GRAND INTERNATIONAL SCHOOL
IGCSE Combined Science (0653)
Chemistry
C1: STATES OF MATTER
Learning objectives:
1. State the distinguishing properties of solids, liquids
and gases
2. Describe the structures of solids, liquids and gases in
terms of particle separation, arrangement and motion.
3. Describe changes of state in terms of melting, boiling,
evaporating, freezing and condensing,
4. Describe the effects of temperature and pressure on
the volume of a gas.
5. Explain changes of state in terms of kinetic particle
theory.
Introduction
o all substances that exist in our universe are made up of
matter.
o matter: any substance that has mass and occupies
space.
o Chemistry is the study of how matter behaves and how
one substance can be changed into another.
o all substances exist in three different forms (physical
states) called states of matter. These are:
1. solids (s).
3. gases (g).
2. liquids (l).
s, l, g are called state symbols
o the three physical states show the differences in the
way they respond to changes in temperature and
pressure.
o The kinetic theory of matter: helps to explain the way in
which matter behaves.
o the theory states that:
1. all matter is made up of tiny, moving particles, invisible
to the naked eye. that are in random molecular motion.
Different substances have different types of particles
(atoms, molecules or ions) which have different sizes.
2. The particles move all the time. The higher the
temperature, the faster they move on average.
3. Heavier particles move more slowly than lighter ones at
a given temperature.
o The kinetic theory can be used as a scientific model to
explain how the arrangement of particles relates to the
physical properties of matter in terms of the movement
of its constituent particles and different energies
involved.
1.1: Structure and properties of solids, liquids and
gases
A: Solids
o fixed volume and definite shape at given temperature.
o particles tightly arranged together in
a fixed position due to stronger
intermolecular forces.
o high density.
o does not flow.
o they are incompressible.
o negligible/no space between particles.
B: Liquids
o has a fixed volume, but no definite shape and size.
o take shape of container into which it is poured.
o particles close to each, but can slide past each other
easily due to weaker intermolecular
forces which makes liquids flow easily.
o moderate to high density.
o more compressible in nature than solids
and less compressible than gases.
o more space between particles
than solids.
C: Gases
o no fixed volume, shape or size.
o spread out (expands) to fill container/space they are
contained in, due to very weak intermolecular forces.
o low density.
o more compressible in nature than solids
and liquids.
o more space between particles
than solids and liquids.
Summary: Flow chart of the nature of matter
Differences in properties of states of matter.
Changes in physical state
o changes in temperature and pressure leads to
physical changes in matter that are more than
contraction or expansion.
o at atmospheric pressure, changes happen by either
lowering or raising temperature.
1. Melting and freezing
o melting point (m.p): temperature at which a
substance (solid) turns into a liquid at a particular
temperature.
o freezing point (f.p): reverse process of melting point
- temperature at which a liquid changes to a solid.
- freezing happens at the same temperature with
melting (melting and freezing of pure water takes
place at °C,
example: Gallium is a metal with melting point slightly
above room temperature and causes it to melt in a
person’s hand
2. Sublimation
o the direct change of state from solid to gas or gas to solid
through bypassing the liquid phase.
o examples: few solids, such as solid carbon dioxide, do not
melt when they are heated at normal pressures. Instead,
they turn directly into gas. Solid carbon dioxide is often called
‘dry ice’ because the surface of the block is dry.
o iodine sublimes and produces a purple vapour but
condenses again on a cold surface.
o this is different to a normal
ice cube, which has a thin
film of liquid water on the surface.
3. Evaporation, boiling and condensation
o evaporation: a process occurring at the surface of a
liquid, involving the change of state from a liquid into a
vapour at a temperature below the boiling point.
o liquid with surface exposed to the air evaporates.
o larger surface area, speeds up evaporation.
o warmer liquids, evaporates faster.
o hot climate in Dead Sea means water evaporates
easily and the sea has high salt concentration
boiling: change of state in which at a certain
temperature, the liquid becomes hot enough to escape
from liquid (body/surface) and change to gas molecules.
boiling point: the temperature at which a liquid boils,
when the pressure of the gas created above the liquid
equals atmospheric pressure.
volatility: the property of how easily a liquid evaporates
water evaporates easily, has low boiling point (100 °C).
water is a volatile liquid. Ethanol, with a boiling point of
78 °C, is more volatile than water. It has a higher
volatility than water and evaporates more easily
condensation: the change of a vapour or a gas into a
liquid and during this process, heat is given out to the
surroundings (cooling happens).
- it is the reverse process of evaporation.
- gas state most affected by changes in pressure. It is
possible, at normal temperatures, to condense a gas
into a liquid by increasing the pressure, without cooling.
Consider a kettle shown below.
o colourless, invisible water vapour escapes from the
kettle.
o water vapour is present in the clear region we can
see at the mouth of the kettle.
o visible cloud of steam is made up of droplets of liquid
o water formed by condensation as vapour cools in air.
o a beaker of boiling water, form bubbles when there
are enough high-energy water molecules to give a
gas with a pressure equal to atmospheric pressure.
o boiling point of a liquid change if surrounding
pressure changes.
o value given for boiling point is stated at the pressure
of the atmosphere at sea level (atmospheric pressure
or standard pressure).
o if surrounding pressure falls, boiling point falls.
o boiling point of water at standard pressure is 100 °C.
o on high mountain, boiling point is lower than 100 °C.
o if surrounding pressure is increased, the boiling point
rises.
o Photocopy pages 7 & 8:
Chemistry for Cambridge IGCSE COURSEBOOK
Heating and cooling curves
o melting point of solids can be measured using the
apparatus shown.
o powdered solid put in narrow melting-point tube for
easy heating.
o water bath used to measure melting points ˂ 100 °C.
o oil bath used to measure melting points ˃ 100 °C.
o on heating, temperature rises and stays constant until
all solid has melted.
o temperature then rises as liquid warms further.
o continuous heating of liquid in same apparatus can be
done to reach boiling point.
o temperature stays same until all liquid has evaporated.
o experiment can be performed in reverse using similar
apparatus to produce a cooling curve, but
thermometer placed in test-tube with solid studied.
o solid is then melted completely and liquid heated.
o heating is then stopped, temperature is noted every
minute as substance cools. This produces a cooling
curve (Figure 1.10). The level (horizontal) part of the
curve occurs where liquid freezes, forming solid.
o experiments show that heat energy is needed to
change a solid into a liquid, or a liquid into a gas.
o During the reverse processes, heat energy given out.
o Photocopy pages 9,10 & 11: EXPERIMENTAL SKILLS 1.1
Chemistry for Cambridge IGCSE COURSEBOOK
1.2 Kinetic particle theory of matter
Existence of atoms and molecules
o elements, compounds react to produce substances
around world (e.g. massive objects – gas giants such
as Jupiter, Saturn).
o matter made up of very small particles called atoms.
o key ideas:
1. each element composed of its own type of atoms.
2. atoms of different elements combine to makes
molecules of compounds
Main points of the kinetic particle theory
o matter made of small particles (different substances
have different particles, atoms, molecules, ions).
o particles in random motion (increase in temperature,
the higher the average energy of particles).
o movement and arrangement of particles is different
for the three states of matter.
o pressure of gas produced by atoms/molecules of gas
hitting the walls increase as more particles collide with
the walls.
Summary of particle arrangement
o Refer to properties of solids, liquids, gases (1.1).
o fluid properties are produced due to ability of particles
to move in the liquid and gas phases.
o particles: widely spread in gas, close together in
liquid or gas.
o intermolecular space: space between atoms or
molecules in a gas or liquid.
o intermolecular spaces larger in gases and reduced by
increasing external pressure (gases compressible).
o intermolecular spaces smaller in liquids (liquids not
very compressible).
Pressure, temperature and volume changes
1. pressure and volume (Boyle's Law)
o increase in external pressure decrease volume, gas
contracts reduces in volume.
o decrease in external pressure increase volume, gas
expands, increases volume.
Boyle’s Law: the volume of a given mass of a gas is
inversely related to its pressure when its temperature is
kept constant.
2. Volume and temperature (Charles' Law)
o increase in temperature of gas increases the volume
of gas. Gas expands.
o decrease in temperature of gas decreases volume, of
gas. Gas contracts.
Charle’s Law: The volume of a given fixed mass of a dry
gas is directly proportional to its absolute temperature at
a constant pressure.
Summary on Gas Laws
o volume of gas easily changed by
conditions of temperature and pressure
because space between rapidly moving
particles is greater than liquids and solids.
o temperature increase: gas particles
moves faster, freely, less interaction,
greater volume.
o temperature lowered: particles move
slowly, more likely interact with each
other, move together to occupy small
volume.
o increase in pressure: particles close
together causing them to interact more.
o decrease in pressure: particles occupy
greater space, interaction between
particles is less.
Differences between evaporation and boiling
Evaporation
Boiling
o
happens at all positions in the liquid at
once because all particles have
enough kinetic energy to move faster
and turn to a gas.
happens at surface of liquid
because not all but only some
particles have enough kinetic
energy to move faster and leave
surface of liquid and turn into a
gas.
o happens at lower temperatures/any happens at higher/specific/fixed
temperature
temperatures called boiling points
o temperature may changes
temperature remains constant
o slow process
quick process
o no bubbles formed
bubbles formed
Interpretation of a cooling curve
o arrangement of particles in states of
matter, help to explain energy changes
involved when a substance is heated/
cooled.
cooling of gas (region A):
o temperature falling, gas cools.
o energy of particles decreasing.
o particles move slowly, interact more
strongly, become close to form liquid.
o Intermolecular forces between particles
increase, heat given out (exothermic).
o results in temperature becoming constant
until gas condenses to liquid.
cooling of liquid (region B):
o temperature falling, liquid cools.
o energy of particles decreasing.
o particles more slow, strong interaction,
liquid becomes solid.
o Intermolecular forces between particles
increase, heat is given (exothermic).
o solid forming, release of energy keeps
temperature constant until freezing ends
formation of solid (region C):
o after solid forms, temperature falls again.
o particles in solid vibrate less strongly as
temperature falls.
Key points (freezing and condensation)
1. when particles come closer together, new
forces of interaction occur.
2. energy is given out.
3. temperature constant until liquid or solid
is totally formed.
NB: freezing/condensation are exothermic
reactions
Heating curve
o when an experiment is carried out in the
opposite direction, from solid to liquid
forms a heating curve.
o temperature constant during melting/
boiling.
o energy absorbed to break bonds
/overcome forces between particles for
free movement (endothermic)
Key words:
o intermolecular forces: the attractive forces
that act between molecules.
o exothermic changes: process/chemical
reaction in which heat energy is produced and
releases to surroundings. Enthalpy for this
change is negative: o endothermic changes: process/chemical
reaction that takes heat from the
surroundings. Enthalpy for this change is
positive.
1.3 Mixtures of substances and diffusion
o chemical world is complex: made up of a
mixture of different substances joined together.
o mixture – two or more substances joined
together, not chemically bonded but separated
by physical means.
- no chemical reaction (no new substance).
- physical change (substance formed change
from one state to another).
- the three states can combine in many ways:
e.g states completely mix to make one
state/phase
o solution – mixture formed when a solute (solid
substance) is dissolved in a solvent (solvent).
o example: solid salt completely dissolve in liquid
water to produce a liquid mixture called salt
solution.
o solute – a solid substance that dissolves in the
liquid (solvent) to form a solution.
o solvent – a liquid substance that dissolves the
solid (solute) to form a solution.
o solution = solute + solvent
o water is a common solvent, liquids that act as
solvents in organic chemistry: organic solvents.
o suspension – mixture in which one phase is
broken down to small particles of insoluble
solid, droplets or bubbles spread throughout
the main phase (liquid).
o example: fine particles of a solid in a liquid
after a precipitation reaction.
o precipitation reaction – reaction in which an
insoluble salt is made from solutions of two
soluble salts.
Solutions
o Earth’s surface made of a mixture of solids and
solvents.
o examples:
- 2/3 of surface covered with salts dissolved in
sea water.
- dissolved gases oxygen and carbon dioxide
important for sea life exists in oceans.
o Common solvents are:
1. water.
2. organic solvents: ethanol, propanone,
trichloroethane.
o Why are organic solvents important
they dissolve substances that do not dissolve
in water.
o soluble – a solute that dissolves in a particular
solvent.
o insoluble – a substance that does not dissolves
in a particular solvent.
o miscible – two liquids form a completely
uniform mixture when combined.
o alloys – mixture made by joining liquid metals
and solidifying them.
- designed to have properties for a certain use.
solder = tin + lid (low melting point).
Examples of miscible/immiscible substances
Miscible
Immiscible
o water + milk
water + oil
o water + ethanol
water + kerosene
o water + vinegar
water + benzene
o water + lemon juice
honey + oil
Solubility of solids in liquids
o dissolving Copper (II) Sulfate in fixed volume of
water, solution becomes more concentrated as
more solid Copper (II) Sulfate is added.
o concentrated solution: higher proportion of
solute, small proportion of solvent (solute >
solvent).
o dilute solution (less concentrated): small
proportion of solute, high proportion of solvent
(solvent > solute).
o saturated solution: more solid solute is added
and a point is reached when no more solute
dissolves in solvent at constant temperature.
- solid dissolves when temperature is increased.
o concentration: measure of how much solute is
dissolved in solvent to make a solution.
o concentration of solute in saturated solution is
the solubility of the solute at that temperature.
o solubility of solids depends on temperature
e.g crystallization.
Solubility of gases in liquids
o gases less soluble in water as temperature
rises.
o solubility from air in water small, but dissolved
oxygen is enough to support aquatic life.
o solubility of gases increase with pressure.
o example: sparkling drinks contain carbon
dioxide dissolved under pressure.
o fizzing happens due to pressure release when
opening lid, container left open drink goes flat.