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Introduction
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Chapter-1. The particulate nature of matter.
• State the distinguishing properties of solids,
liquids and gases
• Describe the structure of solids, liquids and gases
in terms of particle separation, arrangement and
types of motion
• Describe changes of state in terms of melting,
boiling, evaporation, freezing, condensation and
sublimation
• Describe qualitatively the pressure and
temperature of a gas in terms of the motion of
its particles
• Show an understanding of the random motion
of particles in a suspension (sometimes known
as Brownian motion) as evidence for the kinetic
particle (atoms, molecules or ions) model of
matter
• Describe and explain diffusion
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Chapter-2. Experimental techniques.
• Name appropriate apparatus for the
measurement of time, temperature, mass
and volume, including burettes, pipettes and
measuring cylinders
• Demonstrate knowledge and understanding of
paper chromatography
• Interpret simple chromatograms
• Identify substances and assess their purity from
melting point and boiling point information
• Understand the importance of purity in
substances in everyday life, e.g. foodstuffs and
drugs
• Describe and explain methods of purification
by the use of a suitable solvent, filtration,
crystallisation and distillation (including use of
a fractionating column). (Refer to the fractional
distillation of petroleum in section 14.2 and
products of fermentation in section 14.6.)
• Suggest suitable purification techniques, given
information about the substances involved
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Chapter-3. Atoms, elements and compounds.
• State the relative charges and approximate
relative masses of protons, neutrons and
electrons
• Define proton number (atomic number) as the
number of protons in the nucleus of an atom
• Define nucleon number (mass number) as the
total number of protons and neutrons in the
nucleus of an atom
• Use proton number and the simple structure of
atoms to explain the basis of the Periodic Table
(see section 9), with special reference to the
elements of proton number 1 to 20
• Define isotopes as atoms of the same element
which have the same proton number but a
different nucleon number
• State the two types of isotopes as being
radioactive and non-radioactive• State one medical and one industrial use of
radioactive isotopes
• Describe the build-up of electrons in ‘shells’
and understand the significance of the noble
gas electronic structures and of the outer
shell electrons. (The ideas of the distribution
of electrons in s and p orbitals and in d block
elements are not required.)
• Describe the differences between elements,
mixtures and compounds, and between metals
and non-metals
• Describe an alloy, such as brass, as a mixture of a
metal with other elements
• Describe the formation of ions by electron loss
or gain
• Describe the formation of ionic bonds between
elements from Groups I and VII
• Describe the formation of single covalent bonds
in H
2, Cl2, H2O, CH4, NH3 and HCl as the sharing
of pairs of electrons leading to the noble gas
configuration
• Describe the differences in volatility, solubility
and electrical conductivity between ionic and
covalent compounds
• Describe the giant covalent structures of graphite
and diamond
• Relate their structures to their uses, e.g. graphite
as a lubricant and a conductor, and diamond in
cutting tools
• Describe metallic bonding as a lattice of positive
ions in a ‘sea of electrons’ and use this to describe
the electrical conductivity and malleability of
metals
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Chapter-4. Stoichiometry.
• Use the symbols of the elements and write the
formulae of simple compounds
• Deduce the formula of a simple compound from
the relative numbers of atoms present
• Deduce the formula of a simple compound from
a model or a diagrammatic representation
• Construct word equations and simple balanced
chemical equations
• Define relative atomic mass, A
r, as the average
mass of naturally occurring atoms of an element
on a scale where the 12C atom has a mass of
exactly 12 units
• Define relative molecular mass, M
r, as the sum
of the relative atomic masses. (Relative formula
mass or M
r will be used for ionic compounds.)
(Calculations involving reacting masses in simple
proportions may be set. Calculations will not
involve the mole concept.)
• Define the mole and the Avogadro constant
• Use the molar gas volume, taken as 24 dm3 at
room temperature and pressure
• Calculate stoichiometric reacting masses,
volumes of gases and solutions, and
concentrations of solutions expressed in g / dm3
and mol / dm3. (Calculations involving the idea of
limiting reactants may be set. Questions on the
gas laws and the conversion of gaseous volumes
to different temperatures and pressures will not
be set.)
• Calculate empirical formulae and molecular
formulae
• Calculate percentage yield and percentage purity
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Chapter-5. Electricity and chemistry.
• Define electrolysis as the breakdown of an ionic
compound, molten or in aqueous solution, by the
passage of electricity
• Describe the electrode products and the
observations made during the electrolysis of:
– molten lead(II) bromide
– concentrated hydrochloric acid
– concentrated aqueous sodium chloride
– dilute sulfuric acid
between inert electrodes (platinum or carbon)
• State the general principle that metals or
hydrogen are formed at the negative electrode
(cathode), and that non-metals (other than
hydrogen) are formed at the positive electrode
(anode)
• Predict the products of the electrolysis of a
specified binary compound in the molten state
• Describe the electroplating of metals
• Outline the uses of electroplating• Describe the transfer of charge during electrolysis
to include:
– the movement of electrons in the metallic
conductor
– the removal or addition of electrons from the
external circuit at the electrodes
– the movement of ions in the electrolyte
• Describe the production of electrical energy from
simple cells, i.e. two electrodes in an electrolyte.
(This should be linked with the reactivity series in
section 10.2 and redox in section 7.4.)
• Describe, in outline, the manufacture of:
– aluminium from pure aluminium oxide in
molten cryolite (refer to section 10.3)
– chlorine, hydrogen and sodium hydroxide
from concentrated aqueous sodium chloride
(Starting materials and essential conditions
should be given but not technical details or
diagrams.)
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Chapter-6. Chemical energetics
• Describe the meaning of exothermic and
endothermic reactions
• Interpret energy level diagrams showing
exothermic and endothermic reactions
• Describe the release of heat energy by burning
fuels
• State the use of hydrogen as a fuel
• Describe radioactive isotopes, such as 235U, as a
source of energy
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Chapter-7. Chemical reactions.
• Identify physical and chemical changes, and
understand the differences between them
• Describe and explain the effect of concentration,
particle size, catalysts (including enzymes) and
temperature on the rate of reactions
• Describe the application of the above factors to
the danger of explosive combustion with fine
powders (e.g. flour mills) and gases (e.g. methane
in mines)
• Demonstrate knowledge and understanding of a
practical method for investigating the rate of a
reaction involving gas evolution
• Interpret data obtained from experiments
concerned with rate of reaction
Understand that some chemical reactions can
be reversed by changing the reaction conditions.
(Limited to the effects of heat and water on
hydrated and anhydrous copper(II) sulfate and
cobalt(II) chloride.) (Concept of equilibrium is
not required.)
Define oxidation and reduction in terms of oxygen
loss/gain. (Oxidation state limited to its use
to name ions, e.g. iron(II), iron(III), copper(II),
manganate(VII).)
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Chapter-8. Acids, Bases, Salts
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Chapter-9. The Periodic Table
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Chapter-10. Metals
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Chapter-11. Air and Water
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Chapter-12. Sulfur
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Chapter-13. Corbonates
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Chapter-14. Organic Chemistry
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