Advanced Higher Chemistry
Student Learning Outcomes
CHEMISTRY DEPARTMENT
ADVANCED HIGHER GRADE
CHEMISTRY
LEARNING OUTCOMES
UNIT
1
Electronic Structure and the Periodic Table
2
Principles of Chemical Reactions
3
Organic Chemistry
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Advanced Higher Chemistry
Student Learning Outcomes
UNIT
1
Electronic Structure and the Periodic Table
(a) Electronic Structure
(i)
Electromagnetic Spectrum and Associated Calculations
(ii)
Electronic Configuration and the Periodic Table
(iii)
Spectroscopy
(b) Chemical Bonding
(i)
Covalent Bonding
(ii)
Shapes of Molecules and Polyatomic Ions
(iii)
Ionic Lattices, Superconductors and Semiconductors
(c) Some Chemistry of the Periodic Table
2
(i)
The Second and Third Short Periods: Oxides, Chlorides and Hydrides
(ii)
Electronic Configuration and Oxidation States of Transition Metals
(iii)
Transition Metal Complexes
Principles of Chemical Reactions
(a) Stoichiometry
(b) Chemical Equilibrium
(i)
Reactions at Equilibrium
(ii)
Equilibria Between Different Phases
(iii)
Equilibria Involving Ions
(c) Thermochemistry
(i)
Hess’s Law
(ii)
Bond Enthalpies
(iii)
Hess’s Law Applied to Ionic Substances
(d) Reaction Feasibilty
(i)
Entropy
(ii)
Free Energy
(e) Electrochemistry
(f)
533571242
Kinetics
Page 2 of 31
Advanced Higher Chemistry
Student Learning Outcomes
UNIT
3 Organic Chemistry
(a) Permeating Aspects of Organic Chemistry
(i)
Reaction Types
(ii)
Reaction Mechanisms
(iii)
Physical Properties
(b) Systematic Organic Chemistry
(i)
Hydrocarbons and Halogenoalkanes
(ii)
Alcohols and Ethers
(iii)
Aldehydes, Ketones and Carboxylic Acids
(iv)
Amines
(v)
Aromatics
(c) Stereoisomerism
(i)
Geometric Isomerism
(ii)
Optical Isomerism
(d) Structural Analysis
(i)
Elemental Microanalysis and Mass Spectrometry
(ii)
Infra – red and Nuclear Magnetic Resonance Spectrometry and X – ray Crystallography
(e) Medicines
533571242
(i)
Historical Development
(ii)
How a Medicine Functions
Page 3 of 31
Advanced Higher Chemistry
Student Learning Outcomes
ADVANCED HIGHER CHEMISTRY STUDENT LEARNING OUTCOMES
UNIT 1 - ELECTRONIC STRUCTURE AND THE PERIODIC TABLE
(a)
ELECTRONIC STRUCTURE
( i)
Electromagnetic Spectrum and Associated Calculations
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
Electromagnetic radiation may be described in terms of waves.
2
1
2.
Electromagnetic radiation can be specified by its wavelength () and by its
frequency ().
2
1
3.
The electromagnetic spectrum is the range of frequencies or wavelengths of
electromagnetic radiation.
2
3
4.
The unit of measurement of wavelength is the metre or an appropriate
sub-multiple.
3
1
5.
The unit of measurement of frequency is the reciprocal of time in seconds (s -1)
and is called the Hertz (Hz).
3
1
6.
The velocity of electromagnetic radiation is constant and has a value of
approximately 3 x 108 ms-1.
2
1
7.
Velocity, frequency and wavelength are related in the expression: c = .
3
2
8.
Under certain circumstances electromagnetic radiation may be regarded as a
stream of particles, rather than as waves. These particles are known as photons.
12
4
9.
The energy (E) of radiation, and the energy associated with photons, is related to
frequency by Planck’s constant (h) in the expressions:
E = h for one photon
E = Lh for one mole of photons where L is Avogadro’s Constant.
12 – 13
4
65
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
( ii)
Electronic Configuration and the Periodic Table
LEARNING OUTCOME
1.
The emission spectrum of hydrogen provides evidence of energy levels.
19
5–6
65
2.
Quantum theory states that matter can only emit or absorb energy in small fixed
amounts (called quanta).
19
4
65
3.
The energy of a bound electron in an atom is quantised.
19
6
4.
An atom can be considered as emitting a photon of light energy when an electron
moves from a higher energy level to a lower energy level.
20
6
5.
Each line of the emission spectrum represents radiation of a specific wavelength or
frequency from which the difference in energy between the levels can be
calculated.
20
7
6.
Emission spectra of elements with more than one electron provide evidence of
sublevels within each principal energy level above the first.
21
10
7.
The principal energy levels correspond to the principal shells. The second and
subsequent principal shells contain sub-shells that correspond to the sub-levels.
22
10
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70
Advanced Higher Chemistry
( ii)
Student Learning Outcomes
Electronic Configuration and the Periodic Table
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
22
12
71
Heisenberg’s uncertainty principle states that it is impossible to define with
absolute precision, simultaneously, both the position and the momentum of an
electron.
23
11
10.
Electrons, like photons, display the properties of particles and waves.
23
11
11.
Treating bound electrons in atoms as waves leads to regions of high probability of
finding the electrons. These regions are called atomic orbitals.
23
11
72
12.
There are four types of orbitals, namely s, p, d and f, each with a characteristic
shape or set of shapes. Diagrams of the shapes of s and p orbitals can be drawn
and recognised. Diagrams of d orbitals can be recognised.
23 – 25
12 – 13
72
13.
An orbital holds a maximum of two electrons, as required by the Pauli exclusion
principle.
28
14
72
14.
The number of orbitals in each sub-shell is as follows:
23 – 25
13
71 - 72
LEARNING OUTCOME
8.
9.
Sub-shells can be labelled s, p, d and f. The types of sub-shells within each
principal shell are as follows:
Principal shell
Sub-shell(s) present
1
s
2
s and p
3
s, p and d
4
s, p, d, and f
Sub-shell
Number of orbitals
s
one s orbital
p
three p orbitals
d
five d orbitals
f
seven f orbitals
15.
In an isolated atom the orbitals within each sub-shell are degenerate.
24
12
16.
The Aufbau principle states that orbitals are filled in order of increasing energy.
28
16
17.
The relative energies corresponding to each orbital can be represented
diagrammatically for the first four shells of a multi-electron atom.
29
16 – 17
18.
Hund’s Rule states that when degenerate orbitals are available, electrons fill each
singly, keeping their spins parallel before spin pairing starts.
28
17
19.
Electronic configurations using spectroscopic notation and orbital box notation can
be written for elements of atomic numbers 1 to 36.
28 – 30
16
20.
The Periodic Table can be subdivided into four blocks (s, p, d and f) corresponding
to the outer electronic configurations of the elements within these blocks.
30 – 31
18
21.
The variation in first ionisation energy with increasing atomic number for the first
36 elements can be explained in terms of the relative stability of different electron
configurations, and so provides evidence for these electronic configurations.
31 – 32
18 – 19
22.
The relative values of first, second and subsequent ionisation energies can be
explained in terms of the stabilities of the electronic configurations from which the
electrons are being removed.
33 – 34
20
( iii)
Spectroscopy
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70
Advanced Higher Chemistry
Student Learning Outcomes
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
Atomic emission spectroscopy and atomic absorption spectroscopy involve
transitions between electronic energy levels in atoms. Generally, the energy
differences correspond to the visible region of the electromagnetic spectrum, i.e.,
to the approximate wavelength range of 400-700 nm. Some applications use the
ultra-violet region (wavelength range approximately 200-400 nm).
6–7
20 – 21
64 – 65
2.
In emission spectroscopy the sample is energised by heat or electricity causing
electrons to be promoted to higher energy levels. The wavelength of the radiation
emitted as electrons fall back to lower energy levels is measured.
6–7
21
64 – 65
3.
In absorption spectroscopy electromagnetic radiation is directed at the sample.
Radiation is absorbed as electrons are promoted to higher energy levels. The
wavelength of the absorbed radiation is measured.
6–7
21
4.
Each element provides a characteristic spectrum which can be used to identify an
element.
6–7
21
5.
The amount of species can be determined quantitatively if the intensity of emitted
or transmitted radiation is measured.
7–8
21 – 22
(b)
CHEMICAL BONDING
( i)
Covalent Bonding
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
39 – 40
23
97
LEARNING OUTCOME
1.
LEARNING OUTCOME
1.
Non- polar covalent bonding and ionic bonding can be considered as being at
opposite ends of a bonding continuum with polar covalent bonding lying between
these two extremes
65
2.
Different electron models can be used to explain the experimental evidence
associated with covalent bonding.
38
23 – 26
91
3.
Lewis electron dot diagrams represent bonding and non-bonding electron pairs in
molecules and in polyatomic ions.
41
26
87
4.
A dative covalent bond is one in which one atom of the bond provides both
electrons of the bonding pair.
42
27
88
5.
Species such as ozone, sulphur dioxide and the carbonate ion can be represented
by equivalent electron dot diagrams known as resonance structures
42 – 43
27 – 28
93 – 94
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Advanced Higher Chemistry
( ii)
Student Learning Outcomes
Shapes of Molecules and Polyatomic ions
LEARNING OUTCOME
1.
The shapes of molecules or polyatomic ions can be predicted from the number of
bonding electron pairs and the number of non-bonding electron pairs.
2.
The arrangement of electron pairs is linear, trigonal, tetrahedral, trigonal
bipyramidal and octahedral when the total number of electron pairs is 2, 3, 4, 5 and
6, respectively.
3.
Electron pair repulsions decrease in strength in the order:
non- bonding pair/ non-bonding pair > non-bonding pair/ bonding pair > bonding
pair/ bonding pair.
4.
These different strengths of electron pair repulsion account for slight deviations
from expected bond angles in molecules such as NH3 and H2O.
( iii)
Ionic Lattices, Superconductors and Semiconductors
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
44
29
92
45 – 46
29 –- 33
92
44 – 45
29
92
44
30 – 31
92
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
50
34 – 35
125 – 126
50
34 – 35
124 – 126
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
Io n ic L a t t ic e s
LEARNING OUTCOME
1.
2.
( iii)
The geometry of the crystalline structure adopted by an ionic compound depends
on the relative sizes of the ions. This affects the number of ions which can pack
round an ion of opposite charge.
Examples of crystal lattice structures are:
sodium chloride;
caesium chloride.
Ionic Lattices, Superconductors and Semiconductors
Superconductors
LEARNING OUTCOME
1.
Superconductors are a special class of materials that have zero electrical
resistance at temperatures near absolute zero
51
36
2.
Achieving temperatures near absolute zero is difficult and costly so application of
superconduction at these temperatures is impractical
51
36
3.
Recently superconductors have been discovered which have zero resistance up to
temperatures above the boiling point of liquid nitrogen-temperatures, which are
less costly to attain
51
36
4.
Superconductors may have future applications in power transmission and
electrically powered forms of transport.
52
37 – 38
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Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Ionic Lattices, Superconductors and Semiconductors
Semiconductors
SCHOLAR
LTS
NOTES
A covalent element such as silicon or germanium which has a higher conductivity
than that of a typical non-metal but a much lower conductivity than that of a metal
is described as a semiconductor
53
38 + 40
2.
Semiconductors are also referred to as metalloids and occur at the division
between metals and non-metals in the Periodic Table
53
40
3.
The electrical conductivity of semiconductors increases with increasing
temperature
53
38
4.
The electrical conductivity of semiconductors increases on exposure to light. This
is known as the photovoltaic effect.
53
40
5.
Elements such as silicon and germanium have similar structures to diamond but
the covalent bonds are weaker. Thermal agitation of the lattice can result in some
of the bonding electrons breaking free, leaving positive sites called ‘holes’.
53
40
6.
When a voltage is applied to these elements, electrons and holes can migrate
through the lattice.
53
40
7.
Doping pure crystals of silicon or germanium with certain other elements produces
n-type and p-type semiconductors.
54
41 – 42
8.
The type of semiconduction depends on the specific dopant used.
54
42
9
In n-type and p-type semiconductors the main current carriers are surplus
electrons and positive holes respectively.
54
42
10
Crystals of silicon or germanium can be prepared with bands of n-type or p-type
semiconductors. The p-n junction which occurs between a layer of n-type and a
layer of p-type semiconductor has specific electrical properties which form the
basis of the electronics industry.
54
42
Solar cells use the photovoltaic effect to convert sunlight into electricity.
54
43
LEARNING OUTCOME
1.
11
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CHEM in
CONTEXT
Advanced Higher Chemistry
Student Learning Outcomes
(c)
SOME CHEMISTRY OF THE PERIODIC TABLE
( i)
The Second and Third Short Periods: Oxides, Chlorides and Hydrides
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
Melting points, boiling points and electrical conductivities of the oxides, chlorides
and hydrides of the elements of the second and third periods can be explained in
terms of their structure and type of bonding.
60 – 73
44 – 53
177 – 182
2.
Metal oxides tend to be basic and non-metal oxides tend to be acidic but
amphoteric oxides exhibit both acidic and basic properties.
63 – 64
44 – 45
181 – 182
3.
Most ionic chlorides dissolve in water without reaction but some covalent chlorides
are hydrolysed, producing fumes of hydrogen chloride.
66
48
179
4.
Ionic hydrides possess the hydride ion, H─, which acts as a reducing agent.
69
51
5.
In reaction between water and ionic hydrides the products are hydrogen gas and
the hydroxide ion.
69
51
6.
Electrolysis of molten ionic hydrides produces hydrogen gas at the positive
electrode.
68
51
( ii)
Electronic Configuration and Oxidation States of Transition Metals
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
Electronic Configuration
LEARNING OUTCOME
1.
The d block transition metals are metals with an incomplete d sub-shell in at least
one of their ions.
77
54
251
2.
The filling of the d-orbitals follows the Aufbau principle, with the exception of
chromium and copper atoms. These exceptions are due to a special stability
associated with all the d-orbitals being half filled or completely filled.
78
54 – 55
250
3.
When transition metals form ions it is the s electrons which are lost first rather than
the d electrons.
79
55
251
( ii)
Electronic Configuration and Oxidation States of Transition Metals
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
Oxidation States
LEARNING OUTCOME
1.
An element is said to be in a particular oxidation state when it has a specific
oxidation number.
80
55
29
2.
The oxidation number is determined by following certain rules
81
55 – 56
29 – 30
3.
Transition metals exhibit variable oxidation states of differing stability.
82 – 86
57
254
4.
Compounds of the same transition metal but in different oxidation states may have
different colours.
90
57
255
5.
Oxidation can be considered as an increase in oxidation number and reduction can
be considered as a decrease in oxidation number.
87
56
31
6.
Compounds containing metals in high oxidation states tend to be oxidising agents
whereas compounds with metals in low oxidation states are often reducing agents.
89
57
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Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Transition Metal Complexes
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
A complex consists of a central metal ion surrounded by ligands
91
59
202
2.
Ligands are electron donors and may be negative ions or molecules with
non-bonding pairs of electrons. Ligands can be classified as monodentate,
bidentate, etc.
91
59
202
3.
The number of bonds from the ligand to the central metal ion is known as the
co-ordination number of the central ion.
92
60
257
4.
Complexes are written and named according to I.U.P.A.C. rules.
93
60 – 61
202
5.
In a complex of a transition metal the d orbitals are no longer degenerate.
100
62
260
6.
The energy difference between subsets of d orbitals depends on the position of the
ligand in the spectrochemical series.
105
63
203
7.
Colours of many transition metal complexes can be explained in terms of d-d
transitions.
99 – 100
62 – 64
259 – 260
PPA
PPA 1: Preparation of Potassium Trioxalatoferrate(III).
( iii)
Transition Metal Complexes
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
100
64
384
100 – 101
65
385
100
65
101
65
385
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
U.V. And Visible Spectroscopy
LEARNING OUTCOME
1.
The effects of d-d transitions can be studied using ultra-violet and visible
absorption spectroscopy.
2.
Ultra-violet and visible absorption spectroscopy involve transitions between
electronic energy levels in atoms and molecules where the energy difference
corresponds to the ultra- violet and visible regions of the electromagnetic
spectrum.
3.
PPA
The wavelength ranges are approximately 200-400 nm for ultra-violet and
400-700 nm for visible.
An ultra-violet / visible spectrometer measures the intensity of radiation transmitted
through the sample and compares this with the intensity of incident radiation.
PPA 2: Colorimetric Determination of Manganese in Steel.
( iii)
Transition Metal Complexes
4.
Catalysis
LEARNING OUTCOME
1.
Transition metals or their compounds act as catalysts in many chemical reactions
105
66
261
2.
It is believed that the presence of unpaired d electrons or unfilled d-orbitals allows
intermediate complexes to form, providing reaction pathways of lower energy
compared to the uncatalysed reaction
106
67
262
3.
The variability of oxidation states of transition metals is also an important factor.
106
68
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Advanced Higher Chemistry
Student Learning Outcomes
ADVANCED HIGHER CHEMISTRY STUDENT LEARNING OUTCOMES
UNIT 2 – PRINCIPLES OF CHEMICAL REACTIONS
(a)
STOICHIOMETRY
Stoichiometry
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
15 – 16
1.
A quantitative reaction is one in which the substances react completely according
to the mole ratios given by the balanced (stoichiometric) equation.
3–9
1
2.
Volumetric analysis involves using a solution of accurately known concentration in
a quantitative reaction to determine the concentration of another substance.
9
1
3.
A solution of accurately known concentration is known as a standard solution.
A standard solution can be prepared directly from a primary standard.
10
2
9
2
10 – 13
3
11
4.
A primary standard must have, at least, the following characteristics:

high state of purity;

very stable;

high solubility;

reasonably high formula mass.
5.
The volume of reactant solution required to just complete the reaction is
determined by titration.
6.
The equivalence point is the point at which the reaction is just complete. The ‘end
point’ is the point at which a change is observed and is associated with the
equivalence point. An indicator is a substance which changes colour at the endpoint.
10
2
331
7.
Acid / base titrations are based on neutralisation reactions
9
3
11 - 12
8.
Complexometric titrations are based on complex formation reactions.
14
3
11 – 13
4
15 – 16
7-8
E.D.T.A. is an important complexometric reagent and can be used to determine
the concentration of metal ions such as nickel(II).
PPA
PPA 1: The Complexometric Determination of Nickel using EDTA.
9.
Redox titrations are based on redox reactions. Substances such as potassium
manganate(VII) which can act as their own indicators are very useful reactants in
redox titrations
In gravimetric analysis the mass of an element or compound present in a
substance is determined by chemically changing that substance into some other
substance of known chemical composition, which can be readily isolated, purified
and weighed.
PPA 2: The Gravimetric Determination of Water in Hydrated Barium Chloride.
10.
PPA
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Advanced Higher Chemistry
Student Learning Outcomes
(b)
CHEMICAL EQUILIBRIUM
( i)
Reactions at Equilibrium
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
A chemical reaction is in equilibrium when the composition of the reactants and
products remains constant indefinitely.
22
9
289 – 290
2.
The equilibrium constant (K ) characterises the equilibrium composition of the
reaction mixture.
24
10
299
3.
The equilibrium constant can be measured in terms of concentrations or, for
gaseous reactions, in terms of pressure.
For the general reaction:
26
10
300 – 301
24
10
298
4.
a A + b B 
cC+dD
[C] c [D] d
K=
[A] a [B] b
where [A], [B], [C] and [D] are the equilibrium concentrations of A, B, C and D,
respectively, and a, b , c and d are the stoichiometric coefficients in a balanced
reaction equation.
5.
In a homogeneous equilibrium all the species are in the same phase.
27
10
303
6.
In a heterogeneous equilibrium the species are in more than one phase. The
concentrations of pure solids or pure liquids are constant and are given the value 1
in the equilibrium equation.
27
10
303
7.
Equilibrium constants are independent of the particular concentrations or
pressures of species in a given reaction.
26
10
299
8.
Equilibrium constants depend on the reaction temperature.
34
13
299
9.
Le Chatelier’s principle states that when a reaction at equilibrium is subject to
change the composition alters in such a way as to minimise the effects of that
change.
33
13
306
10.
For endothermic reactions a rise in temperature causes an increase in K,
i.e., the yield of the product is increased.
34
14
308
11.
For exothermic reactions a rise in temperature causes a decrease in K,
i.e., the yield of the product is decreased.
34
14
308
12.
The effects of changes in concentration or pressure on the position of equilibrium
can be explained quantitatively in terms of a fixed equilibrium constant.
32 – 33
11 + 15
13.
The presence of a catalyst does not affect the equilibrium constant.
35 – 36
14
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Advanced Higher Chemistry
( ii)
Student Learning Outcomes
Equilibria Between Different Phases
Partition Coefficient
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
When a solute is shaken in two immiscible liquids it partitions itself between the
two liquids in a definite ratio called the partition coefficient.
40
17
294
2.
The value of the partition coefficient depends on the immiscible liquids involved,
the solute and the temperature.
41
17
295 – 296
PPA
PPA 3: Determination of a Partition Coefficient.
( ii)
Equilibria Between Different Phases
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
42
17
296
43
17
296
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
47
20
47 – 50
20
47 – 48
20 – 22
49
22 – 23
Solvent Extraction
LEARNING OUTCOME
1.
2.
Solvent extraction is an application of the partition of a solute between two liquids.
Applications of solvent extraction include the purification of water-soluble organic
acids using a suitable organic solvent.
( ii)
Equilibria Between Different Phases
Chromatography
LEARNING OUTCOME
1.
Chromatographic separations depend on the partition equilibrium between two
phases, one stationary and the other mobile.
2.
There are several types of chromatography. Examples are:
paper chromatography; gas-liquid chromatography
3.
In paper chromatography, the stationary phase is the water held on the paper and
the mobile phase is another solvent.
In gas-liquid chromatography the stationary phase is a liquid held on a solid
support and the mobile phase is a gas.
4.
533571242
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Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Equilibria Involving Ions
Acid / Base Equilibria
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
The Bronsted - Lowry definitions of acid and base state that an acid is a proton
donor and a base is a proton acceptor.
54
23
199
2.
For every acid there is a conjugate base, formed by loss of a proton.
55
23
199
3.
For every base there is a conjugate acid, formed by gain of a proton.
The ionisation of water can be represented by:
55
23
199
59
24
199
56
24
470
59
24
326
4.
H2O(l) + H2O(l)
5.
6.
H3O+(aq) + OH─ (aq)
Water is amphoteric.
The dissociation constant for the ionisation of water is known as the ionic product
and is represented by:
Kw = [H3O+][OH─]
7.
The value of the ionic product varies with temperature.
60
25
326
8.
At 25 oC the value of Kw is approximately 1 x 10-14 mol 2 l –2.
61
25
326
9.
A shorthand representation of H3O+ is H+. Stoichiometric equations and
58
25
198
61
25
325
64
26
329
64
26
329
55
26
199
equilibrium expressions can be written using H+ instead of H3O+ where the
meaning is clear.
10.
11.
12.
13.
14.
15.
16.
The relationship between pH and the hydrogen ion concentration is given by:
pH = -log10 [H+]
The acid dissociation constant of acid HA is given by:
HA(aq) + H2O(l)
H3O+ (aq) + A- (aq)
The acid dissociation constant of acid HA is given by:
[ H3O  ][A  ]
Ka =
[HA]
The conjugate base of an acid of general formula HA is A─.
The dissociation constant of an acid can be represented by pKa where
pKa = - logK a
The relationship of the pH of a weak acid to its dissociation constant is given by:
pHa = ½ pKa - ½ log c
The dissociation in aqueous solution of base of general formula B can be
represented as:
B( aq) + H2O(l)
BH+ (aq) + OH─ (aq)
17.
The conjugate acid of a base of general formula B is BH+.
18.
The dissociation of the conjugate acid of the base can be represented as:
BH+ (aq) + H2O(l)
533571242
B(aq) + H3O+ (aq)
Page 14 of 31
65
66
27
68
28
66
28
67
28
330
Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Equilibria Involving Ions
Acid / Base Equilibria
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
67
28
330
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
72
29
331
72
30
331
In aqueous solution the colour of the acid is distinctly different from that of its
conjugate base.
72
30
331
23.
The colour of the indicator is determined by the ratio of [HIn ] to [In─].
72
30
331
24.
The theoretical point at which colour changes occurs when [H+]= KIn.
75
30
332
25.
75
30
26.
The colour change is assumed to be distinguishable when [HIn] and [In─] - differ by
a factor of 10.
The pH range over which a colour change occurs can be estimated by the
expression:
pH = pKIn ±1
75
30
( iii)
Equilibria Involving Ions
LEARNING OUTCOME
19.
( iii)
The dissociation constant for the conjugate acid is:
[B][H 3O  ]
Ka =
[BH  ]
Equilibria Involving Ions
Indicators
LEARNING OUTCOME
20.
Indicators are weak acids for which the dissociation can be represented as:
HIn(aq) + H2O(l)
21.
22.
H3O + (aq) + In─ (aq)
The acid dissociation constant is represented as K In and is given by the following
expression:
[H3O  ][In ]
KIn =
[HIn]
533571242
Page 15 of 31
Advanced Higher Chemistry
Student Learning Outcomes
Buffer Solutions
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
27.
A buffer solution is one in which the pH remains approximately constant when
small amounts of acid or base are added.
79
33
334
28.
An acid buffer consists of a solution of a weak acid and one of its salts.
79
33
334
29.
In an acid buffer solution the weak acid can supply hydrogen ions when these are
removed by the addition of a small amount of base. The salt of the weak acid
provides the conjugate base, which can absorb excess hydrogen ions produced by
the addition of a small amount of acid.
80
33
335
30.
A basic buffer consists of a solution of a weak base and one of its salts.
79
33
334
31.
In a basic buffer solution the weak base removes excess hydrogen ions and the
conjugate acid provided by the salt supplies hydrogen ions when these are
removed.
82
33
83 – 84
34
335
84 – 85
35
335 – 336
32.
The pH of an acid buffer solution can be calculated from its composition and from
the acid dissociation constant.
[acid]
[H3O+(aq)] = Ka ×
[salt]
or
pH = pKa – log
33.
[acid]
[salt]
The required compositions of an acid buffer solution can be calculated from the
desired pH and from the acid dissociation constant.
533571242
Page 16 of 31
Advanced Higher Chemistry
(c)
THERMOCHEMISTRY
( i)
Hess’s Law
Student Learning Outcomes
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
Thermochemistry concerns the study of changes in energy which occur during
chemical reactions.
90
37
146
2.
The First Law of Thermodynamics states that energy is conserved.
94
38
147
3.
Hess’s law states that the overall reaction enthalpy is the sum of the reaction
enthalpies of each step of the reaction. This is an application of the First Law of
Thermodynamics.
93 – 94
38
151
4.
A thermochemical cycle can be used to calculate an unknown enthalpy value.
94
38
152 – 153
The term ‘standard enthalpy change’
refers to an enthalpy change for a
reaction in which the reactants and products are considered to be in their standard
states at a specified temperature.
95
37
148
6.
The standard state of a substance is the most stable state of the substance under
standard conditions.
95
37
7.
Standard conditions refer to a pressure of one atmosphere and a specified
temperature, usually 298 K (25 °C).
95
37
148
8.
The standard molar enthalpy of combustion refers to the enthalpy change which
occurs when one mole of a substance is burned completely.
95
37
150
9.
Calorimetry is the term used to describe the quantitative determination of the
change in heat energy which occurs during a chemical reaction.
A calorimeter is used to measure the quantity of heat energy given out or taken in
during a chemical reaction.
The standard molar enthalpy of formation refers to the enthalpy change which
occurs when one mole of a substance is prepared from its elements in their
standard states.
The standard enthalpy of formation of a substance can be calculated from
standard enthalpy changes which are experimentally determined.
98
40
150
99
37
150
95
37
148
95 – 96
38 – 39
151
97
39
153
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
For a diatomic molecule, XY, the molar bond enthalpy is the energy required to
break one mole of XY bonds, that is, for the process:
X─Y(g)  X(g) + Y(g)
91
41
2.
Mean molar bond enthalpies are average values which are quoted for bonds which
occur in different molecular environments.
93
41
3.
Bond enthalpies may be calculated from data on enthalpy changes.
92
42
4.
The enthalpy of a reaction can be estimated from a thermochemical cycle involving
bond formation and bond dissociation
Enthalpies of reaction estimated from bond enthalpies may differ from
experimentally determined values.
94
42 – 43
158
93
44
159
5.
10.
11.
12.
13.
( ii)
(Ho)
The standard enthalpy of a reaction can be calculated from tabulated standard
molar enthalpies of formation using the relation:
Ho =
H of (products) - H of (reactants)
Bond Enthalpies
LEARNING OUTCOME
1.
5.
533571242
Page 17 of 31
158
Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Hess’s Law Applied to Ionic Substances
Born – Haber Cycle
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
100
45 – 46
161
101 – 102
45
162
103
45
160
100
46
161
102 – 103
46
161
102
46
149
103
46
85
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
A thermochemical cycle can represent the relation between enthalpy of solution,
enthalpy of lattice formation and enthalpy of hydration for the solution of an ionic
compound.
105 – 106
47 – 48
163
105
47
163
LEARNING OUTCOME
1.
2.
3.
4.
5.
6.
The Born – Haber cycle is a thermochemical cycle applied to the formation of an
ionic crystal.
The Born – Haber cycle can be used to calculate the enthalpy of lattice formation,
which cannot be determined directly by experiment.
The standard molar enthalpy change of lattice formation is the enthalpy change
which occurs when one mole of an ionic crystal is formed from the ions in their
gaseous states under standard conditions.
The cycle is a closed path which includes as steps the different enthalpy changes
involved in the formation of an ionic crystal.
The different enthalpy changes include enthalpy of atomisation, ionisation energy,
bond enthalpy, electron affinity, lattice enthalpy and enthalpy of formation.
The standard molar enthalpy of atomisation of an element is the energy required to
produce one mole of isolated gaseous atoms from the element in its standard
state.
e.g.
7.
½ I 2 (s) I(g)
The electron affinity is usually defined as the enthalpy change for the process of
adding one mole of electrons to one mole of isolated atoms in the gaseous state,
ie, for the change represented by:
E(g) + e -
(iii)
 E - (g)
Hess’s Law Applied to Ionic Substances
Enthalpy of Solution
1.
2.
The hydration enthalpy is the energy released when one mole of individual
gaseous ions becomes hydrated, i.e., the changes represented by:
E n+ (g)  E n+ (aq) and
533571242
E n- (g)  E n- (aq)
Page 18 of 31
Advanced Higher Chemistry
(d)
( i)
Student Learning Outcomes
REACTION FEASIBILITY
Entropy
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
114
49
343
115 – 117
49
344
343
1.
The entropy (S) of a system is the degree of disorder of the system. The greater
the disorder, the greater the entropy.
2.
Entropy increases as temperature increases.
3.
Changes of state involve changes in entropy. Melting and evaporation are
accompanied by increases in entropy.
115
49 – 50
4.
One version of the Third Law of Thermodynamics states that the entropy of a
perfect crystal at 0 K is zero.
115
50
5.
The standard entropy of a substance is the entropy value for the standard state of
the substance.
The change in standard entropy for a reaction system can be calculated from the
standard entropy values for the reactants and products.
117
51
343
117 – 118
51
344
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
One version of the Second Law of Thermodynamics states that the total entropy of
a reaction system and its surroundings always increases for a spontaneous
process.
120
51
341
2.
Heat energy released by the reaction system into the surroundings increases the
entropy of the surroundings, whereas heat absorbed by the reaction system from
the surroundings decreases the entropy of the surroundings.
119
51
344
3.
The change in entropy of the surroundings that occurs as a result of a chemical
reaction can be calculated from the temperature and from the enthalpy change for
the reaction system.
120
51
345
The total entropy change is proportional to the change in free energy (G) of the
reaction system. The direction of spontaneous change is in the direction of
decreasing free energy.
122
52
346
122
52
346
346
6.
( ii)
Free Energy
LEARNING OUTCOME
1.
4.
5.
The change in standard free energy for a reaction is related to the standard
enthalpy and entropy changes by:
G o = H o - TS o
6.
The standard free energy change of a reaction can be calculated from the
standard enthalpy and standard entropy changes for the reaction.
124
52
7.
The standard free energy change of a reaction can be calculated from the
standard free energies of formation of the reactants and products.
124
53
8.
A reaction is feasible under standard conditions if the change in standard free
energy between reactants and products is negative. This means that the
equilibrium composition favours the products over the reactants.
125
54 – 56
9.
Under non-standard conditions any reaction is feasible if G is negative.
123
52
10.
At equilibrium G = 0.
128
53
11.
A reaction will proceed spontaneously in the forward direction until the composition
is reached where G = 0.
127
53
533571242
Page 19 of 31
Advanced Higher Chemistry
( ii)
Student Learning Outcomes
Free Energy
Application of the Concept of Free Energy
LEARNING OUTCOME
SCHOLAR
12.
The feasibility of a chemical reaction under standard conditions can be predicted
from the calculated value of the change in standard free energy (G o).
13.
The temperature at which the reaction becomes feasible can be calculated for a
reaction for which both H o and S o have positive values.
PPA
PPA 4: Verification of a Thermodynamic Prediction.
14.
LTS
NOTES
CHEM in
CONTEXT
56 – 57
131
56 – 57
Ellingham diagrams are plots of variation of free energy change with temperature
and can be used to predict the conditions under which a reaction can occur.
131
57 – 61
15.
Ellingham diagrams can be used to predict the conditions required to extract a
metal from its ores.
131
57 – 61
(e)
ELECTROCHEMISTRY
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
A potential difference is set up when a metal is placed in contact with its ions in
solution.
138 – 139
63
22 – 23
2.
An electrochemical cell is composed of two half cells between which electrical
contact is made by an electrolyte, often in the form of a salt bridge.
140
63
22 – 23
3.
Cell and cell emf conventions should be employed according to I.U.P.A.C.
recommendations.
141 – 144
64
188
148
65
194
141 – 142
66
188
144
64
191
144 – 145
65
191
Electrochemistry
For example, a cell reaction between zinc and copper can be represented as:
Zn(s) Zn2+ (aq) ║ Cu2+ (aq)Cu(s)
The equation for this cell is written as:
Zn(s) + Cu2+(aq) 
Zn2+(aq)
+
Cu(s)
4.
A positive emf is obtained if the reaction takes place in the direction as written.
5.
The emf of a cell (E) is the electric potential difference between the electrodes of
the cell, i.e., E(right) - E(left) when no current is drawn.
6.
Cell emf depends on the concentration, the temperature and the type of cell.
7.
Cell emf values are usually considered at standard conditions.
8.
Standard conditions for the measurement of electrode potentials refer to a situation
in which all pressures are one atmosphere, concentrations of solutions are one
mole per litre and in which temperature is specified normally at 298 K (25 °C).
144
65
191
9.
The absolute value of the electrode potential of a half cell cannot be determined
experimentally.
145
64
190
10.
The standard electrode potential of a half cell is the potential measured against the
standard hydrogen electrode under standard conditions.
145 – 147
64
190
11.
The standard hydrogen electrode potential is given an arbitrary value of 0·00V.
145
64
190
148
66
194 – 195
12.
The emf of a cell under standard conditions
can be calculated from the
tabulated values of standard reduction potentials.
533571242
(Eo)
Page 20 of 31
Advanced Higher Chemistry
Student Learning Outcomes
Electrochemistry
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
13.
The relative strengths of reducing agents and oxidising agents under standard
conditions can be estimated from standard reduction potentials.
148 – 150
65
192 – 193
14.
For a standard cell operated under conditions of thermodynamic reversibility the
standard free energy change for the cell reaction is related to cell emf by the
expression:
151
70
349
154 – 155
71 -72
197 – 198
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
160
75
360
165 – 167
77
360 – 361
G o = -nFE o
15.
(f)
A fuel cell operates like an electrochemical cell, the only difference being that the
fuel for the reaction is provided from external reserves of gas, e.g., the hydrogen /
oxygen fuel cell.
KINETICS
Kinetics
LEARNING OUTCOME
1.
2.
The rate of a chemical reaction normally depends on the concentrations of the
reactants.
For a first order reaction the rate of reaction is proportional to the concentration of
one reactant and the rate can be expressed as:
rate = k [A]
where k is the rate constant and [A] is the concentration of reactant A in mol l -1 .
3.
The order of a reaction with respect to any one reactant is the power to which the
concentration of that reactant is raised in the rate equation.
165 – 167
77
361
4.
The overall order of a reaction is the sum of the powers to which the
concentrations of the reactants are raised in the rate equation.
165 – 167
77
363
5.
In general for a reaction of type:
165 – 167
77
363
166 – 167
78
362
170
80
373
171 – 172
80
374
172
80 – 81
nA + mB product
where the rate equation is of the form:
rate = k [A ]n[B]m
the order of reaction is n with respect to A and m with respect to B and the overall
order is n + m.
6.
The rate constant can be determined from initial rate data for a series of reactions
in which the initial concentrations of reactants are varied.
7.
Reaction mechanisms usually occur by a series of steps.
8.
The rate of reaction is dependent on the slowest step which is called the ‘rate
determining step’.
Experimentally determined rate equations can provide evidence for a proposed
reaction mechanism but cannot provide proof, as other possible reaction
mechanisms may also give the same rate equation.
PPA 5: Kinetics of the Acid - Catalysed Propanone / Iodine Reaction.
9.
PPA
ADVANCED HIGHER CHEMISTRY STUDENT LEARNING OUTCOMES
533571242
Page 21 of 31
Advanced Higher Chemistry
Student Learning Outcomes
UNIT 3 – ORGANIC CHEMISTRY
(a)
PERMEATING ASPECTS OF ORGANIC CHEMISTRY
( i)
Reaction Types
LEARNING OUTCOME
1.
( ii)
Equations can be written for the following reaction types and, given equations,
these reaction types can be identified as:

substitution;

addition;

elimination;

condensation;

hydrolysis;

oxidation;

reduction.
( iii)
CHEM in
CONTEXT
5
3
See index
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
15 – 17
10 – 11
407 – 408
The following reaction mechanisms can be described in terms of electron shifts:
(i)
radical substitution of alkanes;
(ii)
electrophilic addition to alkenes
carbocation mechanism
cyclic ion intermediate mechanism
23 – 26
15 – 19
421
(iii)
nucleophilic substitution
SN1 and SN2
32 – 33
25 – 28
456 – 459
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
44
32 – 33
103
Physical Properties
LEARNING OUTCOME
1.
LTS
NOTES
Reaction Mechanisms
LEARNING OUTCOME
1.
SCHOLAR
The following physical properties are explained in terms of the intermolecular
forces involved:

melting and boiling points;

miscibility with water.
533571242
Page 22 of 31
Advanced Higher Chemistry
Student Learning Outcomes
(b)
SYSTEMATIC ORGANIC CHEMISTRY
( i)
Hydrocarbons and Halogenoalkanes
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
12 – 13
8
CHEM in
CONTEXT
1.
Bonding in alkanes can be described in terms of sp3 hybridisation and sigma
bonds.
2.
Hybridisation is the process of mixing atomic orbitals in an atom to generate a set
of new atomic orbitals called hybrid orbitals.
13
8
3.
A sigma bond is a covalent bond formed by end-on overlap of two atomic orbitals
lying along the axis of the bond.
13
10
419
4.
Alkanes undergo substitution reactions with chlorine and bromine by a chain
reaction mechanism.
The chain reaction includes the following steps:
15
10 – 11
407 – 408
15 – 16
10 – 11
407 – 408
19 – 20
13-14
419
20
14
419
20 – 22
14 – 15
473
23 – 27
15 – 21
422 – 425
23 – 27
17,19 + 21
424
453
5.
(i)
initiation by homolytic fission to produce radicals;
(ii)
propagation;
(iii)
termination.
6.
Bonding in ethene can be described in terms of sp2 hybridisation, sigma and pi ()
bonds.
7.
A pi () bond is a covalent bond formed by the sideways overlap of two parallel
atomic orbitals lying perpendicular to the axis of the bond.
8.
Alkenes can be prepared in the laboratory by:
(i)
dehydration of alcohols using aluminium oxide, concentrated sulphuric acid
or orthophosphoric acid;
(ii)
base-induced elimination of hydrogen halides from
monohalogenoalkanes.
PPA
PPA 1: Preparation of Cyclohexene.
9.
Alkenes undergo:
10.
(i)
catalytic addition with hydrogen to form alkanes;
(ii)
addition with halogens to form dihalogenoalkanes;
(iii)
addition with hydrogen halides according to Markovnikov’s rule to form
monohalogenoalkanes;
(iv)
acid-catalysed addition with water according to Markovnikov’s rule to form
alcohols.
The mechanisms of the above reactions involve:
(i)
for halogenation

cyclic ion intermediate
(ii)
for hydrohalogenation

carbocation intermediate
(iii)
for acid catalysed hydration

carbocation intermediate
11.
Halogenoalkanes are named according to I.U.P.A.C. rules.
28
23
12.
Monohalogenoalkanes can be classified as primary, secondary or tertiary.
29
24 – 25
13.
Monohalogenoalkanes undergo nucleophilic substitution reactions. React
monohalogenoalkanes with alkali and test for halide ion using aqueous ethanolic
silver nitrate solution.
32 – 34
25 – 26
533571242
Page 23 of 31
456
Advanced Higher Chemistry
( i)
Student Learning Outcomes
Hydrocarbons and Halogenoalkanes
LEARNING OUTCOME
14.
Monohalogenoalkanes undergo elimination reactions to form alkenes.
15.
Monohalogenoalkanes react with:
( ii)
(i)
alkalis to form alcohols;
(ii)
alcoholic alkoxides to form ethers;
(iii)
ethanolic cyanide to form nitriles which can be hydrolysed to carboxylic
acids (chain length increased by one carbon atom);
(iv)
ammonia to form amines via alkyl ammonium salts.
LTS
NOTES
CHEM in
CONTEXT
30 – 31
30 – 31
462
34 – 36
28 – 29
460 – 462
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
44 – 45
32 – 33
468
47
33
469
48 – 49
34
456,
424 – 425
50 – 51
34
424 – 425
Alcohols and Ethers
LEARNING OUTCOME
1.
Alcohols exhibit hydrogen bonding and as a result have higher boiling points than
other organic compounds of comparable relative formula mass and shape.
2.
The lower alcohols are miscible with water but as their chain length increases their
solubility in water decreases.
Alcohols can be prepared from:
(i)
alkenes by hydration;
3.
SCHOLAR
(ii)
halogenoalkanes by substitution.
4.
In industry, alcohols (except methanol) can be manufactured by the acid-catalysed
hydration of alkenes.
5.
Alcohols react with some reactive metals to form alkoxides.
51
29 + 34
471
6.
Alcohols can be dehydrated to alkenes.
52
35
473 – 474
7.
Alcohols undergo condensation reactions slowly with carboxylic acids and more
vigorously with acid chlorides to form esters.
53
36
472
8.
Ethers have the general formula R'-O-R ''where R 'and R ''are alkyl groups.
43
37
478
9.
Ethers are named according to I.U.P.A.C. rules.
43
37
10.
Due to the lack of hydrogen bonding, ethers have lower boiling points than the
corresponding isomeric alcohols.
46
37
11.
Ether molecules can hydrogen-bond with water molecules thus explaining the
solubility in water of some ethers of low relative formula mass.
47
38
12.
Ethers are highly flammable and on exposure to air may form explosive peroxides.
55
38
478
13.
Ethers can be prepared by the reaction of halogenoalkanes with alkoxides.
Ethers are used as solvents since they are relatively inert chemically and will
dissolve many organic compounds.
54
29
461
55
38
478
14.
533571242
Page 24 of 31
478
Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Aldehydes, Ketones and Carboxylic Acids
LEARNING OUTCOME
1.
LTS
NOTES
CHEM in
CONTEXT
60 – 62
41 – 42
483 – 484
65 – 66
42
489
67
43
486
68 – 70
43
484 – 485,
487
69
46
487
The following physical properties of aldehydes and ketones can be explained in
terms of dipole-dipole attractions and / or hydrogen bonding:
(i)
higher boiling points than corresponding alkanes;
(ii)
lower boiling points than corresponding alcohols;
(iii)
miscibility of lower members with water.
2.
Tollens’ reagent or Fehling’s solution can be used to distinguish between
aldehydes and ketones. Aldehydes reduce the complexed silver(I) ion and the
complexed copper(II) ion to silver and copper(I) oxide, respectively.
3.
Aldehydes and ketones can be reduced to primary and secondary alcohols,
respectively, by reaction with lithium aluminium hydride in ether.
Aldehydes and ketones undergo:
4.
SCHOLAR
(i)
(ii)
nucleophilic addition with HCN to form cyanohydrins which can be
hydrolysed to hydroxy carboxylic acids.
nucleophilic addition-elimination with hydrazine,
2, 4-dinitrophenylhydrazine to form hydrazones and
2,4-dinitrophenylhydrazones respectively.
5.
These nucleophilic addition-elimination reactions are also described as
condensation since water is formed in the process.
PPA
PPA 2: Identification by Derivative Formation.
6.
The melting points of the resulting 2, 4-dinitrophenylhydrazones are used to
identify unknown aldehydes and ketones.
71
46 – 47
487
7.
Aldehydes are generally more reactive than ketones because the presence of two
alkyl groups in ketones hinders nucleophilic attack and reduces the partial positive
charge on the carbonyl carbon atom.
71
42
484
8.
In pure carboxylic acids hydrogen bonding produces dimers thus explaining the
relatively high boiling points. Dimerisation does not occur in aqueous solution.
63
48
499
9.
Carboxylic acid molecules also form hydrogen bonds with water molecules thus
explaining the appreciable solubility of the lower carboxylic acids in water. As the
chain length increases water solubility decreases.
62
49
10.
Carboxylic acids are weak acids. Their slight dissociation in water can be
explained by the stability of the carboxylate ion caused by electron delocalisation.
63 – 64
49
500
11
Carboxylic acids can be prepared by:
72
50
499
(i)
oxidising primary alcohols and aldehydes;
(ii)
hydrolysing nitriles, esters or amides.
533571242
Page 25 of 31
Advanced Higher Chemistry
( iii)
Student Learning Outcomes
Aldehydes, Ketones and Carboxylic Acids
LEARNING OUTCOME
12.
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
73 – 76
51 – 52
501,
504 – 505
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
82 – 83
53
84
53
516
Reactions of carboxylic acids include:
(i)
formation of salts by reactions with metals, carbonates and alkalis;
(ii)
condensation reactions with alcohols to form esters;
(iii)
reaction with ammonia or amines and subsequent heating of the
ammonium salt to form amides;
(iv)
reduction with lithium aluminium hydride to form primary alcohols.
PPA
PPA 3: Preparation of Benzoic Acid by Hydrolysis of Ethyl Benzoate.
( iv )
Amines
LEARNING OUTCOME
1.
Amines are named according to I.U.P.A.C. rules.
2.
Amines are organic derivatives of ammonia and can be classified as primary,
secondary or tertiary.
3.
Primary and secondary amines, but not tertiary amines, associate by hydrogen
bonding and as a result have higher boiling points than isomeric tertiary amines
and alkanes with comparable relative formula masses.
87 – 88
55
516
4.
Amine molecules can hydrogen- bond with water molecules thus explaining the
appreciable solubility of the lower amines in water.
89
55
517
5.
The nitrogen atom in amines has a lone pair of electrons, which can accept a
proton from water, producing hydroxide ions. Amines are weak bases.
90
55
519
6.
Amines react with aqueous mineral or carboxylic acids to form salts.
91 – 92
55 – 56
519
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Advanced Higher Chemistry
( v)
Student Learning Outcomes
Aromatics
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
Bonding in benzene can be described in terms of sp2
pi () bonds, and electron delocalisation.
100
59
438
2.
Benzene is the simplest aromatic hydrocarbon and its unexpected stability can be
attributed to the presence of delocalised electrons.
102
59
437 – 438
3.
Most reactions of benzene involve attack of an electrophile on the cloud of
delocalised electrons, that is electrophilic substitution.
102
60
441
4.
Benzene resists addition reactions but undergoes electrophilic substitution
reactions. These include:
103 – 107
60 – 63
442 – 445
109 – 112
63 – 65
471
5.
(i)
chlorination and bromination to produce chlorobenzene and
bromobenzene;
(ii)
nitration to produce nitrobenzene;
(iii)
sulphonation to produce benzene sulphonic acid;
(iv)
alkylation to produce alkylbenzenes.
The presence of delocalised electrons in the phenyl group can be used to explain:
(i)
the stronger acidic nature of phenol compared to aliphatic alcohols;
(ii)
the weaker basic nature of the aromatic amine, aminobenzene
(aniline),compared with aliphatic amines.
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Advanced Higher Chemistry
(c)
Student Learning Outcomes
STEREOISOMERISM
Stereoisomers have identical molecular formulae and the atoms are bonded together in the same order but the arrangement of
the atoms in space is different, making them non – superimposable.
( i)
Geometric Isomerism
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
1.
Geometric isomerism is one type of stereoisomerism and can arise due to the lack
of free rotation around a bond, frequently a carbon – carbon double bond.
118
68
420
2.
Geometric isomers are labelled cis and trans according to whether the substituent
groups are on the same side or on different sides of the carbon – carbon double
bond.
119
68
420
3.
Geometric isomers display differences in some physical properties.
120
69
420
4.
Geometric isomerism can also influence chemical properties, for example
cis -but-2- enedioic acid is more readily dehydrated than
trans -but- 2- enedioic acid.
121 – 122
69 – 70
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
127
71 – 72
392
128 – 129
71
392 – 393
( ii)
Optical Isomerism
LEARNING OUTCOME
1.
Optical isomers are non-superimposable mirror images of each other and are said
to be chiral.
2.
Optical isomerism can occur in substances in which four different groups are
arranged around a carbon atom.
3.
Optical isomers have identical physical and chemical properties, except when they
are in a chiral environment, but they have an opposite effect on plane polarised
light and are said to be optically active.
132
72
392 – 393
4.
Mixtures containing equal amounts of both optical isomers are optically inactive.
134
72 – 73
393
5.
In biological systems only one optical isomer of each organic compound is usually
present.
129
73
393
(d)
STRUCTURAL ANALYSIS
( i)
Elemental Microanalysis and Mass Spectroscopy
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
Elemental Microanalysis
LEARNING OUTCOME
1.
Elemental microanalysis (combustion analysis) can be used to determine the
masses of C, H, S and N in a sample of an organic compound in order to find the
empirical formula.
140
75
2.
Other elements in the organic compound have to be determined by other methods.
140
75
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Advanced Higher Chemistry
( i)
Student Learning Outcomes
Elemental Microanalysis and Mass Spectroscopy
Mass Spectroscopy
LEARNING OUTCOME
1.
Mass spectrometry can be used to determine the accurate molecular mass and
structural features of an organic compound.
2.
A conventional mass spectrometer functions in the following manner:

The sample is firstly vaporised and then ionised by being bombarded with
electrons.

Fragmentation can occur when the energy available is greater than the
molecular ionisation energy.

The parent ion and ion fragments are accelerated by an electric field and
then deflected by a magnetic field.

The strength of the magnetic field is varied to enable the ions of all the
different mass/ charge ratios to be detected in turn. A mass spectrum is
obtained.
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
142
76
2–4
142 – 143
77
2–4
144 – 148
78
383
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
166
80
385
3.
Organic compounds can be identified from the very accurate determination of the
relative molecular masses of the parent ion and the ion fragments.
(ii)
Infra – red and Nuclear Magnetic Resonance Spectroscopy and X – ray Crystallography
Infra – red Spectroscopy
LEARNING OUTCOME
1.
Infra – red spectroscopy can be used to identify certain functional groups in an
organic compound.
2.
Infra-red radiation causes parts of a molecule to vibrate. The wavelengths which
are absorbed and cause the vibrations will depend on the type of chemical bond
and the groups or atoms at the ends of these bonds.
161 – 162
80
385
3.
Infra- red radiation is passed through a sample of the organic compound and then
to a detector which measures the intensity of the transmitted radiation at different
wavelengths.
163
81 - 82
385
4.
Infra-red spectra are expressed in terms of wavenumber.
159
80 – 81
385
5.
The unit of measurement of wavenumber which is the reciprocal of wavelength is
cm-1.
160
81
385
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Page 29 of 31
Advanced Higher Chemistry
( ii)
Student Learning Outcomes
Infra – red and Nuclear Magnetic Resonance Spectroscopy and X – ray Crystallography
Nuclear Magnetic Resonance Spectroscopy
LEARNING OUTCOME
1.
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
172 – 173
88
388
Nuclear magnetic resonance spectroscopy (nmr) can give information about:
(i)
the different environments of hydrogen atoms in an organic molecule;
(ii)
how many hydrogen atoms there are in each of these environments.
2.
Hydrogen nuclei behave like tiny magnets and in a strong magnetic field some are
aligned with the field (lower energy) while the rest are aligned against it (higher
energy).
170
84 – 85
387
3.
Absorption of radiation in the radiofrequency region of the electromagnetic
spectrum will cause the hydrogen nuclei to ‘flip’ from the lower energy alignment to
the higher one. As they fall back from the higher to the lower level the emitted
radiation is detected.
170 – 171
85
387
4.
In the nmr spectrum the peak position (chemical shift) is related to the environment
of the proton.
173
86
388
5.
The area under the peak is related to the number of protons in that environment.
173
88
388 – 389
6.
The standard reference substance used in nmr spectroscopy is tetramethylsilane
(TMS) which is assigned a chemical shift equal to zero.
172
86
388
(ii)
Infra - red and Nuclear Magnetic Resonance Spectroscopy and X - Ray Crystallography
SCHOLAR
LTS
NOTES
CHEM in
CONTEXT
X – ray Crystallography
LEARNING OUTCOME
1.
X - ray crystallography can be used to determine the precise three-dimensional
structure of organic compounds.
177
90
116
2.
A crystal of an organic compound acts as a diffraction grating when it is exposed to
X - rays of a single wavelength the atoms of the crystal act as a diffraction grating.
177
90
116 – 117
3.
Electron-density maps are produced from the positions and intensities of the
‘spots’ in the diffraction pattern.
177 – 178
91
118
4.
From the electron- density map the precise location of each atom in the molecule
can be determined, and since heavier atoms have more electrons than lighter ones
each atom in the molecule can be identified.
178
91
118
5.
Since a hydrogen atom has a low electron density X-rays do not easily detect it
180
92
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Advanced Higher Chemistry
(e)
Student Learning Outcomes
MEDICINES
Drugs are substances, which alter the biochemical processes in the body, and those, which have a beneficial effect, are called
medicines.
( i)
Historical Development
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
1.
The first medicines were plant brews.
184
93
2.
Pharmacologically active compounds in plant extracts were identified.
184
93
3.
These compounds and derivatives of them were synthesised where practicable.
184
93
4.
Aspirin is an example of a medicine developed in this way.
184
94
PPA
PPA 4: Preparation of Aspirin.
PPA
PPA 5: Aspirin Determination.
( ii)
How a Medicine Functions
LEARNING OUTCOME
SCHOLAR
LTS
NOTES
Most medicines work by binding to receptors. Receptors are usually protein
molecules that are either on the surface of cells where they interact with small
biologically active molecules or are enzymes that catalyse chemical reactions
(catalytic receptors).
186 – 188
95
1.
2.
That structural fragment of the molecule which confers pharmacological activity on
it is called the pharmacophore.
189
98
3.
The shape of the pharmacophore complements that of the receptor site, allowing it
to fit into the receptor. The functional groups on both are correctly positioned to
interact and bind the medicine to the receptor.
189
98
4.
By comparing the structures of medicines with similar pharmacological activity, the
pharmacophore can be identified.
189
98
5.
Many medicines can be classified as agonists or as antagonists according to
whether they enhance or block the body’s natural responses.
187
97
6.
An agonist will produce a response like the body’s natural active compound.
187
97
7.
An antagonist produces no response but prevents the action of the body’s natural
active compound.
188
97
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CHEM in
CONTEXT
CHEM in
CONTEXT