Chemistry AS Scheme of Work 2011-2012
Unit F321: Atoms, Bonds and Groups
Lesson
Num
1
Lesson Title
Learning Objectives
Atomic
Structure
1) Describe the structure and properties
of the atom
Module
1.1
2) Explain the distribution of mass and
charge in an atom
3) Explain the contribution of protons and
neutrons to the atomic and mass
number.
Prac
2
Ions and
isotopes
1) Be able to deduce the atomic and mass
number.
2) Be able to deduce the ion from its
Starter
Main
Plenary
atomic structure
3) Understand the term isotope and be
able to deduce number of protons and
neutrons.
Prac
3
Masses
1) Understand the term relative isotopic
mass, molecular mass, and relative
formula mass
2) Understand the relevance of 12C as a
standard measurement of mass
3) Be able to calculate the relative atomic
mass given relative abundances of
isotopes.
Prac
4+5
The mole
1) Understand the terms “amount of
substance” and “mole” and “Avogadro’s
constant NA”
2) Define and use the term ‘molar mass’
with units g mol-1
3) Know the terms empirical formula,
molecular formula and be able to
calculate them using composition by
mass and percentage compositions.
Prac
Formation of Potassium Iodide demo
Calculate formula from initial masses.
Atomic mass of metal based on collection
of hydrogen from reaction with acid.
6+7
Reacting
masses
1) Know the value of 1 mole for mass, gas
volume and liquid volumes and
concentrations
2) Be able to establish balanced
compounds from mol calculations
3) Evaluate the terms concentrated and
dilute for solutions
Prac
8+9
Produce different concentration solutions
using volumetric flask and distilled
water
Acids and
Bases
1) Explain that acids release H+ ions
in aqueous soln, and know the
formula for common acids.
2) Know that metal oxides, metal
hydroxides and ammonia are
common bases
3) State that an alkali is a soluble base,
releases OH- ions in soln, and know
the formula of common alkalis
Prac
10
pH probe to demonstrate a precise
measure of pH.
Record as alkali is added to acid.
Salts
1) Describe the reactions of an acid
with carbonates, bases and alkalis
to form a salt
2) Explain that salts are formed, when
the H+ of an acid is replaced by the
metal or NH4+
Prac
3) Explain that a base readily accepts
H+ ions (example include OH- and
NH3)
Produce and ammionium salt from
ammonia and acid reaction
11-12
Salts (II)
1) Explain the terms anhydrous,
hydrated and water of
crystallization
2) Be able to calculate the formula of a
hydrated salt from percentage
composition, mass composition or
experimental data.
3) Carry out acid/base and structured
titrations.
Anhydrous copper sulfate – record mass.
Add water slowly – record mass of
hydrated copper sulfate.
Calculate the formula: CuSO4.5H20
Prac
Titration – establish RFM of washing soda
by titration.
13-14
Oxidation
number
1) Understand the term oxidation
number
2) Be able to explain oxidation and
reduction in terms of electron
transfer and change in oxidation
number
3) Be able to apply the roman numeral
to indicate oxidation number for a
variety of ions and compounds.
TBA
Prac
15
Redox
Reactions
1) Explain why metals generally lose
electrons and non-metals gain
electrons – and what does this
mean for the oxidation number.
2) Describe redox reactions with HCl
and H2SO4
3) Make predictions about oxidation
numbers based on equations.
TBC
Prac
16
Module
1.2
Ionisation
energies
1) Describe the term first ionisation
energy, and second ionisation
energy.
2) Explain ionisation energies in terms
of nuclear charge, electron
shielding and size of atom.
3) Predict an element from first and
successive ionisation energies.
Prac
18 - 19
Electron
configuration
1) Name and order the 4 electron
shells (and sub shells), and know
how the orbitals are filled up to
element 36.
2) Be able to describe and draw s and
p orbitals.
3) Describe the relative energies for
the shells and relate this to s- pand d- block elements.
Prac
20
Ionic bonding
1) Describe ionic bonding as
electrostatic attraction between
oppositely charged ions.
2) Construct ‘dot and cross’ diagrams
for a variety of ionic compounds.
3) Make predictions about ionic
charge of ions based on periodic
table (including polyatomic ions e.g.
SO42-).
Prac
21
Covalent
bonding
1) Describe covalent bonding, with
examples of single and multiple
bonds.
2) Explain how dative covalent
bonding (coordinate) differs from
regular covalent bonding
3) Be able to draw dot and cross
diagrams of Cl2, HCl, H2O, NH3, CH4,
BF3 and NH4+, SF6
Prac
22
Shapes of
molecules and
ions
1) Explain the shape of simple
molecules based on repulsion of
electron pairs around a central
atom.
2) Understand that lone pairs repel
more than bonded pairs
3) Explain the shapes and bond angles
in trigonal planar, tetrahedral,
octahedral, pyramidal, non-linear
and linear.
Balloons modelling
Prac
23
Electronegativi
ty and bond
polarity
1) Describe the term electronegativity
as the ability of an atom to attract
bonding electrons in a covalent
bond
2) Explain how a permanent dipole
may form when covalently bonded
atoms have different
electronegativity
Prac
24
Intermolecular
forces
1) Describe intermolecular forces
based on permanent dipoles,
induced dipoles (van der Waals
forces), and noble gases.
2) Describe hydrogen bonding,
including the role of the lone pair,
in water, ammonia and –OH and –
NH groups.
3) Explain the link between density
and the melting and boiling points
of water – and hydrogen bonding.
Surface tension experiment(?)
Prac
25-26
Bonding and
physical
properties
1) Describe metallic bonding as the
attraction of positive ions to
delocalized electrons.
2) Describe the following structures:
giant ionic lattice, giant covalent
lattices, giant metallic lattices and
simple molecular lattices (eg ice)
3) Interpret physical properties from
different bonding types above:
Melting Points/Boiling Points
Electrical conductivity
Solubility
Prac
27
Module
1.3
Periodic Table
Groups and
Periods
1) Describe the order of the elements
in terms of increasing atomic
number – with periods and groups.
2) Describe periodicity in terms of
repeating patterns across different
periods.
3) Explain why elements in the same
group have similar chemical and
physical properties – with
discussion of electron configuration
Prac
28
Periodic Table
Physical
1) Explain the changes in ionisation
energy of elements across each
properties
period, and down each group (refer
to nuclear charge, atomic radius
and electron shielding).
2) Describe and explain the variations
in melting and boiling points across
periods 2 and 3 in terms of
structure and bonding.
3) Be able to interpret data on
electron configuration, atomic radii,
first ionisation energies, m.p. and
b.p. to describe periodicity
Prac
29 - 30
Group 2 Redox
reactions
1) Describe the redox reactions of the
Group 2 elements – with oxygen
and water – and the use of Ca(OH)2
in agriculture and Mg(OH)2 in
indigestion tablets.
2) Explain the trend in reactivity of
Group 2 elements down the group
due to increasing ease of cation
formation (atomic size, shielding,
nuclear attraction)
3) Explain the reactions of Group 2
oxides and carbonates – including
trends in thermal decomposition of
carbonates and resulting solutions
from Gp 2 oxides with water.
4) Interpret and make
predictionsfrom chemical and
physical properties of Gp 2
compounds.
Prac
Decomposition of Group 2 carbonates –
observe speed of reaction (colour change?)
Reaction of Gp2 oxides with water –
analysis of pH.
31
Group 7
Properties
Titration of indigestion tablets – pH probe
precision results and graph.
1) Explain the trend in boiling points
of Cl2, BR2, and I2 in terms of Van
der Waals’ forces.
2) Explain reactivity of the halides
with reference to negative ions,
atomic size, shielding and nuclear
attraction.
3) Interpret and make predictions
from the chemical and physical
properties of Gp 7 elements and
their compounds.
Prac
32
Group 7 Redox
reactions
1) Describe the redox reactions of Gp
7 elements, including ionic
reactions, in the presence of
organic solvent to illustrate relative
reactivity of Gp 7.
2) Describe the term
disproportionation as a reaction in
which an element is simultaneously
oxidized and reduced.
3) Contrast the benefits and
associated risks of using chlorine in
water treatments.
Prac
Halogen displacement reactions
32
Group 7 –
Characteristic
reactions
1) Describe the precipitation
reactions, including ionic equations,
of the aqueous Cl-, Br- and I- with
silver ions followed by aqueous
ammonia.
2) Describe the use of precipitation
reactions as a test for the different
halide ions.
Prac
Halogen displacement reactions
Unit F322 – Chains, Energy and Resources
Lesson
Num
Lesson Title
Learning Objectives
1
Representing
organic
compounds
1) Be able to recognise the following
representations “empirical formula”,
“molecular formula”, general formula,
Module
Starter
Main
Plenary
2.1
structural formula, displayed formula,
“skeletal formula” correctly
2) Be able to use the above terms correctly
3) Be able to convert one representation
into another
Prac
2
Homologous
series
1) Be able to state the name of the first ten
members of the alkane series
2) Be able to use the term homologous
series and functional group, with examples
3) Use a general formula of a homologous
series to predict the formula of any
member of the series.
Prac
3
IUPAC rules
1) Be able to recognise and and name
common organic functional groups
2) Be able to use the systematic naming
method to name simple compounds
featuring common functional groups
3) Be able to use the systematic naming
method to name more complex
compounds featuring common functional
groups
Prac
4
Isomerism
1) Be able to describe and explain
structural isomers, stereoisomers, E/Z
isomerism and cis-trans isomerisation.
2) Be able to explain the difference
between cis/trans and E/Z
stereoisomerism.
3) Be able to determine possible structural
formulae and stereoisomers of an organic
compounds.
Prac
5+6
Reaction
Mechanisms
1) Describe the different types of
covalent bond fission (homolytic
and heterolytic)
2) Describe a curly arrow as the
movement of an electron pair,
showing breaking or formation of a
covalent bond.
3) Be able to draw reaction
mechanisms to show the movement
of an electron pair with curly
arrows.
Prac
7+8
Percentage
yields and
atom economy
1) Be able to perform percentage yield
calculations.
2) Be able to explain atom economy,
and perform associated calculations
3) Explain that addition reactions
have an atom economy of 100%,
but substitution reactions are less
efficient.
4) Describe the benefits of developing
chemical processes with a high
atom economy.
5) Explain that a reaction may have a
high percentage yield but a low
atom economy.
Prac
9
Hydrocarbons
from crude oil
1) Explain the use of crude oil as a
source of hydrocarbons separated
by fractional distillation and used
as fuels of processing into
petrochemicals.
2) State that alkanes and cycloalkanes
are saturated hydrocarbons, and
explain the tetrahedral shape
around each carbon atoms in
alkanes.
3) Explain the variations in boiling
points of alkanes with different
chain lengths and branching in
terms of van der Waals forces.
Prac
Fractional distillation and investigation of
properties of fractions.
10
Hydrocarbons
as fuels
1) describe combustion of alkanes,
leading to use of fuels in industry,
homes and transport.
2) Explain, with the use of equations,
the difference between complete
and incomplete combustion and the
dangers arising from the problems
of CO.
3) Describe the use of catalytic
cracking to obtain more useful
alkanes and alkenes.
Observe differences between the different
combustion types when using a Bunsen
burner / burning splints.
Prac
11
Hydrocarbons
as fuels (II)
1) Explain that the petroleum industry
processes straight chain
hydrocarbons into branched
alkanes and cycloalkanes to
promote efficient combustion.
2) Be able to contrast the value of
fossil fuels for providing energy and
raw materials with an over reliance
on non renewable fossil fuel
reserves and the importance of
developing renewable plant based
fuels.
3) Be able to contrast the value of
fossil fuels for providing energy and
raw materials with increased CO2
levels from combustion of fossil
fuels leading to climate change and
global warming.
Energy of combustion investigation,
compare straight chain and cycloalkane
combustion.
Prac
12
Substitution
reactions of
alkanes
1) Define the term radical as a species
with an unpaired electron – and
show how this is represented in
equations.
2) Be able to describe the substitution
of alkanes using ultraviolet
radiation, by Cl2 and Br2, to form
halogenoalkanes.
3) Describe how homolytic fission
leads to the mechanism of radical
substitution in alkanes (initiation,
propagation and termination
reactions).
4) Explain the limitations of radical
substitution in synthesis – arising
from further substitution and
formation of a mixture of products.
Prac
13
Alkenes
1) State that alkenes and cycloalkenes
are unsaturated hydrocarbons.
2) Describe the overlap of adjacent porbitals to form a  bond
3) Explain the trigonal planar shape
around each carbon in the C=C of
alkenes.
Prac
14
Addition
1) describe addition reactions of
reactions of
alkenes
alkenes, with conditions – a) with
hydrogen, b) halogens, c) hydrogen
halides, d) steam to form alcohols.
2) Define an electrophile as an
electron pair acceptor
3) Describe how heterolytic fission
leads to the mechanism of
electrophilic addition in alkenes
Prac
15
Iodine and bromine tests for unsaturation.
Polymers from
alkenes
1) Describe the addition
polymerization of alkenes
2) Deduce the repeat unit of an
addition polymer obtained from a
given monomer.
3) Identify the monomer that would
produce a given section of an
addition polymer.
Prac
16 + 17
Industrial
importance of
alkenes
1) Outline the use of alkenes in the
industrial production of organic
compounds – manufacture of
margarine by catalytic
hydrogenation of unsaturated
vegetable oils and a nickel catalyst.
2) Describe the formation of a range of
polymers using unsaturated
monomer units based on the ethene
molecule (H2C=CHCl & F2C+CF2)
Prac
18
Mod
2.2
Extraction of limonene from orange peel.
Alcohol
1) Explain the water solubility and
low volatility of alcohols with
reference to hydrogen bonding.
2) Describe the industrial production
of ethanol by fermentation of
sugars, and the reaction of ethene
with steam in the presence of an
acid catalyst.
3) Outline the use of alcohols in
i)
ii)
iii)
alcoholic drinks
as a solvent
methanol as a petrol
additive
iv)
methanol as a feedstock in
the production of organic
chemicals.
Fermentation of glucose
Prac
19
Reaction of
alcohols
1) Be able to classify alcohols into
primary, secondary and tertiary
alcohols.
2) Be able to describe the combustion
of alcohols.
3) Describe the oxidation of alcohols
using Cr2O72- / H+ including:
i)
ii)
oxidation of primary
alcohols to from aldehydes
and carboxylic acids.
Oxidation of secondary
alcohols to form ketones
iii)
Prac
20
The resistance to oxidation
of tertiary alcohols
Oxidation of ethanol to aldehyde and carboxylic
acid.
Reaction of
alcohols (II)
1) Describe the esterification of
alcohols with carboxylic acids in the
presence of an acid catalyst
2) Describe the elimination of H2O
from alcohols in the presence of an
acid catalyst and heat to form
alkenes.
Prac
Preparation of esters on a test-tube scale.
Elimination of water from cyclohexanol.
21-22
Halogenoalkan
es
1) Define the term nucleophile as an
electron pair donor
2) describe the hydrolysis of
halogenoalkanes as a substitution
reaction.
3) Describe the mechanism of
nucleophilic substitution in the
hydrolysis of primary
halogenoalkanes with hot aqueous
alkali.
4) Explain the rates of hydrolysis of
primary halogenoalkanes in terms
of the relative bond enthalpies of
carbon-halogen bonds.
Compare rates of hydrolysis of
halogenoalkanes
Prac
22
Uses of
halogenoalkan
es
1) describe the uses of chloroethene
and tetrafluoroethene to produce
the plastics PVC and PTFE
2) Explain the reasons for using CFCs
and the environmental damage
caused by them
3) Describe the role of green
chemistry in minimising damage to
the environment by promoting
biodegradable alternatives to CFCs
(HCFCs, CO2 as a blowing agent for
expanded polymers).
Prac
23
Infrared
spectroscopy
1) State that absorption of infrared
radiation causes covalent bonds to
vibrate.
2) Identify alcohols, aldehydes,
ketones and carboxylic acids from
their infrared spectrum.
3) State that modern breathalysers
measure ethanol in the breath
using infrared spectroscopy.
Prac
24 - 25
Mass
spectrometry
1) Describe mass spectrometry a
method to determine relative
isotopic masses, and identifying
elements (for example on the Mars
probe / measuring pollution – eg
lead)
2) Interpret mass spectra of elements
in terms of isotopic abundances.
3) Interpret mass spectra of organic
compounds:
i)
ii)
iii)
to identify molecular mass
identify major fragment ions
predict structures from
peaks
4) explain that a mass spectrum is a
fingerprint for the molecule that can be
identified by computers (and database
library)
Prac
26
Mod 2.3
http://riodb01.ibase.aist.go.jp/sdbs/cgibin/cre_index.cgi?lang=eng
Enthalpy
changes
1) Explain that some chemical
reactions are accompanied by
enthalpy changes – and can be
exothermic (H negative) or
endothermic (H positive).
2) Describe the importance of
oxidation as an exothermic process
in the combustion of fuels and the
oxidation of carbohydrates.
3) Describe that endothermic
processes require an input of heat
energy, eg thermal decomposition
of calcium carbonate.
4) Construct a simple enthalpy profile
diagram for a reaction to show the
difference in the enthalpy of the
reactants compared with products
Prac
27
Direct enthalpy changes of reaction for simple
reactions:
Zn + CuSO4 (exo); NaHCO3 + citric acid (endo);
NaOH + HCl (exo).
Enthalpy
changes (II)
1) Explain qualitatively, using
enthalpy profile diagrams, the term
activation energy
2) Define and use the terms: standard
conditions (298K and 100kPa),
enthalpy change of reaction,
enthalpy change of formation, and
enthalpy change of combustion.
3) Calculate enthalpy changes from
appropriate experimental results
directly, including use of the
relationship: energy change = mcT
Prac
Enthalpy change of combustion of alcohols
28
Bond
enthalpies
1) Explain exothermic and
endothermic reactions in terms of
enthalpy changes associated with
breaking and making bonds.
2) Define the use of term average
bond enthalpy (positive H – bond
breaking of one mole of bonds)
3) Calculate an enthalpy change of
reaction from average bond
enthalpies.
Prac
29+30
Direct enthalpy changes of reaction for simple
reactions:
Zn + CuSO4 (exo); NaHCO3 + citric acid (endo);
NaOH + HCl (exo).
Hess’ law and
enthalpy cycles
1) Use Hess’ law to construct enthalpy
cycles and carry out calculations to
determine enthalpy change of
reaction from enthalpy changes for
combustion.
2) Use Hess’ law to construct enthalpy
cycles and carry out calculations to
determine enthalpy change of
reaction from enthalpy changes for
formation.
3) Use Hess’ law to construct enthalpy
cycles and carry out calculations to
determine enthalpy change of
reaction from an unfamiliar
enthalpy cycle.
Prac
31
Indirect enthalpy change of reaction:
2KHCO3 → K2CO3 + H2O + CO2 indirectly using
HCl.
Collision
theory
1) describe qualitatively, in terms of
collision theory, the effect of
concentration changes on the rate
of a reaction;
2) explain why an increase in the
pressure of a gas, increasing its
concentration, may increase the
rate of a reaction involving gases;
Prac
32 - 33
Rate graphs for gas products, eg CaCO3 + HCl;
Mg + HCl
Catalysts
1) State that a catalyst speeds up a
reaction without being consumed
by the overall reaction.
2) Explain that catalysts:
i)
ii)
iii)
iv)
affect the conditions that are
needed, usually lowering
temperatures and reducing
energy demand (lower CO2
and fossil fuels)
allow scientists to use
different reactions with
better atom economy
are often enzymes – and
work at room temperature
and pressures
are of great economic
importance using 4
examples.
3) Explain, using enthalpy profile
diagrams, how a catalyst allows a
reaction to follow a different route
with a lower activation energy
(increasing reaction rate)
Prac
34
Boltzmann
distribution
1) Explain qualitatively the Boltzmann
distribution, and its relationship
with the activation energy.
2) Describe qualitatively, using the
Boltzmann distribution, the effect
of temperature changes on the
proportion of molecules exceeding
the activation energy and rate of
reaction.
3) Interpret how catalysts affect the
Boltzmann distribution.
Prac
35
Dynamic
Equilibrium
1) State le Chatelier’s principle
2) Explain that a dynamic equilibrium
exists when the rate of forward
reaction is equal to the rate of the
reverse reaction.
3) Apply le Chateliers’s principal to
deduce the effect of changes in
temperature, concentration and
pressure on a homogeneous system
in equilibrium
4) Explain (using data) the importance
in the chemical industry of a
compromise between chemical
equilibrium and reaction rates.
Prac
Changing equilibrium position with heat:
[Cu(H2O)6]2+ ⇌ CuCl42Changing equilibrium position with
concentration: Fe3+ and SCN–