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Williamwood High School
CfE Advanced Higher Chemistry
Learning Outcomes
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Unit 1 Inorganic and Physical Chemistry
Key areas: Electromagnetic radiation and atomic spectra
Electromagnetic waves
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Electromagnetic radiation may be described in terms of waves and can be specified by its
wavelength (  ) in metres or an appropriate sub-multiple (e.g. nanometres) and by its
frequency (f) in s-1
 which is the reciprocal of time in seconds also called the Hertz (Hz).
The velocity (c) of electromagnetic radiation is constant and has a value of approximately
 3 x 108 ms-1.
The electromagnetic spectrum is the range of frequencies or wavelengths of
electromagnetic radiation.
Velocity, frequency and wavelength are related in the expression: c = f  .
Dual nature of electromagnetic radiations- waves and particles
Under certain circumstances electromagnetic radiation may be regarded as a stream of
particles, rather than waves. These particles are known as photons.
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
 E = L h
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for one photon
for one mole of photons
(where L is Avogadro’s constant).
Atomic emission
The emission spectrum of hydrogen provides evidence of energy levels.
Quantum theory states that matter can only emit or absorb energy in small fixed amounts
(called quanta).
The energy of a bound electron in an atom is quantised.
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
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.
Each element produces a unique pattern of frequencies of radiation in its emission
spectrum.
Atomic Spectroscopy
Atomic emission spectroscopy and atomic absorption spectroscopy are used to identify
and quantify the elements present in a sample.
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 400-700nm. Some applications use the ultra-violet region (wavelength region
approximately 200-400nm).
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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.
In absorption spectroscopy electromagnetic radiation is directed at the atomised sample.
Radiation is absorbed as electrons are promoted to higher energy levels. The wavelength
of the absorbed radiation is measured.
Each element provides a characteristic spectrum which can be used to identify an
element.
The concentration of an element within a sample can be determined quantitatively if the
intensity of emitted or transmitted radiation is measured.
Key areas: Atomic orbitals, electronic configurations and the
Periodic Table
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Atomic orbitals
Emission spectra of elements with more than one electron provide evidence of sublevels
within each principal energy level above the first.
Any electron within an atom can be uniquely identified by a set of four quantum numbers:
 The principal quantum number (n)
 The angular momentum quantum number (l)
 The magnetic quantum number (m)
 The spin magnetic quantum number (s)
The principal energy levels correspond to the principal shells (principal quantum number
(n) and takes the values 1, 2, 3……). The second and subsequent principal shells contain
subshells which correspond to the sublevels (second quantum number (l) defines the
subshell type i.e. s, p, d or f).
The types of subshell within each principal shell are as follows:
 First shell
s subshell
 Second shell
s and p subshells
 Third shell
s, p and d subshells
 Fourth shell
s, p, d and f subshells
An atomic orbital is a region or volume of space within which there is a high (95%)
probability of finding an electron.
There are four types of orbitals (within subshells), s, p, d and f, each with characteristic
shapes. Diagrams of d orbitals can be recognised.
Electronic configurations
An orbital holds a maximum of two electrons, as required by the Pauli Exclusion Principle
which states that no two electrons have the same set of four quantum numbers.
The third quantum number (m) determines the actual orbital in which the electron is
located. The number of orbitals in each subshell is as follows:
 s subshell
one s orbital
 p subshell
three p orbitals
 d subshell
five d orbitals
 f subshell
seven f orbitals
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In an isolated atom the orbitals within each subshell are degenerate ( have the same
energy)
The Aufbau Principle states that orbitals are filled in order of increasing energy.
The relative energies corresponding to each orbital can be represented diagrammatically
for the first four shells of a multi-electron atom.
The fourth quantum number (s) is related to the electron’s spin direction.
Hund’s Rule states that when degenerate orbitals are available, electrons fill each singly,
keeping their spins parallel before spin pairing starts.
Electronic configurations using spectroscopic notation and orbital box notation can be
written for elements of atomic numbers 1 to 36.
The Periodic Table and ionisation energies
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.
The variation in first ionisation energy with increasing atomic number for the first 36
elements can be explained in terms of relative stability of different electron configurations,
and so provides evidence for these electronic configurations.
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.
Shapes of molecules and polyatomic ions
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Valence Shell Electron Pair Repulsion Theory (VSEPR)
VSEPR theory provides a quick method to predict the shape of a molecule or polyatomic
ion from the numbers of bonding electron pairs and non-bonding electron pairs.
 Arrangement of electron pairs
Total number of electron pairs
Linear
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Trigonal planar
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Tetrahedral
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Trigonal bipyramidal
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Octahedral
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Electron pair repulsions decrease in strength in the order: non-bonding pair/non-bonding
pair> non-bonding pair/bonding pair> bonding pair/bonding pair.
Differences in strength of electron pair repulsion account for slight deviations from
expected bond angles in molecules such as NH3 and H2O.
Key area: Transition metals
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Electronic configurations of transition metal atoms and ions
The d-block transition metals are metals with an incomplete d subshell in at least one of
their ions.
The filling of the d orbitals follows the Aufbau Principle, with the exception of chromium
and copper atoms. These exceptions are due to special stability associated with all the d
orbitals being half filled or completely filled.
When transition metals form ions it is the s electrons which are lost first rather than the d
electrons.
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Oxidation states of transition metals
An element is said to be in a particular oxidation state when it has a specific oxidation
number.
The oxidation number is determined by following certain rules listed on page ___ of notes.
Transition metals exhibit variable oxidation states of differing stability.
Compounds of the same transition metal but in different oxidation states may have
different colours.
Oxidation can be considered as an increase in oxidation number and reduction can be
considered as a decrease in oxidation number.
Compounds containing metals in high oxidation states tend to be oxidising agents
whereas compounds with metals in low oxidation states are often reducing agents.
Dative covalent bonding of ligands in transition metal complexes
A complex consists of a central metal ion surrounded by ligands.
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.
The total number of bonds from the ligand to the central metal ion is known as the coordination number of the central ion.
Complexes are written and named according to IUPAC rules.
Colour in transition metal complexes
In a complex of a transition metal the d orbitals are no longer degenerate.
The energy difference between subsets of d orbitals depends on the position of the ligand
in the spectrochemical series.
Colours of many transition metal complexes can be explained in terms of d-d transitions.
Light is absorbed when electrons in a lower energy d orbital are promoted to a d orbital of
higher energy.
If light of one colour is absorbed, then the complementary colour will be observed.
UV and visible absorption of transition metal complexes
The effects of d-d transitions can be studied using ultra-violet and visible absorption
spectroscopy.
Ultra-violet and visible absorption spectroscopy involves transitions between electronic
energy levels in atoms and molecules where the energy difference corresponds to the
ultra-violet and visible regions of the electromagnectic spectrum.
The wavelength ranges are approximately 200-400nm for ultra-violet and 400-700nm 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.
Catalysis by transition metals
The presence of unpaired d electrons or unfilled d orbitals allows intermediate complexes
to form, providing reaction pathways with lower activation energies compared to the
uncatalysed reaction.
This allows transition metals or their compounds act as catalysts in many chemical
reactions.
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.
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The variability of oxidation states of transition metals is also an important factor.
Homogeneous and heterogeneous catalysts should be explained in terms of changing
oxidation states.
Key area: Chemical equilibrium
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Equilibrium expressions, factors affecting equilibrium and calculation of the
composition of equilibrium mixtures
A chemical reaction is in equilibrium when the composition of the reactants and products
remains constant indefinitely.
The equilibrium constant (K) characterises the equilibrium composition of the reaction
mixture.
The equilibrium constant can be measured in terms of concentrations or, for gaseous
reactions, in terms of pressure.
For the general reaction:
aA + bB ⇌
cC + dD
The equilibrium constant, K
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= [C]c [D]d
[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.
In a homogeneous equilibrium all the species are in the same phase.
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.
Equilibrium constants are independent of the particular concentrations or pressures of
species in a given reaction.
Equilibrium constants depend on the reaction temperature.
For endothermic reactions a rise in temperature causes an increase in K, i.e., the yield of
the product is increased.
For exothermic reactions a rise in temperature causes a decrease in K, i.e., the yield of
the product is decreased.
The presence of a catalyst does not affect the value of the equilibrium constant.
The equilibrium constants have no units.
The effects of changes in concentration or pressure on the position of equilibrium can be
explained quantitatively in terms of a fixed equilibrium constant.
Ionic product of water and calculations of pH
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The ionisation of water can be represented by:
H2O(l) + H2O(l) ⇌ H3O+ (aq) + OH - (aq)
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Water is amphoteric.
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The dissociation constant for the ionisation of water is known as the ionic product and is
represented by:
Kw = [H3O+][OH -]
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The value of the ionic product varies with temperature.
At 25o C the value of Kw is approximately 1 x 10-14 mol2 l-2 .
A shorthand representation of H3O+ is H+. Stoichiometric equations and equilibrium
expressions can be written using H+ instead of H3O+ where the meaning is clear.
The relationship between pH and the hydrogen ion concentration is given by
pH = -log10 [H+] and conversely [H+] = 10– pH.
Bronsted-Lowry acids and bases
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The Bronsted–Lowry definitions of acid and base state that an acid is a proton donor and
a base is a proton acceptor.
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For every acid there is a conjugate base, formed by loss of a proton.
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For every base there is a conjugate acid, formed by gain of a proton.
A soluble salt of a strong acid and a strong base dissolves in water to produce a neutral
solution.
A soluble salt of a weak acid and a strong base dissolves in water to produce an alkaline
solution.
A soluble salt of a strong acid and a weak base dissolves in water to produce an acidic
solution.
Soaps are salts of weak acids and strong bases.
The pH of a salt solution can be explained by reference to the equilibria between the acids
and alkalis involved.
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Strong and weak acids and bases
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In aqueous solution, strong acids/bases are completely dissociated into ions but weak
acids/bases are only partially dissociated.
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Examples of strong acids include hydrochloric acid, sulfuric acid and nitric acid. Ethanoic,
carbonic and sulfurous acids are examples of weak acids.
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Solutions of metal hydroxides are strong bases. Ammonia and amines are examples of
weak bases.
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The weakly acidic nature of solutions of carboxylic acids, sulfur dioxide and carbon dioxide
can be explained by reference to equations showing the equilibria. The weakly alkaline
nature of a solution of ammonia or amines can be explained by reference to an equation
showing the equilibrium.
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The acid dissociation constant is represented by Ka or by pKa where pKa = –log10Ka
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The dissociation in aqueous solution of an acid of general formula HA can be represented
as:
HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)
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The acid dissociation constant of acid HA is given by:
[H 3 O  ][A  ]
Ka 
[HA]
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The conjugate base of an acid of general formula HA is A-.
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The dissociation constant of an acid can be represented by pKa where
pKa = - log10Ka
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The approximate pH of a weak acid can be calculated from the dissociation constant:
pH = ½ pKa – ½ log10c
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The dissociation in aqueous solution of base of general formula B can be represented as:
BH+(aq) + OH-(aq)
B(aq) + H2O(l)
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The conjugate acid of a base of general formula B is BH+.
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The dissociation of the conjugate acid of the base can be represented as
BH+(aq) + H2O(l)
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B(aq)
+
H3O+(aq)
The dissociation constant for the conjugate acid is
Ka = B H3O+
BH+(aq)
Buffer solutions
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A buffer solution is one in which the pH remains approximately constant when small
amounts of acid or base are added.
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An acid buffer consists of a solution a weak acid and one of its salts.
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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.
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A basic buffer consists of a solution of a weak base and one of its salts.
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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.
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The pH of an acid buffer solution can be calculated from its composition and from the acid
dissociation constant:
pH = pKa –log10
[acid]
[salt]
Indicators
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Indicators are weak acids for which the dissociation can be represented as
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HIn(aq)
+
H2O(l)
H3O-(aq)
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+
In-(aq)
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The acid dissociation constant is represented as KIn and is given by the following
expression:
[H 3O  ] [In  ]
K In 
[HIn]
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In aqueous solution the colour of the acid is distinctly different from that of its conjugate
base.
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The colour of the indicator is determined by the ratio of HIn to In-.
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The theoretical point at which colour changes occurs when H+ = KIn
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The colour change is assumed to be distinguishable when HIn and In- differ by a factor
of 10.
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The pH range over which a colour change occurs can be estimated by the expression
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pH = pKIn  1
Key area: Reaction feasibility
Standard enthalpy of formation
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The standard enthalpy of formation of a substance (Hof) is the enthalpy change when one
mole of a substance is formed from its elements in their standard states.
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The standard enthalpy of a reaction can be calculated from tabulated standard molar
enthalpies of formation using the relationship
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Hof
=
Hof(products) - Hof(reactants)
Knowledge of standard conditions is important.
Entropy
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The entropy (S) of a system is the degree of disorder of the system. The greater the
disorder, the greater the entropy.
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Entropy increases as temperature increases.
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Changes of state involve changes in entropy. Melting and evaporation are accompanied
by increases in entropy. One version of the Third Law of Thermodynamics states that the
entropy of a perfect crystal at 0 K is zero.
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The standard entropy of a substance is the entropy value for the standard state of the
substance.
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The change in standard entropy for a reaction system can be calculated from the standard
entropy values for the reactants and products:
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So
=
So(products) - So(reactants)
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Free energy
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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.
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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.
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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.
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The total entropy change is normally expressed in terms of Gibbs free energy (G). The
direction of spontaneous change is in the direction of decreasing free energy. The change
in standard free energy for a reaction is related to the standard enthalpy and entropy
changes by:
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Go = Ho - T So
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The standard free energy change of a reaction can be calculated from the standard
enthalpy and standard entropy changes for the reaction.
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The standard free energy change of a reaction can be calculated from the standard free
energies of formation of the reactants and products using the relationship:
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Go
=
Go(products) - Go(reactants)
Applications of the concept of free energy
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The feasibility of a chemical reaction under standard conditions can be predicted from the
calculated value of the change in standard free energy (Go).
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The temperature at which the reaction becomes feasible can be calculated for a reaction
for which both Ho and So have positive values.
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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. Under non-standard conditions any reaction is
feasible if G is negative.
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At equilibrium G = 0.
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A reaction will proceed spontaneously in the forward direction until the composition is
reached where G = 0
Key area: Kinetics
Determining the order of a reaction
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The rate of a chemical reaction normally depends on the concentrations of the reactants.
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For a first order reaction the rate of reaction is proportional to the concentration of one
reactant and the rate can be expressed as:
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rate = k A
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Where k is the rate constant and A is the concentration of reactant A in mol l-1.
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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.
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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.
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The order of a reaction can only be determined from experimental data.
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In general for a reaction of type
nA +
mB
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products
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Where the rate expression is of the form:
rate = k An Bm
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The order of reaction is n with respect to A and m with respect to B and the overall
order is n+ m.
The rate constant can be determined from initial rate data for a series of reactions in which
the initial concentrations of reactants are varied.
Reaction mechanisms
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Reaction mechanisms usually occur by a series of steps.
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The rate of reaction is dependent on the slowest step which is called the ‘rate determining
step’.
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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.
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Unit 2 Organic Chemistry and Instrumental Analysis
Key area: Molecular orbitals
Formation of bonding molecular orbitals
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When atoms approach each other, their separate sets of atomic orbitals merge to form a
single set of molecular orbitals.
Some of the molecular orbitals, known as ‘bonding molecular orbitals’, occupy the region
between two nuclei.
The attraction of positive nuclei to negative electrons occupying bonding molecular
orbitals is the basis of bonding between atoms. Each molecular orbital can accommodate
a maximum of two electrons.
Hybridisation
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Hybridisation is the process of mixing atomic orbitals within an atom to generate a set of
new atomic orbitals called hybrid orbitals.
Bonding in alkanes can be described in terms of sp3 hybridisation and sigma bonds.
Bonding in alkenes can be described in terms of sp2 hybridisation and both sigma and pi
bonds.
Bonding in alkynes can be described in terms of sp hybridisation with both sigma and pi
bonds.
A sigma bond () is a covalent bond formed by end-on overlap of two atomic orbitals lying
along the axis of the plane.
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.
The bonding continuum
Van der Waal’s forces of attractions describe all intermolecular forces of attraction
including London dispersion forces, permanent dipole to permanent dipole and hydrogen
bonds.
London dispersion forces are weak intermolecular forces of attraction that arise due to a
temporary uneven distribution of electrons within atoms and molecules.
In a non-polar covalent bond, the bonding molecular orbital is symmetrical about the
midpoint between two atoms.
Polar covalent bonds result from bonding molecular orbitals which are asymmetric about
the midpoint between two atoms.
Ionic compounds represent an extreme case of asymmetry with the bonding molecular
orbitals being almost entirely located around just one atom.
Absorption of visible light by organic molecules
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Most organic molecules appear colourless because the energy difference between the
highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital
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(LUMO) is relatively large resulting in the absorption of light in the ultraviolet region of the
spectrum.
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Coloured organic compounds contain delocalised electrons within molecular orbitals which
extend across several atoms. This is known as a conjugated system.
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A conjugated system exists when a molecule has alternating single and double bonds.
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The more atoms the delocalised molecular orbital spans, the smaller the energy gap
between the delocalised orbital and the next unoccupied orbital and hence the lower the
frequency of light (or longer the wavelength or lower the energy of radiation) absorbed by
the compound.
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When the wavelength of the light absorbed is in the visible region, the organic substance
will appear coloured. Molecules in which the structural formula contains alternate double
bonds will exhibit molecular orbitals containing delocalised electrons which will extend the
conjugated section of the molecule.
Chromophores
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The chromophore is the group of atoms within a molecule which is responsible for the
absorption of light in the visible region of the spectrum (i.e. responsible for its colour).
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This is usually a reasonably large part of the molecule - usually involving a conjugated
system.
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Light can be absorbed when electrons in a chromophore are promoted from one
molecular orbital to another. If the chromophore absorbs light of one colour, the compound
will exhibit the complementary colour.
Key area: Molecular structure
Structural formulae
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Structural formulae and skeletal formulae can be drawn for organic molecules with no
more than 10 carbon atoms in their longest chain.
Conversion between molecular, structural and skeletal organic molecules can be
demonstrated.
In a skeletal formula neither the carbon atoms, nor any hydrogens attached to the carbon
atoms, are shown. The presence of a carbon atom is implied by a ‘kink’ in the carbon
backbone, and at the end of a line.
Key area: Stereochemistry
Stereoisomerism
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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 nonsuperimposable.
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Geometric isomerism
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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.
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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. Geometric
isomers display differences in some physical properties.
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Geometric isomerism can also influence chemical properties, for example cis-but-2enedioic acid is more readily dehydrated than trans-but-2-enedioic acid.
Optical isomerism
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Optical isomers are non-superimposable mirror images of asymmetric molecules and are
said to be chiral molecules or enantiomers.
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Optical isomerism can occur in substances in which four different groups are arranged
around a carbon atom.
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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.
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Racemic mixtures contain equal amounts of both optical isomers and are optically
inactive.
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In biological systems only one optical isomer of each organic compound is usually
present.
Key area: Synthesis
Reaction types
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Equations can be written for the following reaction types and, given equations, these
reaction types can be identified.
substitution
addition
elimination
condensation
hydrolysis
oxidation
reduction
dehydration
Synthetic routes
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Synthetic routes, with no more than three steps, can be devised from a reactant to a
product.
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Molecular structures can be examined and possible reactions can be determined.
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Bond fission
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When an organic reaction takes place, bonds are broken and formed. If, when the bond
between atoms breaks, each atom retains one electron from the former covalent bond,
then two free radicals are formed. This is known as homolytic fission.
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Reactions involving free radicals tend to result in formation of very complex mixtures of
products, thus making them unsuitable for synthesis.
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If, when the bond between atoms breaks, one atom retains both of the electrons from the
former covalent bond, then an ion pair is formed. This is known as heterolytic fission.
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Reactions proceeding via heterolytic fission tend to produce far fewer products and are
therefore better suited for synthesis.
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Heterolytic fission will be favoured when the bond between the atoms is polar.
Electrophiles and nucleophiles
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In reactions involving heterolytic bond fission, attacking groups are classified as
'nucleophiles' or 'electrophiles'.
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Nucleophiles are atoms or groups of atoms which are attracted towards atoms bearing a
(partial) positive charge. Nucleophiles are capable of donating and sharing an electron
pair to form a new bond.
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Electrophiles are atoms or groups of atoms which are attracted towards atoms bearing a
(partial) negative charge. Electrophiles are capable of accepting an electron pair.
Curly arrow notation
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Double-headed and single headed curly arrows are used to show the movement of
electron during a reaction.
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The base of the arrow shows the original location of the pair of electrons. The head of the
arrow indicates the destination of the pair of electrons.
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An arrow starting at the middle of a covalent bond indicates that heterolytic bond fission is
occurring.

When an arrow is drawn with the head pointing to the space between two atoms, this
indicates that a covalent bond will be formed between the two atoms.

A double-headed arrow indicates the movement of an electron pair and a single-headed
arrow indicates the movement of a single electron.
Halogenoalkanes

Halogenoalkanes are named according to IUPAC rules.

Monohalogenoalkanes can be classified as primary, secondary or tertiary.

Monohalogenoalkanes undergo nucleophilic substitution reactions.

Monohalogenoalkanes undergo elimination reactions to form alkenes.
14

Monohalogenoalkanes react with:
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.

Monohaloalkanes can also undergo elimation reactions to form alkenes.
Alkenes

Alkenes can be prepared in the laboratory by
i.
dehydration of alcohols using aluminium oxide, conc. sulfuric acid or
orthophosphoric acid.
ii.
base-induced elimination of hydrogen halides from monohalogenoalkanes.

Alkenes undergo
i.
catalytic addition with hydrogen to form alkanes
ii.
addition with halogens to form dihalogenoalkanes
iii.
addition with hydrogen halides according to Markownikoff’s rule to form
monohalogenoalkanes
iv.
acid-catalysed addition with water according to Markownikoff’s rule to form
alcohols.

The mechanisms of the above electrophilic addition reactions involve
i.
ii.
iii.
for halogenation, cyclic ion intermediate
for hydrohalogenation, carbocation intermediate
for acid catalysed hydration, carbocation intermediate.
All of the above mechanisms can be represented using curly arrows.
Reaction mechanisms

The following reaction mechanisms can be described in terms of electron shifts.

radical substitution of alkanes.

electrophilic addition to alkenes

carbocation mechanism

cyclic ion intermediate mechanism.

nucleophilic substitution: SN1 and SN2

The dominance of an SN1 or SN2 mechanism for a particular haloalkane can be
explained in terms of steric hindrance and the inductive stabilisation of an
intermediate carbocation.

An SN2 reaction proceeds via a single five-centred transition state, whereas an SN1
reaction occurs in two steps via a carbocation.
15
Physical Properties

The following physical properties are explained in terms of the intermolecular forces
involved
melting and boiling points
miscibility with water
polarity
Alcohols

Alcohols exhibit hydrogen bonding and as a result have higher boiling points than other
organic compounds of comparable relative formula mass and shape.

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.
ii.
halogenoalkanes by substitution.

In industry, alcohols (except methanol) can be manufactured by the acid-catalysed
hydration of alkenes.

Alcohols react with some reactive metals to form alkoxides.

Alcohols can be dehydrated to alkenes.

Alcohols undergo condensation reactions slowly with carboxylic acids and more vigorously
with acid chlorides to form esters.

Alcohols undergo oxidation with oxidizing agents.
i
Primary alcohols are oxidized to aldehydes and then to carboxylic acids.
ii
Secondary alcohols are oxidized to ketones.
iii
Tertiary alcohols are not readily oxidized.
Ethers

Ethers have the general formula R1-O-R2 where R1 and R2 are alkyl groups.

Ethers are named according to IUPAC rules.

Due to the lack of hydrogen bonding, ethers have lower boiling points than the
corresponding isomeric alcohols.

Ether molecules can hydrogen bond with water molecules thus explaining the solubility in
water of some ethers of low relative formula mass. The solubility decreases as the
molecular size increases.

Ethers are highly flammable and on exposure to air may form explosive peroxides.

Ethers can be prepared by the reaction of halogenoalkanes with alkoxides.

Ethers are useful solvents for many organic compounds due to relative stability with
respect to oxidation and reduction.
16
Aldehydes, ketones and carboxylic acids

The following physical properties can be explained in terms of dipole-dipole attractions
and/or hydrogen bonding

i.
ii.
iii.

Tollens’ reagent or Benedict’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), respectively.

Aldehydes and ketones can be reduced to primary and secondary alcohols, respectively,
by reaction with lithium aluminium hydride in ether.

The reaction with 2,4-dinitrophenylhydrazine to form a 2,4-dinitrophenylhydrazone is an
example of a condensation reaction.

The melting points of the resulting 2,4-dinitrophenylhydrazones are used to identify
carbonyl compounds.

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.

In pure carboxylic acids hydrogen bonding produces dimers thus explaining the relatively
high boiling points. Dimerisation does not occur in aqueous solution.

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.

Carboxylic acids are weak acids. Their slight dissociation in water can be explained by the
stability of the carboxylate ion caused by electron delocalisation.

Carboxylic acids can be prepared by:

i.
ii.

Reactions of carboxylic acids include:

i.
ii.
iii.
iv
higher boiling points than corresponding alkanes.
lower boiling points than corresponding alcohols.
miscibility of lower members with water.
oxidising primary alcohols and aldehydes.
hydrolysing nitriles, esters or amides.
formation of salts by reactions with metals, carbonates and alkalis.
condensation reactions with alcohols to form esters.
reaction with ammonia or amines and subsequent heating of the ammonium salt to
form amides.
reduction with lithium aluminium hydride to form primary alcohols.
Amines

Amines are named according to IUPAC rules.

Amines are organic derivatives of ammonia and can be classified as primary, secondary
or tertiary.
17

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.

Amine molecules can hydrogen bond with water molecules thus explaining the
appreciable solubility of the lower amines in water.

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.

Amines react with aqueous mineral or carboxylic acids to form salts.
Aromatics

Bonding in benzene can be described in terms of sp2 hybridisation, sigma and pi bonds
and electron delocalisation.

Benzene C6H6 is the simplest aromatic hydrocarbon and its unexpected stability can be
attributed to the presence of delocalised electrons.

A benzene ring in which one hydrogen atom has been substituted by another group is
known as the phenyl group and has the formula -C6H5.

Most reactions of benzene involve attack of an electrophile on the cloud of delocalised
electrons. This is called electrophilic substitution.

Benzene resists addition reactions but undergoes electrophilic substitution reactions.
These include:
i.
halogenation e.g. chlorination and bromination to produce chlorobenzene and
bromobenzene
ii.
nitration to produce nitrobenzene
iii.
sulfonation to produce benzene sulfonic acid
iv.
alkylation to produce alkylbenzenes.
Key area: Experimental determination of structure
In organic chemistry a number of experimental techniques are carried out to verify that the
correct chemical structure has been synthesised.
Elemental microanalysis

Elemental microanalysis (combustion analysis) can be used to determine the masses of
C, H, O, S and N in a sample of an organic compound in order to find the empirical
formula.

Other elements in the organic compounds have to be determined by other methods
18
Mass spectrometry

Mass spectrometry can be used to determine the accurate formula mass and structural
features of an organic compound.

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.




Organic compounds can be identified from the very accurate determination of the relative
molecular masses of the parent ion and the ion fragments.
Infra red spectroscopy

Infra-red spectroscopy can be used to identify certain functional groups in an organic
compound.

The vibrational energy of a chemical bond is quantised.

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.

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.

Wavelength is expressed in spectra as wavenumber.

The unit of measurement of wavenumber is the reciprocal of length in centimetres (cm-1).
Proton Nuclear magnetic resonance spectroscopy (1 H NMR)

Proton nuclear magnetic resonance spectroscopy (proton 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

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).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.

In the NMR spectrum the peak position is related to the environment of the proton.

The area under the peak is related to the number of protons in that environment.
19

An interaction with H atoms on a neighbouring carbon atom can result in the splitting of
NMR peaks into 'multiplets'. The number of H atoms on the neighbouring carbon will
determine the number of lines within a multiplet.

Hydrogen nuclei behave like tiny magnets and in a strong magnetic field some are aligned
with the field (lower energy) whilst the rest are aligned against it (higher energy).
Absorption of radiation in the radio-frequency region of the electromagnetic spectrum will
cause the hydrogen nuclei to ‘flip’ from the lower to the higher energy alignment. As they
fall back from the higher to the lower level, the emitted radiation is detected.

The standard reference substance used in NMR spectroscopy is tetramethylsilane (TMS)
which is assigned a chemical shift value equal to zero.

Low resolution proton NMR spectra can be drawn and analysed.

High resolution proton NMR spectra can be analysed.
Key area: Pharmaceutical chemistry
Effect of drugs on the body

Drugs are substances which alter the biochemical processes in the body and those which
have a beneficial effect are called medicines.

A medicine usually contains the drug plus other ingredients.
Classification of drugs

Many medicines can be classified as agonists or as antagonists according to whether they
enhance or block the body’s natural responses.

An agonist will produce a response like the body’s natural active compound.
An antagonist produces no response but prevents the action of the body’s natural active
compound.
How drugs work

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).

The structural fragment of a drug molecule which confers pharmacological activity upon it
normally consists of different functional groups correctly orientated with respect to each
other.

The overall shape and size of the drug has to be such that it fits a binding site. The
functional groups on both the drugs and the receptor are positioned such that the drugs
can interact with and bind to the receptor.

The types of interaction between drugs and binding sites can be identified.

By comparing the structures of drugs that have similar effects on the body, the structural
fragment that is involved in the drug action can be identified.
20

Many drugs act as enzyme inhibitors by binding to the enzyme’s active site and blocking
the reaction normally catalysed there.

% solution is the mass of solute made up to 100 cm3 of solution.

% by volume is the number of cm3 of solute made up to 100cm3 of solution.

The unit ppm stands for parts per million and refers to 1 mg per kg.
21
Unit 3
Researching Chemistry
Overarching principles
Throughout each of the key areas the learners must be able to apply the following
principles:

Precision

Accuracy

Uncertainties

Units
Key area: Stoichiometric calculations

Stoichiometry is the study of quantitative relationships involved in chemical reactions. The
ability to balance and interpret equations enabling calculations to be carried out involving
any of the above skills/techniques is an important part of chemistry at this level and is
examinable in the course assessment.

A quantitative reaction is one in which the substances react completely according to the
mole ratios given by the balanced (stoichiometric) equation.
Experimental determination of percentage yield

Theoretical percentage yields can be calculated and compared with actual yield . Mass
transfer or mechanical losses, purification of product, side reactions, equilibrium position
and purity of reactants, inevitably lowers the percentage yield.
Key area: Gravimetric analysis


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.
Key techniques to perform a gravimetric analysis include ‘weighing by difference’ and
‘heating to constant mass’.
Key area: Volumetric analysis

Volumetric analysis involves using a solution of accurately known concentration in a
quantitative reaction to determine the concentration of another substance.
22

A solution of accurately known concentration is known as a standard solution.

A standard solution can be prepared directly from a primary standard.

A primary standard must have, at least, the following characteristics:

high state of purity

stability in air and in solution

solubility

reasonably high formula mass.

The volume of reactant solution required to just complete the reaction is determined by
titration.

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 end-point.

Acid/base titrations are based on neutralisation reactions.
Controls

Use of a control validates a technique and may consist of carrying out a determination on
a solution of known concentration.
Complexometric titrations

Complexometric titrations are based on complex formation reactions. EDTA is an
important complexometric reagent and can be used to determine the concentration of
metal ions such as nickel(II).

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.
Back titrations

Back titration is used to find the number of moles of a substance by reacting it with an
excess volume of reactant of known concentration. The resulting mixture is then titrated to
work out the number of moles of the reactant in excess. From the initial number of moles
of that reactant the number of moles used in the reaction can be determined, making it
possible to work back to calculate the initial number of moles of the substance under test.

A back titration is useful when trying to work out the quantity of substance in an insoluble
solid.
Key area: Practical skills and techniques
Standard solutions and calibration curves

Preparation of standard solutions using accurate an accurate dilution technique.
23

Colorimetry uses the relationship between colour intensity of a solution and the
concentration of the coloured species present.

Calibration curves can be constructed to determine an unknown using solution of
appropriate concentration.

The concentration of the ‘unknown’ solution is determined from its absorbance and by
referring to the calibration curve. The straight line section of the calibration graph should
cover the dilution range likely to be used in the determination.
Preparation and purification of an experimental product

Distillation, reflux, vacuum, filtration, recrystallization and separating funnels are practical
skills that can be used in the preparation and purification of an experimental product.








Distillation: The boiling point of a compound — determined by distillation — is welldefined and thus is one of the physical properties of a compound by which it is
identified.
Refluxing: Used to apply heat energy to a chemical reaction mixture over an
extended period of time. The liquid reaction mixture is placed in a round-bottomed
flask along with anti-bumping granules with a condenser connected at the top. The
flask is heated vigorously over the course of the chemical reaction; any vapours
given off are immediately returned to the reaction vessel as liquids when they
reach the condenser.
Vacuum filtration: Using a Buchner, Hirsch or sintered glass funnel. These
methods are carried out under reduced pressure and provide a faster means of
separating the precipitate from the filtrate. The choice of filtering medium depends
on the quantity and nature of the precipitate.
Recrystallisation: Lab technique used to purify solids, based upon solubility. The
solvent for recrystallisation must be carefully selected such that the impure
compound is insoluble at lower temperatures, yet completely soluble at higher
temperatures. The impure compound is dissolved gently in the minimum volume of
hot solvent then filtered to remove insoluble impurities. The filtrate is allowed to
cool slowly to force crystallisation. The more soluble impurities are left behind in
the solvent.
Use of a separating funnel: Solvent extraction can be an application of the partition
of a solute between two liquids. It is based on the relative solubility of a compound
in two different immiscible liquids, usually water and an organic solvent. The
equilibrium constant is called the partition coefficient. The two solvents form two
separate layers in the separating funnel and the lower layer is run off into one
container and the upper layer is poured out into another container. The quantity of
solute extracted depends on the partition coefficient and on the number of times
that the process is repeated.
Thin-layer chromatography
Thin –layer chromatography (TLC) uses a fine film of silica or aluminium oxide spread
over glass or plastic to assess the purity of experimental products.
Rf values can be calculated and under similar conditions a compound will always have the
same Rf value within experimental error.
Since a pure substance will show up as only one spot on the developed chromatogram,
TLC can be used to assess the purity of a product prepared in the lab.
24




Chromatographic separations depend on the partition equilibrium between two phases;
one stationary and the other mobile.
Other examples of chromatography are:
 paper chromatography
 gas liquid chromatography
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.
Determination of melting point and mixed melting point

The melting point of an organic compound is one of several physical properties by which it
can be identified. A crystalline substance has a sharp melting point falling within a very
small temperature range. Determination of the melting point can also give an indication of
the purity of an organic compound, as the presence of impurities lowers the melting point
and extends its melting temperature range.

Since impurities lower the melting point, the technique of mixed melting point
determination can be used as a means of identifying the product of a reaction
25