Section 1 a) Specification

1.1 understand the arrangement, movement and energy of the particles in each

of the three states of matter: solid, liquid and gas

As shown in the diagram:

Solids have a fixed shape and volume. Their particles are arranged regularly and are tightly packed. They do not flow, and the particles vibrate about a fixed point. They have the strongest forces of attraction and are usually the most dense.

Liquids have a fixed volume, but their shape changes to fit the container. The particles have weaker forces of attraction than solids, but stronger than gases. The particles can flow over each other, and they are irregularly arranged.

Gases have no fixed shape or volume, they can be compressed, and they change shape to fit the container. They move randomly in different directions and collide with other particles. There are wide gaps between the particles and almost no forces of attraction between them.

1.2 understand how the interconversions of solids, liquids and gases are

achieved and recall the names used for these interconversions

Melting- Solid --> Liquid
Achieved by heating, increased thermal energy means the particles gain kinetic energy.

Boiling- Liquid --> Gas
Achieved by heating, increased thermal energy means the particles gain kinetic energy.

Freezing- Liquid --> Solid
Achieved by cooling, decreased thermal energy means the particles lose kinetic energy.

Condensing- Gas --> Liquid
Achieved by cooling, decreased thermal energy means the particles lose kinetic energy.

Evaporation- Liquid--> Gas
Some particles on the surface move faster than average and have enough energy to break the bonds and become a gas.

Sublimation- Solid--> Gas
Achieved by a sudden increase n temperature or high pressure conditions. e.g dry ice. 

1.3 explain the changes in arrangement, movement and energy of particles

during these interconversions.

Melting- The solid particles vibrate at an increasing speed with increased temperature until the forces of attraction can no longer hold them together, and it changes state.

Boiling- The liquid is heated to the point that all the forces of the attraction within it are broken and bubbles of gas are formed throughout the liquid.

Evaporation- Faster particles on the surface of the liquid have enough energy to break free from their bonds and evaporate to form a gas.

Condensing- The gas is cooled and the particles move slowly enough that they can form attractions between each other and form a liquid.

Freezing- If a liquid is cooled, the particles will begin to lose kinetic energy and their movement will slow until the forces of attraction will hold them in place as a solid.

Subliming- The particles gain kinetic energy very quickly, allowing all forces of attraction to break quickly and form a gas.

Section 1 a) Summary

Particles of a substance can be present in different physical states, the states of matter. They are solid, liquid and gas, and increasing or decreasing the kinetic energy of the particles changes which state the substance will appear in. 

The properties of the different states can be clearly seen in the table below:

When a solid is heated to its melting point, the particles gain enough kinetic energy to break the forces of attraction between the particles, so the particles are no longer in a regular, densely packed and fixed pattern. The space between the particles has increased enough that the physical state has changed.

This is the same for boiling, the state of matter changes from liquid to gas due to increased movement of particles. Freezing and condensing work similarly, but the opposite way.

Sublimation is the change from a solid to a gas, skipping the liquid phase. This can happen due to a sudden increase in temperature, or at high pressure that means the particles spread quickly, forming a gas instead of melting.

Changes in state are also affected by pressure, the melting and boiling point can be increased by increasing the pressure. this is because increased pressure forces particles closer together and means it is more difficult for them to spread out at the same temperature.

Section 1 a) Key Words

Boiling: The change of state from liquid to gas with an increase in thermal energy, to the substance's boiling point. The liquid becomes a gas because the particles are able to spread apart enough to change from a liquid phase to a gaseous phase.

Boiling Point: The temperature at which boiling occurs in a certain liquid, e.g. water's boiling point is 100ºC at the pressure of 1 atm., as boiling point increases as pressure increases due to the pressure exerted on the particles changing their ability to spread and boil.

Condensing: The change of state from gas to liquid. This happens when a gas cools, often upon colliding with a solid surface, so the particles collect and become more dense until they change physical state.

Freezing: A change of state from liquid to solid. When the particles lose kinetic energy they move less and become more densely packed. Attractions between the particles cause bonds or intermolecular forces form between the particles to keep them in fixed positions.

Gas: A state of matter in which the particles move randomly in different directions, colliding with other particles and surfaces and changing direction. The particles are spread over a wide space; it is the least dense state of matter. It has no fixed shape or volume, meaning it can be compressed and squeezed, and it flows.

Liquid: A state of matter involving randomly organised particles that are fairly closely packed, but less densely than in a solid. It has a fixed volume but no fixed shape; it will take the shape of its container. The particles are able to move and flow over each other.

Melting: The change of state from solid to liquid. Due to an increase in thermal energy, the kinetic energy of the particles increases, making the bonds or intermolecular forces break and allowing the particles to be freed from their fixed positions and flow over one another.

Solid: A state of matter in which the particles are closely and regularly packed and vibrate about a fixed point. It does not flow and has a fixed shape and volume, and it is the most dense state of matter.

Sublimation: The change of state from solid to gas, skipping the liquid phase. (Deposition is the same but changing state the opposite direction, from gas to solid). Examples include carbon dioxide, which sublimes into dry ice, or when the sun heats a field of snow. Sublimation occurs when the total pressure of the atmosphere is less than the vapor pressure of the compound, and the temperature has not yet reached the substance's melting point. The particles spread out quickly to change state from solid to gas quickly.

Section 1: Principles of Chemistry Specification

a) States of matter

1.1 understand the arrangement, movement and energy of the particles in each
of the three states of matter: solid, liquid and gas

1.2 understand how the interconversions of solids, liquids and gases are
achieved and recall the names used for these interconversions

1.3 explain the changes in arrangement, movement and energy of particles
during these interconversions.

b) Atoms

1.4 describe and explain experiments to investigate the small size of particles
and their movement including:
i dilution of coloured solutions
ii diffusion experiments

1.5 understand the terms atom and molecule

1.6 understand the differences between elements, compounds and mixtures

1.7 describe experimental techniques for the separation of mixtures, including
simple distillation, fractional distillation, filtration, crystallisation and paper
chromatography

1.8 explain how information from chromatograms can be used to identify the
composition of a mixture.

c) Atomic structure

1.9 understand that atoms consist of a central nucleus, composed of protons
and neutrons, surrounded by electrons, orbiting in shells

1.10 recall the relative mass and relative charge of a proton, neutron and electron
1.11 understand the terms atomic number, mass number, isotopes and relative
atomic mass (Ar)

1.12 calculate the relative atomic mass of an element from the relative
abundances of its isotopes

1.13 understand that the Periodic Table is an arrangement of elements in order of
atomic number

1.14 deduce the electronic configurations of the first 20 elements from their
positions in the Periodic Table

1.15 deduce the number of outer electrons in a main group element from its
position in the Periodic Table.

d) Relative formula masses and molar volumes of gases

1.16 calculate relative formula masses (Mr) from relative atomic masses (Ar)

1.17 understand the use of the term mole to represent the amount of substance

1.18 understand the term mole as the Avogadro number of particles
(atoms, molecules, formulae, ions or electrons) in a substance

1.19 carry out mole calculations using relative atomic mass (Ar) and relative
formula mass (Mr)

1.20 understand the term molar volume of a gas and use its values
(24 dm3 and 24,000 cm3) at room temperature and pressure (rtp) in
calculations.

e) Chemical formulae and chemical equations

1.21 write word equations and balanced chemical equations to represent the
reactions studied in this specification

1.22 use the state symbols (s), (l), (g) and (aq) in chemical equations to
represent solids, liquids, gases and aqueous solutions respectively

1.23 understand how the formulae of simple compounds can be obtained
experimentally, including metal oxides, water and salts containing water of
crystallisation

1.24 calculate empirical and molecular formulae from experimental data

1.25 calculate reacting masses using experimental data and chemical equations

1.26 calculate percentage yield

1.27 carry out mole calculations using volumes and molar concentrations.

f) Ionic compounds

1.28 describe the formation of ions by the gain or loss of electrons

1.29 understand oxidation as the loss of electrons and reduction as the gain of
electrons

1.30 recall the charges of common ions in this specification

1.31 deduce the charge of an ion from the electronic configuration of the atom
from which the ion is formed

1.32 explain, using dot and cross diagrams, the formation of ionic compounds by
electron transfer, limited to combinations of elements from Groups 1, 2, 3
and 5, 6, 7

1.33 understand ionic bonding as a strong electrostatic attraction between
oppositely charged ions

1.34 understand that ionic compounds have high melting and boiling points
because of strong electrostatic forces between oppositely charged ions

1.35 understand the relationship between ionic charge and the melting
point and boiling point of an ionic compound

1.36 describe an ionic crystal as a giant three-dimensional lattice
structure held together by the attraction between oppositely
charged ions

1.37 draw a diagram to represent the positions of the ions in a crystal of
sodium chloride.

g) Covalent substances

1.38 describe the formation of a covalent bond by the sharing of a pair of
electrons between two atoms

1.39 understand covalent bonding as a strong attraction between the bonding
pair of electrons and the nuclei of the atoms involved in the bond

1.40 explain, using dot and cross diagrams, the formation of covalent compounds
by electron sharing for the following substances:
i hydrogen
ii chlorine
iii hydrogen chloride
iv water
v methane
vi ammonia
vii oxygen
viii nitrogen
ix carbon dioxide
x ethane
xi ethene

1.41 understand that substances with simple molecular structures are gases or
liquids, or solids with low melting points

1.42 explain why substances with simple molecular structures have low melting
and boiling points in terms of the relatively weak forces between the
molecules

1.43 explain the high melting and boiling points of substances with giant covalent
structures in terms of the breaking of many strong covalent bonds

1.44 draw diagrams representing the positions of the atoms in diamond
and graphite

1.45 explain how the uses of diamond and graphite depend on their
structures, limited to graphite as a lubricant and diamond in cutting.

h) Metallic crystals

1.46 understand that a metal can be described as a giant structure of positive
ions surrounded by a sea of delocalised electrons

1.47 explain the electrical conductivity and malleability of a metal in terms of its
structure and bonding.

i) Electrolysis

1.48 understand that an electric current is a flow of electrons or ions

1.49 understand why covalent compounds do not conduct electricity

1.50 understand why ionic compounds conduct electricity only when molten or in
solution

1.51 describe experiments to distinguish between electrolytes and nonelectrolytes

1.52 understand that electrolysis involves the formation of new substances when
ionic compounds conduct electricity

1.53 describe experiments to investigate electrolysis, using inert electrodes, of
molten salts such as lead(II) bromide and predict the products

1.54 describe experiments to investigate electrolysis, using inert
electrodes, of aqueous solutions such as sodium chloride, copper(II)
sulfate and dilute sulfuric acid and predict the products

1.55 write ionic half-equations representing the reactions at the electrodes during
electrolysis

1.56 recall that one faraday represents one mole of electrons

1.57 calculate the amounts of the products of the electrolysis of molten
salts and aqueous solutions.