect the iron from corrosion.
* Latin word
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Regardless of the origin of the symbol, the first letter is always a capital and the second
letter, if there is one, is always written in lowercase. So the symbol for chlorine is written as
Cl, not as CL. Symbols are also printed and cursive writing is not acceptable.
Exercise 7 – Elements
Identify the following elements and give their symbols:
A brittle, yellow non-metal that is commonly known as “flower of sulphur” and can be bought
at a pharmacy.
Sulphur, S
An element that is plated onto iron to prevent it from rusting.
Zinc, Zn
A rare and expensive element that was previously used to make cutlery but has now been
substituted to a great extend by stainless steel.
Silver, Ag
A gas that we can’t live without.
Oxygen, O
A silvery, liquid metal that is sometimes used in thermometers.
Mercury, Hg
3.4
Compounds
Most of the pure substances we know, like sugar, salt, water, are compounds that are
composed of two or more different elements which are always combined in the same fixed
ratio. Water is a compound consisting of hydrogen and oxygen. The ratio of hydrogen to
oxygen is always the same in water.
Table salt is a compound consisting of the elements sodium (Na) and chlorine (Cl). These
elements have reacted chemically to produce table salt (sodium chloride, NaCl).
The
elements in the compound are not present in the same form as we find them as elements.
The element sodium is a shiny grey, very reactive metal and is stored under oil to prevent it
from reacting with the oxygen in the air, it is shiny grey. Chlorine is a pale green, poisonous
gas at room temperature. Sodium chloride is not shiny like metal and doesn't have the
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characteristic properties of metals, neither is it a gas at room temperature and it is not
poisonous.
When elements react to form compounds, their individual properties are lost and in their
place we find the unique properties of the compound. The formation of a compound is a
chemical change.
3.5
Mixtures
We are very familiar with mixtures, because most materials we use everyday consist of
mixtures. Stainless steel is a mixture of the elements iron, chromium, nickel and carbon.
Coffee in is a mixture of coffee powder (which is a mixture by itself), water, sugar and milk
(which is also a mixture).
There is a difference between the way pure compounds combine to form mixtures and the
way they combine to form compounds.
When we form a compound, we said that the
elements lose their characteristics and the compound has its own characteristics.
In a
mixture each substance retains its characteristics. The sugar molecules in the bowl are the
same as the sugar molecules in the cup of coffee or the sugar molecules in a sweet
pumpkin. The only difference is that the sugar molecules in the coffee and the pumpkin are
mixed with other substances. Mixtures can have variable composition, in other words we
can have more or less of the one compound. Some people like their coffee sweeter than
other people. Compounds always have the same composition.
Mixtures can be separated by physical means, i.e. filtration, distillation, decantation, etc.
We distinguish between homogeneous mixtures and heterogeneous mixtures.
Exercise 8 – Experiment
Pour water into one glass and dissolve salt in a second glass with water. Ask a friend to
show you the salt water.
Is it possible to distinguish between the salt water and the ordinary tap water? No
Pour oil and water into a third glass. Is it possible to distinguish the tap water and the oil and
water mixture? Yes
The mixture of salt and water is called a homogeneous mixture. A homogeneous mixture
has the same properties throughout the sample. Such mixtures are often called solutions.
Solutions are not only liquids, like salt water but can also be a homogeneous gas mixture,
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like air or a homogeneous solid mixture, like brass which is a solid solution of copper and
zinc.
A heterogeneous mixture consists of two or more regions called phases that differ in
properties. Examples of heterogeneous mixtures are a mixture of ice and water; milk which
might appear smooth in texture to the unaided eye, but magnification reveals fat and protein
globules within the liquid or a mixture of oil and water.
4
ATOMS, IONS AND MOLECULES
4.1
Why do scientists believe that matter is made of particles?
John has learned in the school that scientists say matter consists of particles which cannot
be seen. He wants to know how they can say that, if they can’t see the particles.
The most convincing support that matter consists of particles is found in the behaviour of
gases. We all know that sensation when you walk into the house after a long days work and
smell the food cooking in the kitchen. How did the smell reach you?
Take a not too highly inflated inner tube of a tyre and stand on it so that the tube is flat at that
point. What happens? Can you do that if you have filled the tube with water?
Scientists have studied the behaviour of gases in detail. In the light of the experimental
facts, they have formulated a theory. The best way to explain the expansibility of gases is to
assume that matter consists of particles. Certainly, no continuous elastic material, like a
sheet of rubber, can be stretched indefinitely. If matter didn’t exist of particles, something
would have been stretched form the kitchen to the front door so that you could smell the
food. This sounds absurd, therefore scientists think that matter is not continuous but consist
of small particles that can move from the kitchen to the front door.
The fact that any gas will diffuse into a space already occupied by another gas suggests that
gases are composed of particles too small to be seen. These particles are in constant, rapid,
chaotic motion, colliding with each other and with the walls of the container. There are also
large spaces between the particles.
Diffusion not only occurs in gases but also in liquids, although the spaces between the
particles in a liquid are much smaller than the spaces between the particles of a gas.
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Exercise 9 - Experiment
Take a glass of water and very carefully add one drop of food colouring with a
medicine dropper to the water. Watch the colour spreading out till the water is coloured.
Would this have been possible if matter was continuous? No
The fact that gases may be compressed into small fractions of their original volumes lead to
the conclusion that the spaces between the particles must be large compared to the
dimensions of the particles themselves. Therefore we can stand on the tube filled with air
and compress the gas so that the tube becomes flat, we can’t do it with a tube filled with
water.
Do not confuse small bits of a substance with particles. There is a big difference between
particles of wheat and pieces of wheat. Particles refer to sub-microscopic parts while grains
or pieces refer to small bits that are still visible (maybe you have to use a microscope to see
them). In fact these visible pieces contain billions of invisible particles of the substance.
4.2
Atoms
In the previous section we saw that scientists believe that matter
consists of particles.
These particles are atoms, molecules and
ions.
An atom is the smallest particle of an element that retains the
chemical properties of that element. The atom is a neutral particle.
Each atom consists of sub-atomic particles. There are three principal
sub-atomic particles, namely protons, electrons and neutrons.
Protons are positively charged particles.
Electrons are negatively charged particles.
Neutrons are neutral particles.
The atom of each element contains a different number of protons in the nucleus. Hydrogen
(H) has 1 proton, while carbon (C) has 4 protons and oxygen has 6 protons. An atom is a
neutral particle and has the same number of electrons as there are protons in the nucleus.
So, an atom of hydrogen has 1 proton (+1) and 1 electron (-1). An atom of carbon has 4
protons (+4) and 4 electrons (-4), while an atom of oxygen has 6 protons (+6) and 6
electrons (-6). In the periodic table the number of protons in the atom of a specific element is
the same as the atomic number.
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The atom of an element is written as the symbol for that element. If I want to indicate an
atom of hydrogen, then I would write the symbol H.
4.3
Ions
Ions form when atoms lose or gain electrons. The symbol for an electron is e–.
4.3.1 Cations
Cations are usually formed when an atom of a metal loses an electron or electrons. The
atom then has more protons (positive charge) than electrons (negative charge). The result is
a positively charged ion called a cation. The symbol for an ion is the symbol for the element
with the charge on the ion indicated as a superscript. The charge is always written as the
number followed by the sign, e.g. Ca2+, Fe3+.
A sodium atom can lose 1 e– to form the sodium ion which is written as Na+.
Na atom
→
Na+ ion + e–
(11 protons + 11 electrons)
→
(11 protons + 10 electrons)
The charge on the sodium ion is +11
→
10 = +1
A magnesium atom loses 2 e– to form the magnesium ion, Mg2+ (Mg → Mg2+ + 2e–). The
magnesium atom contains 12 protons and 12 electrons. The magnesium ion contains 12
protons and 10 electrons. So, the charge on the ion is +12 – 10 = +2. Written as 2+.
How can you predict the number of electrons lost?
The periodic table (see addendum) is divided into rows and columns. The columns are
called groups. In the periodic table at the back of this module, the groups are numbered at
the top as group I, group II, etc. These are called the main group elements. In the middle,
there are 10 columns without numbers.
These elements are the so-called transition
elements.
It is easy to determine the charge on an ion of the main group elements.
The metals of groups I, II and III form positive ions with a charge equal to the group number.
For example Li, Na and K are all in group I and forms the ions Li+, Na+ and K+. The elements
of group II, i.e. Be, Mg, Ca, and Ba form the ions Be2+, Mg2+, Ca2+, and Ba2+. Aluminium in
group III form the ion Al3+.
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There is no rule to determine the charges of the ions of the transition metals, i.e. Mn, Fe, Cu,
Zn and Hg. They have to be memorised. Further, some of these metals can form more than
one ion, for example iron can form the ions Fe2+ and Fe3+.
Positive ions are named by the following rules:

For a metal cation the name of the ion is the name of the element plus the word
“ion”. For example, Na+ is referred to as the sodium ion, Al3+ is referred to as the
aluminium ion.

Some metals, especially the transition metals, can form more than one type of ion.
The charge of the ion is indicated by a Roman numeral in parenthesis directly after
the ion’s name. For example, Fe2+ is the iron(II) ion and Fe3+ is the iron(III) ion.
The following is a short list of cations and their preferred oxidation states:
Oxidation state
Element
Li+
Lithium
+
K
Potassium
Ba2+
Barium
Ca2+
Calcium
Na+
Sodium
Mg2+
Magnesium
Al3+
Aluminium
Mn2+
Manganese
Zn2+
Zink
Cr3+
Chromium
Fe2+
Iron
H+
Hydrogen
4.3.2 Anions
Non-metals gain electrons to form negative ions called anions.
A chlorine atom gains 1 electron to form the chloride ion:
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Cl + e– → Cl–
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The chlorine atom has 17 protons and 17 electrons, while the chloride ion has 17 protons
and 18 electrons (+17 – 18 = -1)
An oxygen atom gains 2 electrons to form the oxide ion:
O + 2e– → O2–
The oxygen atom has 8 protons and 8 electrons and the oxide ion has 8 protons and 10
electrons (+8 – 10 = -2).
The number of electrons gained is given by the following rule:
The non-metals of groups V, VI and VII form negative ions with a charge equal to 8 – group
number. For example nitrogen in group V forms the N3– ion (8 – 5 = 3). Oxygen and sulphur
in group VI form the ions O2– and S2– (8 – 6 = 2). The elements in group VII form the ions F–,
Cl–, Br– and I– (8 – 7 = 1).
The following is a short list of anions and their preferred oxidation states:
Oxidation state
Element
F–
Fluoride
Cl–
Chloride
O2–
Oxygen
I–
Iodide
Br–
Bromide
S2–
Sulphide
Negative ions are named by the following rules:

A mono-atomic ion (i.e. an ion consisting of one atom) is named by adding –ide to
the stem of the non-metal from which it is derived. So, the negative ion of chlorine is
called the chloride ion. Likewise we get nitride (from nitrogen), oxide (from oxygen),
fluoride (from fluorine), phosphide (from phosphorus), sulphide (from sulphur),
bromide (from bromine) and iodide (from iodine).

Polyatomic ions are ions consisting of two or more atoms. Their names must simply
be memorised. The more common polyatomic ions are given in the following table:
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Polyatomic ions
NH4+
ammonium
CO32-
carbonate
OH-
hydroxide
SO42-
sulphate
NO3-
nitrate
Cr2O72-
dichromate
NO2-
nitrite
ClO3-
chlorate
MnO4 -
permanganate
PO43-
phosphate
4.3.3 Hydrogen
Hydrogen is classified as a non-metal.
In metal hydrides the hydrogen atom gains an
–
electron to form the anion, H , like the non-metals. Hydrogen can, however, also lose an
electron to form H+, as for example in hydrochloric acid.
4.3.4 Ionic compounds
An ionic compound consists of a cation and an anion.
For an ionic compound to be
electrically neutral, that is to have no overall charge, the numbers of the positive and
negative ions must be such that the positive and negative charges balance.
Sodium chloride has sodium ions (Na+) and chloride ions (Cl–) ions. If these ions are present
in the ratio of 1:1 the total charge will be +1 + (-1) = 0. Therefore the formula of sodium
chloride is written as NaCl.
Aluminium oxide has the ions Al3+ and O2–. To have a compound with the same number of
positive charges and negative charges, the smallest common denominator is 6. So two Al3+
ions will give a total positive charge of 6 and three O2– ions will give a total negative charge
of 6. The two and the three are written as subscripts after the symbol and the formula for
aluminium oxide is Al2O3.
Magnesium hydroxide contains the ions Mg2+ and OH–. The ions will combine in the ratio of
1:2, because 1 magnesium ion will give a positive charge of 2 and two hydroxide ions will
give a negative charge of 2. Since the hydroxide ion is a polyatomic ion, we put the OH in
brackets and then the subscript, i.e. Mg(OH)2.
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When there is only one of an atom, there are no subscripts. So sodium chloride is written as
NaCl and not as Na1Cl1 and Mg(OH)2 is not written as Mg1(OH)2.
In writing an ionic compound the cation is given first and then the anion.
The name of the ionic compound is built from the names of the cation and the anion. The
cation is also named before the anion. For example, NaCl is sodium chloride, Al2O3 is called
aluminium oxide, Mg(OH)2 is called magnesium hydroxide.
4.4
Molecules
Not all atoms lose or gain electron(s) when it combines with another atom to form an ionic
compound.
Many familiar compounds, like water and sugar, are not ionic, they are
molecular. Molecules are electrically neutral and they form when atoms share electrons.
Molecules vary in size from simple molecules consisting of two atoms, e.g. nitrogen gas, N 2,
carbon monoxide gas, CO, and hydrogen gas, H2, to molecules containing millions of atoms,
e.g. plastics and living organisms.
The attractions that hold atoms together in a compound are called a bond. In molecules
these bonds are strong enough that the group of atoms making up the molecule move
together and therefore the molecule behave as a single particle.
Molecules are formed when two non-metals combine. Hydrogen forms compounds with all
the non-metals, except the noble (inert) gases (group VIII in the periodic table).
For
compounds of hydrogen with the elements of group VI and VII, the H atom is generally
written and named first in the formula (see the following table).
Formula
Name
H2O
Water
H2S
Hydrogen sulphide
HF
Hydrogen fluoride / Hydrofluoric acid
HCl
Hydrogen chloride / Hydrochloric acid
HBr
Hydrogen bromide
HI
Hydrogen iodide
The molecules formed by the elements of groups IV and V and hydrogen is written with
hydrogen last as can be seen in the following table:
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Formula
Name
NH3
Commonly known as ammonia
PH3
Commonly known as phosphine
CH4
Methane
The formula of compounds of two non-metals, not hydrogen, are generally written by putting
the elements in order of increasing group number, as illustrated in the following table:
Formula
Name
NF3
Nitrogen trifluoride
NO
Nitrogen monoxide
BCl3
Boron trichloride
Carbon has the unique property to form strong bonds between two carbon atoms. Therefore
there are millions of compounds formed by carbon. Some of the simple carbon compounds
are given in the next table:
Formula
Name
CH4
Methane
C4H10
Butane (the gas in lighters)
C2H5OH
Ethanol (the alcohol in beer and wine)
C6H6
Benzene
CCl4
Carbon tetrachoride
4.5
Acids and bases
Acids have characteristic properties:

Acids react with carbonates, like limestone, to produce bubbles of CO2 and a salt
and water

Aqueous solutions of acids conduct electricity
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
Acids change the colour of an indicator, e.g. adding lemon juice, which contains an
acid, to tea, change the colour of the tea

Dilute solutions of acids have a sour taste, e.g. vinegar and lemon juice.
In a
laboratory we never taste an acid
An acid is defined as a substance that, when dissolved in water, increases the concentration
of hydrogen ions, H+, in the water.
Some well known acids are given in the following table:
Acid
Formula
Hydrochloric acid
HCl
Hydrofluoric acid
HF
Sulphuric acid
H2SO4
Nitric acid
HNO3
Carbonic acid
H2CO3
Acetic acid
CH3COOH (Vinegar is a very dilute solution of acetic acid)
Oxalic acid
C2O4H2 (HOOCCOOH)
Bases also have characteristic properties:

Solutions of bases change the colour of an indicator

Aqueous solutions of bases conduct electricity

Bases react with acids

Aqueous solutions of bases feel soapy. Do not touch a base because concentrated
solutions of bases can cause severe burns
A base is defined as a substance that, when dissolved in water increases the concentration
of hydroxide ions, OH– in the water. Bases that dissolve in water are called alkalis.
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Some well known bases are:
Base
Formula
NaOH
Sodium hydroxide
KOH
Potassium hydroxide
NH3
Ammonia (found in cleansing agents like Handy Andy)
Acids react with bases to give a salt and water. For example
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O()
In this equation the acid is HCl where the (aq) means a watery solution. The base is NaOH
and the salt is NaCl. The word “salt” in chemistry means an ionic compound formed by the
reaction of an acid with a base. The cation comes from the base and the anion comes from
the acid, e.g. NaCl has the cation Na+, coming from the base NaOH, and the anion Cl–,
coming from the acid HCl.
The term neutralisation reaction is given to a reaction between an acid and a base.
4.6
pH
Scientists use the pH scale to indicate whether a solution is acidic or basic. The pH of a
solution is a number between 0 and 14 that indicates the degree of acidity or basicity of a
solution.

A solution with a pH less than 6 is an acidic solution

A solution with a pH of 7 is a neutral solution. Pure water has a pH of 7

A solution with a pH greater than 7 is a basic solution
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CHAPTER 2: THE NATURE OF WATER
On completion of this chapter, you should be able to:

Explain the chemical composition of water

Discuss the phases changes of water

Discuss the solvent characteristics of water

Explain the thermal properties of water and associated density changes

Explain what osmosis and diffusion are and the impact on nature
1
INTRODUCTION
Water is very common in our lives and we are so used to it that the unusual properties of
water easily escape our notice.
Exercise 10
1
Water makes up 60% of our body mass and there is a reason for it.
2
Do you know what the reason could be?
3
How long can a person survive without food?
4
How long can a person survive without water?
We can survive for more than a month without food but without fresh water we will die in a
matter of days. The water in our body keeps our body temperature constant and transports
the nutrients through our body.
In this chapter we will learn more about the properties of water.
1.1
The chemical composition of water
A water molecule consists of two hydrogen atoms and one oxygen atom. The formula of
water is H2O and its structure can be represented as follows:
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1.2
Phases of water
Water is the only chemical substance on the earth’s surface that can be found abundantly in
all three phases. Solid water is known as ice, liquid water is referred to as water and water in
the gaseous phase is known as water vapour.
Exercise 11
1
Would you freeze water in a sealed glass jar? Explain why.
It wouldn’t be a good idea to seal water in a glass jar because the jar will break.
Water expands when it freezes. In other words, the same mass of water occupies a greater
volume in the solid state than in the liquid state.
The expansion of water can be explained in terms of what happens to the molecules. In
liquid water, the molecules are free to move and the water molecules can get relatively close
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to each other. When water freezes, the water molecules arrange themselves in a six-sided
crystalline structure that contains many open spaces (see figure).
Therefore the same number of water molecules occupies a greater space when it is in ice
than when it is in water. Consequently, ice is less dense than water (d = m/V). That is why
ice floats on water.
For most substances, the solid is denser than the liquid.
This property of water for
expanding upon freezing is quite rare. If it wasn’t for this, the earth would have looked quite
different, because the ice at the poles would not have been floating on the sea but would
have sunk to the bottom of the sea.
1.3
Phase changes
When a substance goes from one phase to another, we talk about phase changes. Water
(liquid) becomes ice when it is sufficiently cooled. Water forms water vapour when the water
is heated.
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Phase changes are always accompanied by

An absorption or release of heat energy

A change in the arrangement of the particles
Melting occurs when the crystal structure of a solid collapses and the solid is converted to a
liquid. This occurs at a specific temperature, called the melting point of the substance. Heat
energy is added to a solid to melt it. The melting point of ice is 0oC.
Crystallisation is the opposite process than melting. If a liquid is cooled down enough, that
and heat energy is being removed, it becomes a solid. During crystallisation, the water
molecules that were moving randomly in the liquid become fixed in specific positions in the
crystal.
Evaporation is the change of a substance from a liquid to a vapour.
Boiling is evaporation beneath a liquid surface.
Condensation is the opposite process than evaporation. Condensation occurs either when
a vapour is cooled down to below the boiling point of the liquid, or when the pressure is
increased to above the vapour pressure of the substance with the result that the vapour then
becomes a liquid. During condensation, the water molecules that were moving randomly in
free space, now become limited to the bottom of the container or space it is located in.
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Have you ever seen dry ice or sometimes called steam ice? Although it is called ice, it is not
water.
Dry ice is solid carbon dioxide.
Carbon dioxide does not melt at ordinary
temperatures but evaporates, going directly from the liquid to the gas phase. This is called
sublimation.
Ice can also sublimate, although it is a slower process than the sublimation of dry ice. Ice
does not release water molecules as readily as liquid water does. Sublimation of ice does
however account for the loss of significant portions of snow and ice, especially on sunny, dry
mountain tops. It is also why ice cubes left in the freezer for a long time tend to get smaller.
Exercise 12
1
What happens to the water on the street after rain, or the water in an open container
if you leave it long enough?
The streets dry off and the water in the container evaporates till eventually the container is
empty.
2
How can evaporation be described in terms of the behaviour of the molecules?
The water molecules do not all have the same kinetic energy. Some water molecules have more kinetic energy and move faster than others. The fast moving water molecules can overcome the force of attraction from the other water molecules and escape from the liquid. They have evaporated and are not in the liquid any more. So even at
room temperature there are molecules that move fast enough to escape from the liquid.
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Exercise 13
The dams in South Africa have a large surface area. Taking our climate into
consideration, is this a good or a bad thing? Motivate your answer.
Large surface areas allow more evaporation to occur. If you pour a cup of water in a big pot
or long glass, the water in the pot will evaporate faster than the water in the glass. More
molecules are closer to the surface in the pot than in the glass, therefore it is easier for the
molecules in the pot to escape. South Africa, with its warm weather and shortage of water,
actually needs deeper dams to reduce the amount of evaporation. Unfortunately, due to the
landscape, it is not so easy to build deep dams.
Exercise 14 – Experiment
Boil water in a pot on the stove. Use a thermometer to measure the temperature
of the water from start till it is boiling. Once the water boils, keep it boiling for another 2
minutes.
1
What happens to the temperature? What would you look for to decide whether the
water is boiling?
Bubbles would indicate whether the water boils.
2
What is in these bubbles? Air? Water vapour? What happens when water boils?
When water is heated, the liquid water gets enough energy to move further from each other.
The molecules that have enough energy to overcome the force of attraction of other water
molecules are now in the gas phase. So inside the bubbles is water in the gas phase, i.e.
water vapour and not air that is in the bubbles. As the bubble grows, more water molecules
in the gas phase are entering the bubble and the pressure of the vapour pushes the liquid
apart, so the size of the bubble increases. The air pressure on the liquid is causing a
pressure on the bubble. If the air pressure is more than the pressure inside the bubble, the
bubble will collapse. The only way the bubble can exist is when the air pressure is equal to
the pressure inside the bubble. Now the water boils.
The temperature at which a liquid boils is defined as the temperature at which the vapour
pressure of the liquid is equal to the atmospheric pressure at that moment.
The atmospheric pressure varies, especially from place to place. The atmospheric pressure
at sea level is more than the atmospheric pressure in Gauteng. Water boils at 100oC in
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Durban and Cape Town, but due to the decrease in air pressure in Pretoria, water boils at
96oC in Pretoria.
Have you seen that once the water started boiling the temperature remained constant? A
warmer stove plate would make the water boil faster but not at a higher temperature. The
boiling point of the water depends only on the air pressure.
Exercise 15
Have you ever boiled food in a pressure cooker? How does the time to cook the food
compare to the time needed in an open pot? Have you ever wondered why this is so?
We said that the boiling point of water depends on the atmospheric pressure. In a pressure
cooker, the pot is sealed so that no water vapour can escape. This causes an increase in
the pressure inside the pot. The water inside the pot will now boil at a higher temperature
and therefore the time of cooking is reduced.
1.4
The solvent characteristics of water
Water is the most familiar solvent.
A few characteristics of water make it a very good
solvent.
First of all water is a liquid over a wide and commonly encountered temperature range, 0 to
100oC.
Secondly, water is a very good solvent for ionic compounds.
Water reacts with some covalent compounds to produce ionic solutions, for example acids
like hydrogen chloride and a base like ammonia.
1.5
The unusual thermal properties of water
The heat capacity of an object is defined as the amount of heat energy required to raise the
temperature of the object by one degree Celsius. Water has an unusually high specific heat,
higher than the values for almost all known materials.
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The following table has been included to give you an idea of the position of water relative to
some other substances:
Substance
Specific heat (in J.g-1.C-1)
Copper
0,387
Iron
0,4998
Olive oil
2,0
Water
4,1796
What does this mean practically?
We need very little heat energy to raise the temperature of copper, more to raise the
temperature of iron and a lot of heat energy to raise the temperature of water. In the kitchen,
you might have experienced it already. A copper pot on a hot plate will become warmer than
the water in the pot on the same plate.
The adult body is about 60% water by mass. An infant’s body is about 80% water by mass.
It is thus relatively easy for the human body to maintain a steady temperature of 37°C even
when the outside temperature fluctuates a lot.
The tendency of liquid water to resist change in temperature improves the climate in many
places.
For example, islands which are surrounded by water do not have the extreme
temperatures observed in the interior of a continent.
1.6
The change in the density of water with a change in
temperature
Most substances expand when the temperature of the substance is
increased. A typical example is telephone wires that sag more on a hot day
than on a cold day. We use the principle when we put a metal lid on a
glass jar that is stuck in hot water so that it expands and become loose.
The reason for the expansion of the substance is that the molecules vibrate
faster at a higher temperature and move further apart because of that.
Water above 4°C expands as it is heated because of greater molecular motion. Water
between 0° and 4°C contracts as the temperature is increased from 0°C to 4°C. This is
because water between 0° and 4°C has small ice crystals in the liquid. Upon warming, the
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ice crystals collapse, resulting in a smaller volume for the liquid water. So water is at its most
dense at 4°C.
This is of great importance in nature. In winter, as the temperature of the water drops, the
colder, less dense water moves to the bottom off the lake. This process keeps on till all the
water in the lake is at 4°C. The colder water that forms at the top is now less dense and
stays at the top. In this way, the bottom of a deep lake stays at 4°C and fish and other water
species can survive in a dam or lake even if it is frozen at the top.
1.7
Osmosis and diffusion
In living organisms, membranes of various kinds keep mixtures and solutions organised and
separated.
Yet some substances, like nutrients, have to be able to pass through
membranes. These membranes must have a selective permeability. In a process called
osmosis, the membrane only allows solvent molecules to pass through. During osmosis,
the solvent molecules move through the membrane from a solution with a lower
concentration to a solution with a higher concentration.
In Exercise 9, the food colour slowly mingled with the water till the glass was one colour.
This proved as evidence for the theory that matter consists of particles that is moving. This
is called diffusion.
Diffusion is the gradual mixing of the molecules of two or more substances by random
molecular motion.
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CHAPTER 3: TEMPERATURE, ENERGY & HEAT
On completion of this chapter, you should be able to:

Explain the nature of energy and identify different forms of it

Define energy, temperature and heat

Explain the principles of heat transfer
1
ENERGY
Energy is defined as the capacity to do work. You can do work today because you have the
energy to do so. This energy has been provided by the food you have eaten. Food energy
is chemical energy. Energy is stored in chemical compounds in the food and released when
the compounds undergo chemical reactions of metabolism in your body.
Energy can be classified as kinetic or potential energy. Kinetic energy is associated with
motion and is given by Ek = ½mv2 where m is the mass of the particle and v is the speed of
the particle. Examples of kinetic energy are:

Thermal energy. That is the energy of sub-microscopic particles,
atoms, molecules or ions, in motion. All matter has thermal energy

Mechanical energy of macroscopic objects.
For example, the
energy that a soccer ball in motion has or the energy of a moving
car

Electrical energy.
The energy transported by the electrons in a
conductor

Sound energy. The energy transported by pressure waves through a
substance.
Potential energy is the energy that results from an object’s position. Examples of potential
energy are:

Chemical potential energy. The molecules in, for example, paraffin,
have chemical potential energy.
This chemical potential energy is
converted into heat energy when the paraffin burns
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
Gravitational energy. The energy an object possess because of its
height above the ground, for example water at the top of a waterfall has
potential energy. When it falls, the potential energy of the water can be
used to generate electricity as in a hydro-electric power station.
Energy cannot be created or destroyed. It can only be changed from one form to another.
This is known as the law of conservation of energy.
When we use wood to make a fire to heat our homes, the energy which was stored in the
wood (chemical potential energy) is converted to the same amount of energy, now in the
form of heat for your home, the thermal energy of the air as well as the thermal energy of the
gases going up in the chimney, has increased. We very often say the energy of the wood
has been used up. That is not true. The energy source, that is the wood, has been used up
but the energy has been converted to other forms of energy.
Exercise 16
A battery stores chemical potential energy. Into what types of energy can this potential
energy be converted?
I can think of using a battery in a torch or a radio or turning a motor. In a torch the chemical
potential energy of the batteries is converted to light and heat (the bulb gets hot). In a radio it
is converted to sound and in the motor it is converted to work and heat.
2
TEMPERATURE AND HEAT
We have seen in Chapter 2 that according to the kinetic theory, particles in matter are in
constant motion and that the higher the temperature, the faster the particles move. Since the
kinetic energy of a particle is associated with the speed of the particle, a faster moving
particle has a higher kinetic energy. Not all particles move at the same speed. In a beaker
water at room temperature, there are some molecules that move faster and some that move
slower. If we look at that same beaker of water at 60°C there are still some molecules that
move faster than others but on average the molecules in the hot water move faster than the
molecules in the cold water. The temperature of any object is directly proportional to the
average kinetic energy of the particles of the object.
A change in temperature can be
measured by using a thermometer.
There is a difference between temperature and thermal energy. Thermal energy is the total
energy of all the atoms, molecules or ions in that object. For a given substance, its thermal
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energy depends not only on the temperature of the substance but also on the amount of the
substance. A cup full of water will have less thermal energy than a kettle full of water at the
same temperature. However, the average kinetic energy of the water in the cup and of the
water in the swimming pool is the same. A cup full of hot coffee may contain less thermal
energy than a bathtub full of warm water, even though the coffee is at a higher temperature.
Exercise 17
What happens when you put your cold hands in warm water? First explain what happens to
the temperature of the water and that of your hands. Then try to explain this in terms of the
motion of the particles of the water and that of your hands.
When you put your hands in warm water, your hands are getting warmer and the water is
getting colder. In other words, the temperature of your hands increases and the temperature
of the water decreases. In terms of the motion of the particles, if the temperature of your
hands increases, it means the average kinetic energy of the particles in your hands
increases, which means that the particles in your hands are moving faster. The temperature
of the water decreases. So the average kinetic energy of the water decreases. This means
that the particles in the water moves slower.
There is difference between temperature and heat. Heat is energy that flows from an object
at a higher temperature to an object at a lower temperature.
Exercise 18 – Experiment
Take two beakers and fill them to about a third with water. Heat the water in one beaker.
Measure the temperature of both beakers. Pour the water of the one beaker into the other
beaker. Measure the temperature immediately and then every 2 minutes till the temperature
stays constant. What happens to the temperature? How does this constant temperature
compare with the original temperature of the two beakers?
The temperature of the mixture decreases till it is constant. This constant temperature is
higher than the temperature of the cold water and lower than the temperature of the hot
water.
This simple experiment demonstrated some important concepts:
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
Heat transfer always occurs from a hotter object to a cooler object. Cold can not be
transferred. From a human perspective, if you are receiving heat, you experience
warmth, if you give away heat, you experience cooling

Transfer of heat continues till both objects are at the same temperature

The quantity of heat lost by a hotter object is equal to the quantity of heat gained by
a cooler object when they are in contact
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CHAPTER 4: INTRODUCTORY PRINCIPLES OF
PHYSICS
On completion of this chapter, you should be able to:

Describe the origins of SI units

Work with different SI decimal multipliers and associated SI units

Use conversion factors or general principles to convert from various units and unit
basis to SI units

Explain the meaning of various physical quantities, their SI units, symbols,
applications and formulas where applicable.
1
THE IMPORTANCE OF MEASUREMENT
Have you ever gone to buy a fridge without measuring the space? After all the trouble to get
it home, you discover that it is 1 cm too big. You could have measured the space by using
your arm (“From my middle finger to my elbow, plus my palm, and the length of a thumb”). If
you personally went to buy the fridge, it might have worked; but what if you wanted to order
it? You could also have used a measuring tape to measure the space, and compared it to
the measurement of the fridges at the store. But if the measurement divisions on these two
tapes were not the same, the measurement would be useless.
We therefore need to standardise measurement in order for it to be useful. All the “things”
that we can measure are called quantities which then give us quantitative information. To
measure a quantity we need to compare it to a standard unit. If we measure length, we
compare the length to the standard unit for length, which is the meter. The length is then
given as a number of times the standard unit: 5 m means that the object has a length five
times the standard unit.
All units are man-made. The kilogram is defined as the mass of a certain block of metal kept
in a laboratory in Paris. They could have made the block smaller, making the unit of mass
(kg) smaller or they could have made the block bigger, making the unit of mass larger. But,
as we all know, scientists have chosen that specific block of metal to indicate a mass of 1 kg.
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1.1
SI units
1.1.1 SI base units
The system chosen by scientists for recording and reporting measurements are called the
Système International d’Unitès (International System of Units), abbreviated as SI. It is also
known as the metric system.
All SI Units are derived from seven base units. The seven base units are given in the
following table:
Physical Quantity
Unit
Symbol
Length
meter
m
Mass
kilogram
kg
Time
second
s
Electric current
ampere
A
Temperature
kelvin
K
Amount of substance
mole
mol
Luminous intensity
candela
cd
The size of each base unit is very precisely defined. In the SI, only the kilogram is defined by
using an object (as described above). The other base units are established in terms of
“reproducible physical phenomena”.
For instance, the meter is defined as exactly the
distance light travels in a vacuum in 1/299 792 458 of a second.
1.1.2 Other common SI units
The following commonly used SI units
Physical quantity
Unit
Symbol
Length
meter
m
Mass
kilogram
kg
Time
second
s
Velocity
meter / second
m/s
Acceleration
meter / second2
m/s2
Force
Newton
N = kg·m/s2
Momentum
Newton / second
N/s
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Physical quantity
Unit
Symbol
Energy, work
joule
J = Nm
Power
watt
W = J/s
Volume
meter3
m3
Density
kilogram / meter3
kg/m3
Torque
Newton meter
Nm
Pressure
Pascal
Pa = N/m2
1.1.3 Rules for writing units

All units, when written out in full, starts with a small letter.

When symbols are used, the symbol is written with a small letter, except when the
unit is named after a scientist. In that case, the first letter of the symbol is a capital
letter. For example, the symbol for second is “s”, but the symbol for kelvin is K (after
a scientist) and the symbol for ampere is A (after another scientist).

No other abbreviations are acceptable. You should not use sec for second – it must
be s – or amp for ampere – it must be A.

When two units are multiplied, a small dot ( . ) must be written between them. This
is to differentiate between, for example, ms (millisecond) and m·s (meter second).
1.2
Decimal multipliers
It is difficult to measure very small objects and very large objects with the same unit. For
example, the meter is too large to measure the size of bacteria, while it is too small to
measure the distance between the stars. In the SI system, we form units that more closely
suit our needs by modifying the basic units with decimal multipliers. Most of us are familiar
with the kilometre (= 1 000 m) and the centimetre (= 0,01 m).
The following table give some of the most commonly used decimal multipliers:
Prefix
Symbol
Multiplication factor
tera
T
1012
giga
G
109
mega
M
106
kilo
k
103
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Prefix
Symbol
Multiplication factor
deci
d
10–1
centi
c
10–2
milli
m
10–3
micro
µ
10–6
nano
n
10–9
pico
p
10–12
The prefix “kilo” indicates 103 or 1 000.
So 1 km = 1 x 103 m (substitute k for 103) = 1 000 m
and 1 kg = 1 x 103 g = 1 000 g.
The prefix “centi” indicates 10–2.
So 1 cm = 1 x 10–2 m = 0,01 m.
When calculating physical quantities, it is often helpful to write all expressions in terms of
base SI units and scientific notation. This reduces the chance of getting confused with units.
For example, if you want to calculate the speed of an object that moved 9 Mm in 30 ms, you
cannot simply divide 9 by 30! Instead, it must first be converted to SI units (m and s):
9 x 106 m / 30 x 10-3s = 300 x 106 m/s = 300 Mm/s
Note that the answer is given as Mm/s, not m/µs (which amounts to the same thing) – this is
because a prefix is never written after the division sign. The only exception is the kilogram –
it is never written as “g”; this is because it is a base unit.
1.3
Converting between units
We often need to change units in which a physical quantity is expressed. We do so by a
method called chain-link conversion.
Step 1:
Set up a conversion factor. For example, 1 min is exactly the same time interval
as 60 s. So the conversion factors are
1 min
1
60 s
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and
60 s
1
1 min
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In a conversion factor the number and unit are treated as one, because
1
 1 is not true.
60
Including the units does make it true.
Step 2:
These ratios can now be used in your calculation because multiplying any
quantity by 1 leaves it unchanged. For example
If you introduce a conversion factor in such a way that the units do not cancel, invert the
factor and try again.
The Americans and the British do not use the metric system. So apparatus coming from
there still use the Imperial (British) system of measurement. Therefore conversions of inches
to centimetres or gallons to cubic meters are often needed. Tables with different conversion
factors are available from which you can get the conversion factor.
The instructions stated that the two points have to be 2,5 inches apart. Your ruler only
shows centimetres. From a conversion table you see that:
1 in = 2,54 cm
Reasoning: Since you have inches, you would use the conversion factor with the inches at
the bottom and the centimetres at the top so that the inches can cancel.
Calculation:
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Exercise 19
Show in the same fashion as the above example how you would do the following
conversions:
a) How many seconds are there in 5 hours?
b) How many cubic centimetres is 5 US gallon?
(1 US gallon = 3,785 litre and 1 litre = 1 000 cm3. Note that US gallons and British gallons
are not the same)
Solutions
a)
b)
Common Imperial Units and Symbols
Physical Quantity
Units
Symbols
Length
Inch
foot
mile
in, “
ft, ‘
mi
Mass
US gallon
British
pound gallon
lb
Speed
mile / hour
mi/h, mph
Force
pound
lbwt
Volume
Energy
Btu
Pressure
pound / inch2
lbwt/in2, psi
Temperature
degrees Fahrenheit
°F
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Conversion Factors
Physical Quantity
Metric Conversions
Imperial to Metric
Metric to Imperial
Length
1 km = 1000 m
1 m = 100 cm
1 m = 1000 mm
1 cm = 10 mm
1 in = 2,54 cm
1 ft = 30,48 cm
1 mi = 1,609 km
1 cm = 0,3937 in
1 km = 0,6214 mi
Volume
1 l = 1000 cm3
1 l = 1000 ml
1 US gallon = 3,785 l
1 l = 61,02 in3
1 l = 0,03532 ft3
1 m3 = 35,32 ft3
Mass
1 t* = 1000 kg
1 kg = 1000 g
* metric ton
1 lb = 0,4536 kg
1 kg = 2,2046 lb
Speed
1 km/h = 0,2778 m/s
1 mi/h = 1,609 km/h
1 km/h = 0,6214 mi/h
Density
1g/cm3 = 1000 kg/m3
Force, weight
1 t = 9,81 kN*
* on Earth
Energy
1 kW·h = 3600 kJ
1 J = 9,481 x 10-4 Btu
1 kW·h = 3413 Btu
Pressure
1 atm = 101,25 kPa
1 atm = 14,7 psi
Temperature
K = °C - 273,15
1.4
1000 kg/m3 = 62,4
lb/ft3
1 lbwt = 4,448 N
°C = (°F - 32) x 5/9
1 N = 0,2248 lbwt
1 t = 2205 lbwt
°F = (°C x 9/5) + 32
Temperature
Temperature is a property of matter, and is proportional to the average kinetic energy of the
atoms, molecules or ions present. It determines whether, and in which direction, heat energy
can be transferred from one body to another.
To measure temperature we use a thermometer. A thermometer consists of a glass tube
with a very thin tube connected to a small reservoir (the bulb at the end) filled with a liquid.
In the past mercury was used, but due to its toxicity, coloured alcohol is now used. As the
temperature increases, the liquid in the reservoir expands and its length in the column
increases.
Thermometers are calibrated in degrees centigrade (Celsius) or degrees
Fahrenheit. At atmospheric pressure, water freezes at 0°C and boils at 100°C.
The formula to convert from centigrade to Fahrenheit is:
F
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9
C  32
5
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1.
Convert 25°C to Fahrenheit.
F =
9
C + 32
5
=
9
×25 + 32
5
=
77
So 25°C is equal to 77°F
2.
What temperature in centigrade corresponds to 263 °F?
9
C + 32
5
9
C + 32
5
9
C
5
9
C
5
F
=
263
=
263 – 32
=
231
=
5
9
=
C
128,3
=
C
231×
So 263°F is equal to 128,3°C
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Exercise 20
Complete the following table:
°C
100
0
85
°F
235
260
-10
-40
Solution
°C
100
0
85
112.7
126.6
-23.3
-40
°F
212
32
185
235
260
-10
-40
2
PHYSICAL QUANTITIES
Scientists also measure physical quantities. In this unit we are going to look at the physical
quantities length, time, mass, weight, gravitational force, volume density, pressure and
temperature.
Any measurement of a physical quantity must always include a number followed by a unit
that tells us what was measured.
Activity: If I would say: The animal is 3. What do you think I was referring to?
I could have referred to his age, meaning the animal is 3 years old. Or I could have referred
to his length (3 m) or his mass (3 kg). A number without a unit is meaningless.
2.1
Length
In 1792, the newborn Republic of France established a new system of weights and
measures. It used the metre as its basis and they defined the metre as one ten-millionth of
the distance from the North Pole to the equator. For practical reasons they later redefined
the metre as the distance between two fine lines engraved on a platinum-iridium bar.
Objects can be destroyed or lost and therefore the metre is now defined as the distance that
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light travels in 1/299 792 458 of a second. This time interval was chosen so that the speed
of light is exactly 299 792 458 m.s–1.
The SI base unit for length is the metre (m). The metre is too large for
most laboratory measurements and therefore the centimetre (cm) is used
for smaller lengths.
The kilometre (km) is used for longer lengths.
These units are related to the metre as follows:
1 000 cm = 1 m
1 000 m = 1 km
2.2
Time
Time has two aspects. First of all we want to know the time of the day so that we can order
events in sequence. Secondly we want to know how long an event
lasts. A time standard has to answer to both these needs. The
earth’s rotation, which determines the length of the day, has been
used as standard. Today the second is defined by using a cesium
clock.
The SI Unit for time is the second (s) and the symbol for time is a small letter t.
2.3
Mass
The mass of an object is a measure of the quantity of matter in a given object. It is also a
measure of an object’s inertia, which is the resistance the object has to any change in its
motion
We can’t measure the mass of an object directly. We know from experience that
it is more difficult to stop an object with a large mass like a man, than an
object, moving at the same speed, with a smaller mass like a child. If we
want to measure the mass of an object, we need to measure how difficult
it is to set the object in motion or how difficult it is to change the motion
of the moving object.
The symbol for mass is a small letter m. The SI unit for mass is the
kilogram (kg). Smaller quantities are measured in g, where 1 000 g = 1
kg.
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2.4
Weight
The weight of a body is the gravitational force exerted on the object by the nearest most
massive body.
For us on earth, the nearest massive body is the
earth.
Therefore the weight of an object is the
gravitational force exerted by the earth on the object.
When the astronauts landed on the moon, the
nearest massive body was the moon, therefore the
weight of the astronauts on the moon would be the
gravitational force exerted by the moon on the
astronauts.
Since the gravitational force of the earth on an object is greater than the
gravitational force of the moon on the same objects, the weight of an object on the moon is
less than the weight on the earth. The gravitational force of the moon is about 1/6 of the
gravitational force of the earth.
The symbol for weight is a capital W and the SI unit is a derived unit called the Newton (N).
Weight = mass x gravitational acceleration
∴ 1 N = 1 kg x m.s-2 = 1 kg.m.s-2
The weight of an object on the earth can be calculated by multiplying its mass with the value
of the gravitational acceleration (g), which is 9,8 m.s-2. Thus,
Weight = mass x 9,8
or in symbolic form:
W = mg
Very often the value of g is rounded off, so that 10 m.s-2 is used.
The words mass and weight are often used interchangeably even though they are different
things. Mass refers to how much matter there is in a given object; weight refers to the force
with which the object is attracted by the earth (or other planet). The mass does not change
from place to place but the weight of an object can, depending on the magnitude of the
gravitational force. When we measure mass we call the procedure weighing. When we
weigh an object we are comparing the sample’s weight with the weights of standard masses.
When both weigh the same, their masses must be the same. This is an example where the
meaning of the word used in daily life is different to the meaning scientists couple to it. In
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daily life we would say that the weight of a person is 80 kg. The kg is not the unit for weight
but for mass.
2.5
Gravitational force
Any two objects in the universe exert a force of attraction on one another. This force is
directly proportional to the product of the masses of each other and inversely proportional to
the square of the distance between them.
We are not physically aware of these forces of attraction between two objects, except the
force between the earth and an object. The reason is that the force between two people, for
example, is too small to move either one of them. The earth has a mass of about 6,0 x 1024
kg. Since the mass of the earth is so big, it exerts a big force on objects close to it or on its
surface. Therefore objects fall towards the earth.
The SI unit for all forces is the Newton (N).
2.6
Volume
Volume (also called capacity) is a quantification of how much space
an object occupies. The SI unit for volume is the cubic metre (m3).
Volume is therefore expressed as a derived unit of length (m) and
cannot be applied to 1 dimensional (e.g. lines) or two dimensional
(e.g. surfaces) but only to three dimensional objects.
2.7
Density
Density is a measure of the mass per unit volume of a substance and can be applied to
gases, liquids and solids. This indicates that the higher an object's density is, the more mass
there is per volume of the substance. Specific volume is the inverse of density and indicates
the volume taken up by a given mass of the substance.
The average density of an object equals its total mass divided by its total volume. An object
as dense as iron has less volume than an equal mass of some less dense substance such
as water.
Density is not only dependent of the substance but is also influenced by
temperature. The densities between the same substance in gaseous, liquid and solid form
differs widely.
The symbol used for density is ρ – the Greek letter rho and the formula is:
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Where:
ρ is the object's density (in kg/m3)
m is the object's total mass (in kg)
V is the object's total volume (in m3)
The SI unit of density is therefore kilogram per cubic metre (kg/m3).
The most dense naturally occurring substance on Earth is iridium, at about is 22 650 kg/m3.
The following table shows the densities of various substances:
Substance Density in kg/m3
Iridium
22 650
Platinum
21 450
Gold
19 300
Tungsten
19 250
Mercury
13 580
Lead
11 340
Copper
8 920
Iron
7 870
Diamond
3 500
Aluminium
2 700
Seawater
1 025
Water
1 000
Ethyl alcohol
790
Gasoline
730
Air
1,2
2.8
Pressure
Pressure is the size of a force exerted to a surface, and the concentration of
that force on a given area. A finger pressed against a wall will not make
any lasting impression; but the same finger pushing a thumbtack can
easily damage the wall, even though the force applied is the same,
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because the point concentrates that force into a smaller area.
The definition of pressure (symbol: p or P) is the measure of the normal component of force
that acts on a unit area and is represented by the following formula:
Where:
p is the pressure (in Pa or the equivalent N/m2)
F is the normal component of the force (in N)
A is the area (in m2)
Pressure is sometimes measured not as an absolute pressure, but relative to atmospheric
pressure; this is normally referred to as “gauge pressure”. An example of this is the air
pressure in a tire of a car, which could be 2 atmospheres (202 650 Pa), but is actually 3
atmospheres, 2 atmospheres above atmospheric pressure plus atmospheric pressure.
In industry this “gauge pressure” (pressure above atmospheric pressure) is used widely.
Apart from Pascal, other units commonly used are kPa, atmosphere (1 atm. = 101 325 Pa),
bar (1 bar = 100 000 Pa), psi (1 pound per square inch = 14 696 Pa).
2.9
Temperature
Temperature is the property of matter that determines whether heat energy can be
transferred from one body to another, the direction of that transfer and whether it is
proportional to the average kinetic energy of the atoms, molecules or ions present.
To measure temperature we use a thermometer. A thermometer consists of a glass tube
with a very thin tube connected to a reservoir (at the end
of the thermometer) filled with a liquid.
Previously
mercury was used as liquid but due to the toxicity of
mercury they now use coloured alcohol (red).
As the
temperature increases, the liquid in the reservoir expands
and its length in the column increases. Thermometers are
calibrated in degrees Celsius or degrees Fahrenheit (the British system). The freezing point
of water if 0°C and the boiling point of water is 100°C at standard temperature and pressure.
The SI unit for temperature is the Kelvin. Kelvin temperatures are used in all calculations
which involve temperature.
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Thermometers are never marked in Kelvin, so we need to
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convert the Celsius reading to a Kelvin. The relationship between the Kelvin temperature
(symbol TK) and the Celsius temperature (symbol tC), is
TK = tC + 273,15
Notice that the size of 1 K is exactly the same as the size of 1°C. The only difference
between these two temperature scales is the zero point. On the Kelvin scale there is no
negative temperature as on the Celsius scale. The zero point on the Kelvin scale is called
absolute zero.
Exercise 21
Complete the following table:
Physical Quantity
Length
Time
Mass
Weight
Facilitator Guide
SI unit
Description
m (metre)
The metre is defined as the
distance that light travels in
1/299 792 458 of a second
s (second)
The earth’s rotation, which
determines the length of the
day, has been used as
standard.
kg (Kilogram)
The mass of an object is a
measure of the quantity of
matter in a given object. It is
also a measure of an object’s
inertia, which is the
resistance the object has to
any change in its motion
If we want to measure the
mass of an object, we need
to measure how difficult it is
to set the object in motion or
how difficult it is to change
the motion of the moving
object.
N (Newton)
The weight of a body is the
gravitational force exerted on
the object by the nearest
most massive body
Therefore the weight of an
object is the gravitational
force exerted by the earth on
the object
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Physical Quantity
Gravitational force
Volume
Density
Pressure
Temperature
Facilitator Guide
SI unit
Description
N (Newton)
Any two objects in the
universe exert a force of
attraction on one another.
This force is directly
proportional to the product of
the masses of each other
and inversely proportional to
the square of the distance
between them.
m3 (cubic metre)
Volume (also called capacity)
is a quantification of how
much space an object
occupies.
ρ (rho)
Density is a measure of the
mass per unit volume of a
substance and can be
applied to gases, liquids and
solids.
This indicates that the higher
an object's density is, the
more mass there is per
volume of the substance.
Specific volume is the
inverse of density and
indicates the volume taken
up by a given mass of the
substance.
p (Pascal)
Pressure is the size of a
force exerted to a surface,
and the concentration of that
force on a given area.
The definition of pressure
(symbol: p or P) is the
measure of the normal
component of force that acts
on a unit area and is
represented by the following
formula
TK (temperature Kelvin)
tc (Celcius temperature)
Temperature is the property
of matter that determines
whether heat energy can be
transferred from one body to
another, the direction of that
transfer and whether it is
proportional to the average
kinetic energy of the atoms,
molecules or ions present.
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CHAPTER 5: MECHANICAL SEPARATION
TECHNIQUES
On completion of this chapter, you should be able to:

Explain the principles of a filtration process

Explain the principles of a decanting process

Explain the principles of a centrifugal separation process
1
INTRODUCTION
Mechanical separation processes can be divided into three groups – sedimentation,
centrifugal separation and filtration.
In sedimentation, two immiscible (non-mixing) liquids such as oil and water, or a liquid and a
solid, differing in density, are separated by allowing them to come to equilibrium under the
action of gravity, the heavier material settling at the bottom with respect to the lighter
material. Examples of methods using this principle to separate different products include
decanters, settling tanks and separation drums.
Since the separation under gravitational force may be a slow process for certain products, it
is often speeded up by applying centrifugal forces to increase the rate of sedimentation; this
is called centrifugal separation.
Examples of methods using this principle to separate
different products include hydrocyclones and centrifugal screens.
Filtration is the separation of solids from liquids, by causing the mixture to flow through fine
pores which are small enough to stop all (or some of) the solid particles, but large enough to
allow the liquid (and small particles) to pass. Different types of sand filters and other solid
particle filters (e.g. anthracite), cloth filters and a range of pressure filters all use the same
principle to separate particles of different size.
Mechanical separation of particles from a fluid uses the forces acting on the particles as
driving force for its separation action. The forces can be direct restraining forces such as in
filtration, or indirect as in impingement filters. They can come from gravitational or centrifugal
action, which can be thought of as negative restraining forces, moving the particles relative to
the containing fluid. So the separating action depends on the character of the particle being
separated and the forces on the particle which cause the separation. The important
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characteristics of the particles are size, shape and density; and of the fluid are viscosity and
density.
The reactions of the different components to the forces set up relative motion between the
fluid and the particles, and between particles of different character. Under these different
motions, particles and fluid accumulate in different regions and can be gathered and
removed. Examples are in the filter cake and the filtrate tank in the filter press; in the
discharge valve in the base of the cyclone and the air outlet at the top; in the outlet streams
of a centrifuge.
The forces considered are gravity, combinations of gravity with other forces, centrifugal
forces, pressure forces in which the fluid is forced away from the particles, and finally total
restraint of solid particles where normally the fluid is of little consequence. The velocities of
particles moving in a fluid are important for several of these separations.
2
CENTRIFUGAL SEPARATION
Centrifugal separation describes the process used to remove solid materials included in
liquid suspensions but problematic to the use of the liquid. Centrifugal separation works on
the basis of more dense materials moving to the outside of a spinning container in
preference to lighter materials. Here the denser materials are then carefully removed, while
the light materials are removed through a central removal mechanism.
Examples of such materials to be removed by centrifugal separation are normally sand, grate
and other foreign materials.
Three common centrifugal separation processes are important:

Pressure screening

Cyclone separators

Centrifugal separators
Each of these three processes will be discussed shortly.
2.1
Pressure screening
Pressure screens work on the basis of a partial barrier being created by a perforated screen
(mostly holes or slots) which limits passage of unwanted particles. Although the principle is
similar to screening building sand through a sieve, modern pressure screens are developed
to be very selective (reject specific debris) as well as being capable of processing large
volumes due to the added pressure offered by the centrifugal force.
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In principle, particles which are smaller than the screen aperture, can either pass through the
screen or be rejected. The probability of acceptance or rejection is dependent on the relative
size of the apertures and the retention time of the particle in the screen area. Rejection of
small particles can be better understood by studying attempts to throw a tennis ball through a
set of parallel bars which are, say, one and a half times the size of the tennis ball apart.
Then try a smaller ball [e.g. a squash ball] – the tennis ball is retained by the parallel bars
more times than is expected.
Figure 0.1
2.2
A pressure screen
Cyclone separators
Products used to perform cyclone separation are generally known as hydrocyclones or
centricleaners. These cleaners are either used to remove small, dense debris (e.g. sand and
grit) from a liquid stream or separate liquid which is entrained (carried along) with a gas
stream.
Due to the injection speed and angle of the inlet feed stream, an external vortex carries
particles (or liquid in the case of gas-liquid cyclones) downward on the outside of the cleaner
vessel. A central vortex flows upwards (towards the main outlet of the vessel) and loses
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remaining particles (or liquid) in the contact area with the downward vortex.
A high
percentage of debris (or liquid) is disposed through the bottom outlet of the cyclone.
2.3
Centrifugal separator
The best known centrifugal separator is the cream separator used to separate fresh whole
milk into cream and skim milk. Formerly the separation was made by the gravity method,
allowing the cream to rise to the top of a pan and then skimming it off. The first mechanical
cream separator based on the principle of centrifugal force was produced in 1880.
Whole milk is conducted into a bowl, commonly through a central tubular shaft. A spindle
rotates the bowl at a rate of from 6,000 to 9,000 rpm, and a series of identical conical disks
separates the milk into vertical layers. The heavier skim milk collects on the outer
circumference of the rapidly whirling bowl, and the lighter cream tends to remain in the
centre. The pressure of the whole-milk supply above the bowl then forces the cream and
skim milk out of the machine and into separate collecting vessels. The cream separator
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makes it possible to control the amount of fat (called butterfat) remaining in the milk. The
gravity method ordinarily leaves one fourth of the fat in the milk, while the cream separator
leaves only 0.01% to 0.02% of the fat in the skim milk. Since the latter process is much faster
than the gravity method, there is also less chance for harmful bacterial action.
3
FILTRATION
Sand filtration is the flow of water through a bed of granular media to remove any particulate
matter.
The filter process operates on two basic principles, mechanical straining and
physical adsorption. Sand filtration is a process for separating suspended impurities from
water by passing the fluid through a bed of granular material. Water fills the pores of the
filter medium, and the impurities are adsorbed on the surface of the grains or trapped in the
openings. The key to this process is the relative grain size of the filter medium.
Sand filters need to be cleaned often, usually daily, by reversing the flow of water through the
entire filter bed, referred to as backwashing. Sand filters can also be cleaned by the total
removal of the top layer of filtration media and its contents of entrapped impurities.
In its most basic form, the filter is composed of three components, the inlet, the filter bed and
the outlet. The wastewater inflow is distributed over the media bed and then gravity pulls it
through the filter. It is then collected by the outlet component and distributed to storage
tanks or pump stations.
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4
DECANTING
Decanting is a process for the separation of mixtures, carefully pouring or removing a
solution from a container, leaving the precipitate (sediments) in the bottom of the container.
Usually a small amount of solution must be left in the container and care must be taken to
prevent a small amount of precipitate from flowing with the solution out of the container. It is
generally used to separate a liquid from an insoluble solid, e.g. potable (drinkable) water,
where the clean water is segregated from the silt and sand.
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