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Scheme of work – Cambridge International AS Level Physical Science (8780)
CHEMISTRY – SECTION I
Unit 1: Theoretical Chemistry
Recommended prior knowledge
A reasonable standard at IGCSE Science or Chemistry is assumed for this unit.
Context
This unit underpins all succeeding units, so needs to be studied first.
Outline
The unit covers most of the calculations needed for the Chemistry section of this syllabus, involving the mole and empirical formulae. It also includes atomic
structure, bonding, intermolecular forces, and the behaviour of gases.
Syllabus ref
Learning objectives
Suggested teaching activities
Learning resources
Chemistry for Advanced Level (ISBN
9780719586026) 1.1-1.7
Advanced Chemistry (ISBN
9780521423328) 37
AS Level and A Level Chemistry (ISBN
9780521544719) 2.1, 2.6
Candidates should be able to:
C1(a)
define the terms relative atomic,
isotopic, molecular and formula
masses, based on the 12C scale
Definitions of the mole (= an Avogadro number of particles), and molar
mass (= the mass of one mole of substance).
C1(b)
define the term mole in terms of
the Avogadro constant
Definition of Avogadro number (L = number of atoms in 12.0 g of 12C)
Definitions of relative atomic mass (Ar), isotopic mass, molecular
and formula masses (Mr) in terms of 1/12 the mass of 1 mole of 12C
atoms.
Some revision of IGCSE ideas of the mole.
Simple determinations of L in the lab could include the oil-drop expt. (If
we make assumptions about the density and the Mr of the oil, and the
shape of its molecule (cubic), we can calculate the volume of 1 mole,
and the volume of 1 molecule. Hence we can calculate L.)
A more accurate estimate of L comes from electro-deposition of copper
during electrolysis of CuSO4(aq), making assumptions about valency
v2 2Y07
Cambridge International AS Level Physical Science (8780)
site 3 (Avogadro)
site 6 (atomic, molecular mass)
site 7 (what is a mole?)
1
Syllabus ref
Learning objectives
Suggested teaching activities
Learning resources
and Mr of Cu, and the charge on the electron.
C1(c)
analyse mass spectra in terms of
uniatomic species in terms of m/e
values (knowledge of the working
of the mass spectrometer is not
required)
Knowledge that a mass spectrum plots relative abundance against mass
number. Translating a table of relative abundances and mass numbers
into a plotted mass spectrum, and vice versa.
C1(d)
calculate the relative atomic mass
of an element given the relative
abundances of its isotopes, or its
mass spectrum
Measuring peak heights from a given mass spectrum, and calculating
the relative atomic mass from the weighted mean of the peaks [Ar =
Σaimi where ai = the fractional abundance of isotope of m/e (mass) mi]
Searching databases and periodic table tabulations will provide a host of
examples of elements having several isotopes. Calculating their relative
atomic masses is a useful exercise.
C1(e)
define the terms empirical and
molecular formulae
Definitions of empirical formula (= simplest whole number mole ratio of
atoms of each element in a molecule) and molecular formula (= actual
number of atoms of each element in a molecule).
C1(f)
calculate empirical and molecular
formulae, using combustion data
or composition by mass
Calculation of empirical formulae using mass data, percentage data or
combustion data. Calculation of molecular formula from empirical
formula and relative molecular mass.
Class sets of experimental data (e.g. on the heating of Mg in air), or
sheets showing a series of results, could be analysed using computer
spreadsheets or by a graphical method.
Various worksheets containing suitable data for simple compounds such
as Fe2O3 and other oxides, CaCO3 and other salts, C2H6 and other
hydrocarbons (mass, % and combustion data), CH3CO2H etc.
Experimental work could include the combustion of Mg in a weighed
crucible; passing H2(g) or CH4(g) over heated CuO; heating BaCl2.nH2O
crystals to find n.
C1(g)
write balanced equations
Revise the valencies of ions from IGCSE. Revise the four state
symbols. Provide a worksheet of essential formulae to learn (e.g. simple
covalent compounds, common salts, acids and alkalis). Students should
practice both balancing equations where the formulae are given, and
also constructing balanced chemical equations, given the word
equations.
v2 2Y07
Cambridge AS Level Physical Science (8780)
Chemistry for Advanced Level 2.5
Advanced Chemistry 29
AS Level and A Level Chemistry 2.2,
2.3, 2.4, 2.5
Chemistry for Advanced Level 1.8
Advanced Chemistry 37
AS Level and A Level Chemistry 2.7,
2.8
Teaching AS Practical Skills 4
Chemistry for Advanced Level 1.9
AS Level and A Level Chemistry 2.9
site 9 (equations, redox)
2
Syllabus ref
Learning objectives
Suggested teaching activities
Learning resources
C1(h)
perform calculations, including
use of the mole concept,
involving:
(i) reacting masses (from
formulae and equations)
Use of the equations
moles = mass (g)/Mr or moles = mass (g)/Ar
moles = pressure (Pa) × gas vol (m3) / (R × temperature (K)
moles = vol of solution (dm3) x concentration (mol dm–3)
Chemistry for Advanced Level 1.101.13
Advanced Chemistry 39, 40
AS Level and A Level Chemistry 2.10
Teaching AS Practical Skills 2, 3
(ii) volumes of gases using the
ideal gas equation, pV = nRT
State and use the general gas equation pV = nRT in calculations,
including the determination of Mr, where Mr = mRT/pV [units: m in g; T in
K; p in Pa and V in m3]
Experiments weighing known volumes of gases in gas syringes (the use
of a gently heated oven extends this to volatile liquids) can be used
here. Plugging results into a spreadsheet allows the effects of
experimental error to be seen clearly.
site 3 (calculations)
Chemistry for Advanced Level 4.13
Advanced Chemistry 34
AS Level and A Level Chemistry 4.4
Teaching AS Practical Skills 5, 6
(iii) volumes and concentrations
of solutions
These equations are used across the syllabus, but initial experiments
and/or calculations could include the following:
Acid-base and redox titrations etc. e.g.
• calculate n in Na2CO3.nH2O;
• [CH3CO2H] in vinegar;
• %CaCO3 in limestone (back titration with HCl & NaOH);
• %Fe2+ in iron pills (& investigation of Fe2+/MnO4- ratio)
C1(i)
deduce stoichiometric
relationships from calculations
such as those in C1(h)
C2(a)
describe the number and relative
energies of the s, p and d orbitals
for the principal quantum
numbers 1, 2 and 3 and also the
4s and 4p orbitals
Distinguish between shells (containing a maximum of 2, 8, 18 etc
electrons), subshells (s, p, d, containing a maximum of 2, 6, 10
electrons), and orbitals (containing a maximum of 2 electrons each).
The number of subshells increases by 1 for each successive shell. The
number of orbitals in each subshell is fixed as s=1; p=3; d=5 etc. Within
each shell the relative energies are d > p > s. The 4s is lower in energy
than the 3d from element number 19 (K) and above.
Chemistry for Advanced Level 2.8
Advanced Chemistry 12
AS Level and A Level Chemistry 1.11
C2(b)
describe the shapes of s and p
orbitals
Spherical and dumb-bell shaped. The three p orbitals are at right angles
to each other.
Chemistry for Advanced Level 2.9
Advanced Chemistry 11
AS Level and A Level Chemistry 1.10
C2(c)
state the electronic configuration
of atoms and ions for elements 1
to 36
The aufbau principle. There are various mnemonics for remembering
the order in which the orbitals are filled. Use the convention 1s2 2s2 2p4
etc (for simple atoms this could be extended to include 2px22py12pz1)
Chemistry for Advanced Level 2.13
Advanced Chemistry 12
AS Level and A Level Chemistry 1.12
v2 2Y07
Cambridge AS Level Physical Science (8780)
site 4 (atomic structure)
site 15
3
Syllabus ref
Learning objectives
Suggested teaching activities
Learning resources
C2(d)
explain and use the term
ionisation energy
Define I.E. in terms of the energy required to remove 1 electron from
each atom/ion in 1 mol of gaseous atoms/ions.
C2(e)
explain the factors influencing the
ionisation energies of elements
Factors include proton number; shielding by inner shells; distance of
outermost electron from the nucleus.
Chemistry for Advanced Level 2.14
Advanced Chemistry 13
AS Level and A Level Chemistry 1.7,
1.8
C2(f)
explain the trends in ionisation
energies across Period 3 and
down the groups of the Periodic
Table (see also Section C7)
I.E. increases across a period and decreases down a group. Slight
decreases: from Gp II to Gp III due to the p subshell experiencing more
shielding than the s subshell; and from Gp V to Gp VI due to repulsion
between electrons sharing same orbital.
C3(a)
describe ionic (electrovalent)
bonding, as in sodium chloride
and magnesium oxide, including
the use of 'dot-and-cross'
diagrams
Emphasise the electrostatic attraction between two oppositely charged
ions; the giant nature of ionic structures (e.g. the cubic lattice of NaCl
and MgO). Unless otherwise stated, outer shells only need to be drawn.
Usually only the electrons on the product ions need to be shown, but the
use of dots and crosses to show which electrons have been transferred
from metal to non-metal is recommended. Resultant charges on the ions
should be shown. Other examples could be used, including LiF, Li2O,
MgF2
Chemistry for Advanced Level 4.6-4.8
Advanced Chemistry 18
AS Level and A Level Chemistry 3.1
describe covalent bonding as
involving a shared pair of
electrons including co-ordinate
(dative covalent) bonding as, for
example, in the formation of the
ammonium ion and in the Al2Cl6
molecule
Bonding descriptions should include the use of 'dot-and-cross' diagrams,
for example in:
(i) covalent bonding, as in hydrogen; oxygen; chlorine; hydrogen
chloride; carbon dioxide; methane; ethene
(ii) co-ordinate bonding, as, in the formation of the ammonium ion and
in the Al2Cl6 molecule.
Chemistry for Advanced Level 3.2-3.8
Advanced Chemistry 14
AS Level and A Level Chemistry 3.3,
3.4
site 7 (N-ch1-05, 06, 07)
C3(b)
site 4 (bonding and structure)
Dot-and-cross structures for the molecules mentioned (outer shells
only). Emphasise that bonds are stable entities, so give out heat when
they form. This stability is due to attraction of the bonding electrons to
two nuclei rather than just one.
The use of two dots (or two crosses) in a dative bond will show which is
the donor atom. In “line = bond” diagrams dative bonds can be shown
as a bond with an arrow (e.g. Cl→Al).
C3(c)
v2 2Y07
explain the shapes of, and bond
angles in, molecules by using the
qualitative model of electron-pair
Describe the principles of VSEPR. The bond angle trend in CH4, NH3
and H2O (109½o, 107½o, 104½o) should be explained (lone pair
repulsion > bond pair repulsion) and learned. Students should be
Cambridge AS Level Physical Science (8780)
Chemistry for Advanced Level 3.9
Advanced Chemistry 17
AS Level and A Level Chemistry 3.8
4
Syllabus ref
Learning objectives
Suggested teaching activities
Learning resources
repulsion (including lone pairs),
using simple examples such as:
CO2 (linear); BF3 (trigonal planar);
CH4 (tetrahedral); NH3
(pyramidal); H2O (bent or V
shaped); SF6 (octahedral)
exposed to a range of different molecules and ions in order to develop
experience in the prediction of shapes. For example, AlCl3, CH2O, SiCl4,
PCl3, C2Cl4. Further examples could include BeF2, SF2, SO3 , HOCl,
and, for the advanced student, ClF3, ClF5 and XeF4,
The hands-on use of molecular models will help students understand the
3-D shapes of molecules.
site 2 (VSEPR)
site 7 (N-ch1-15)
site 14 (animated molecules)
C3(d)
describe covalent bonding in
terms of orbital overlap, giving σ
and π bonds
σ bonds are formed by the overlap of an s orbital with another s orbital
(e.g. in H2); overlap of an s orbital with a p orbital (e.g. CH4 – there is no
need to introduce the idea of hybridisation); or end-on overlap of two ptype orbitals (e.g. F2 or C2H6). π bonds are formed by the sideways
overlap of two p orbitals (e.g. C2H4). Explain that this results in restricted
rotation around the C=C bond.
Chemistry for Advanced Level 3.13
Advanced Chemistry 15
site 14 (double bond formation)
site 4 (bonding and structure)
C3(e)
explain the shape of, and bond
angles in, the ethane and ethene
molecules in terms of σ and π
bonds (see also Section C12 )
Assume all bond angles are 109½o in ethane and 120o in ethene. The
bonding in ethene should be considered as a planar framework of five σ
bonds, on which is superimposed a π bond. See also C12.
Chemistry for Advanced Level 3.13,
22.3
C3(f)
(i) explain and use the term
electronegativity
Electronegativity explained as the ability/strength of an atom to attract
the bonding pair of electrons in a covalent bond.
Factors to include are proton number, atomic radius and shielding.
(ii) explain the factors influencing
the electronegativity of
elements
Electronegativity increases across Period 3 due to increasing proton
number and decreasing atomic radius. Electronegativity decreases down
Group VII due to increased atomic radius and shielding.
(iii) explain the trend in
electronegativity across
Period 3 and down Group VII
of the periodic table (see also
Section C7)
C3(g)
v2 2Y07
explain the terms bond energy,
bond length and bond polarity
(attributed to differences in
electronegativity) and use them to
compare the reactivities of
covalent bonds (see also
Bond energy in terms of the breaking of 1 mole of bonds in the gas
phase (i.e. the endothermic change); bond length as the distance
between the centres of adjacent atoms; bond polarity in terms of
differing electronegativities of the two bonded atoms. Reactivity is
associated with weak or long bonds (e.g. F-F, or O-O in H2O2 compared
Cambridge AS Level Physical Science (8780)
Chemistry for Advanced Level 3.10,
5.5
Advanced Chemistry 19, 42
AS Level and A Level Chemistry 3.7
5
Syllabus ref
Learning objectives
Suggested teaching activities
Sections C4 and C14)
to C-C in ethane; C-I compared to C-Cl). See also Sections C4 and
C14.
describe hydrogen bonding
(attributed to a large difference in
electronegativity), using ammonia
and water as simple examples of
molecules containing N-H and OH groups
Represent the hydrogen bond as a dotted line. Include the lone pair and
the partial charges on diagrams thus: OδHδ+― . Mention the lack of
hydrogen bonding in HCl, H2S etc.
C3(i)
outline the importance of
hydrogen bonding to the physical
properties of substances,
including ethanol and water
A simple understanding of the role of hydrogen bonding in generating
the relatively high melting point of water, the boiling points of water and
of ethanol and the solubility of ethanol in water. An understanding of the
ice structure of water is not required.
C3(j)
describe the intermolecular
forces, based on permanent and
induced dipoles, as in CHCl3(l);
Br2(l) and the liquid noble gases
Distinguish between polar bonds and polar molecules (e.g. the
molecules of CO2 and CCl4 are not polar, although they contain polar
bonds). Induced dipole – dipole forces (often referred to as van der
Waals' forces) are experienced by all molecules, and their strength
depends on the total number of electrons. The additional attraction of
permanent dipoles is usually weaker.
Chemistry for Advanced Level 3.11,
3.13
Advanced Chemistry 20
AS Level and A Level Chemistry 3.10
C3(h)
Learning resources
Chemistry for Advanced Level 3.15
Advanced Chemistry 21
AS Level and A Level Chemistry 3.12,
3.13
site 7 (N-ch1-09)
site 3 (intermolecular forces)
C3(k)
describe metallic bonding in
terms of the attraction between a
lattice of positive ions and
delocalised electrons
The strength of metallic bonding explained in terms of the charge on the
metal cation, the number of delocalised electrons and the size of the
cation. I.e. the more outer-shell electrons there are, the more electrons
are delocalised. Using Na, Mg and Al as examples, the cations in the
lattice are Na+, Mg2+ and Al3+. As all three are isoelectronic, but the
proton number increases across the period, the ionic radii decrease. So,
in aluminium, the Al3+ ion is the most highly charged, it has contributed
the most delocalised electrons and it is the smallest ion. Thus, the
strength of the metallic bond, and hence the melting point, increases so
that the m.pt. Al > Mg > Na.
Chemistry for Advanced Level 4.11
Advanced Chemistry 22
AS Level and A Level Chemistry 3.9
C3(l)
describe, in simple terms, the
lattice structure of a solid which
is:
(i) ionic, as in sodium chloride,
magnesium oxide
A lattice is described as a regular three-dimensional array of ions,
molecules or atoms
The physical properties of the lattice explained in terms of:
Electrostatic attractions between ions of opposite charge. The
alternating oppositely charged ions in 3 dimensions in ionic solids allows
Chemistry for Advanced Level 4.4, 4.5,
4.9, 4.11
Advanced Chemistry 18, 32
AS Level and A Level Chemistry 3.2,
3.11, 4.5
v2 2Y07
Cambridge AS Level Physical Science (8780)
6
Syllabus ref
Learning objectives
(ii) simple molecular, as in iodine
(iii) giant molecular, as in
graphite; diamond; silicon(IV)
oxide
(iv) metallic, as in copper
(the concept of unit cell is not
required)
Suggested teaching activities
Learning resources
a strong attraction between them.
The regular zigzag array of I2 molecules, with weak induced dipoledipole (van der Waal’s) forces between the molecules causes the m.pt.
to be low.
The need to break many strong covalent bonds results in the high m.pts.
of giant molecular solids. The continuous, 3-dimensional, tetrahedral,
strongly-bonded covalent structure of diamond results in the great
hardness of diamond. The layer nature of graphite (briefly mention the
“2-dimensional metal” nature of the electron delocalisation within each
layer being responsible for its electrical conductivity, but no detail is
required, since delocalisation in benzene is not covered in this syllabus).
The strength of metallic bonding due to the close packing of metal atoms
(“ions”) in their sea of delocalised electrons.
The use of molecular model kits is helpful, as are some interactive CDs.
Teaching AS Practical Skills 7
C3(m)
describe, interpret and/or predict
the effect of different types of
bonding (ionic bonding; covalent
bonding; hydrogen bonding; other
intermolecular interactions;
metallic bonding) on the physical
properties of substances
See also C3(l). Distinguish between giant structures (diamond, metals,
salts) that have high melting points, and simple molecular structures
(water, CO2 etc) that have low melting points. The increase of induced
dipole forces (and hence the increased m.pts. and b.pts.) with number of
electrons. A plot of the trends in b.pts. down Groups 4, 5, 6 and 7
against group number shows this, and also the effect of H-bonding in
NH3, H2O and HF. The enhanced b. pts. and solubility in water of
hydrogen –bonded compounds. The conductivity of metals and
molten/dissolved salts.
Chemistry for Advanced Level 4.4, 4.5,
4.10, 4.11, 4.12
AS Level and A Level Chemistry 3.14
C3(n)
suggest, from quoted physical
data, the types of structure and
bonding present in a substance
See C3(l and C3(m). Such information might be m.pts. or b.pts.,
solubilities, conductivities.
Chemistry for Advanced Level 4.4, 4.5,
4.9
AS Level and A Level Chemistry 3.15
C3(o)
show understanding of chemical
reactions in terms of energy
transfers associated with the
breaking and making of chemical
bonds
See C4(d). Bond-breaking being an endothermic process, and bondmaking an exothermic process. Reactions that break weak bonds or
make strong bonds are favoured.
Chemistry for Advanced Level 5.1, 5.5
Advanced Chemistry 42
C3(p)
describe and explain the
formation of salts from their
elements as redox processes in
OILRIG (oxidation is loss; reduction is gain (of electrons)) is a useful
mnemonic. Practice working out oxidation number in molecules, ions
and empirical formulae, using the equation Σ(O.N.) = overall charge.
Chemistry for Advanced Level 7.1-7.4
Advanced Chemistry 41
AS Level and A Level Chemistry 6a.1
v2 2Y07
Cambridge AS Level Physical Science (8780)
7
Syllabus ref
v2 2Y07
Learning objectives
Suggested teaching activities
terms of electron transfer and/or
of changes in oxidation number
(oxidation state)
Using ionic equations to show the redox reactions occurring in examples
such as Mg + O2; Cl2 + KI; Cl2 + FeCl2; Zn + CuSO4, all of which could
be carried out as practical work.
This concept should also be used more generally in the explanation of
redox reactions.
Cambridge AS Level Physical Science (8780)
Learning resources
8