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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
AQA A2-LEVEL
Student Guide to A2 Unit 5
Energetics, Redox and Inorganic Chemistry
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LUND 7 May 2017
1
A2 Chemistry
Unit 5
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Enthalpy Changes and Hess’s Law



enthalpy changeH = heat exchange at constant pressure (open container)
enthalpy change cannot be measured directly but can be determined experimentally
H varies with temperature and pressure so Standard Conditions (Ho298) are required:
temperature
pressure
physical state


298 K (25ºC i.e. nominal room temperature)
100kPa don’t put ‘1 atmosphere’ in the exam!!!
at room temperature and most stable allotrope (e.g. graphite)
the enthalpy of formation of an element is by definition zero
here are two enthalpy changes that you learned in AS Chemistry:
the enthalpy of
is the enthalpy change that occurs when one mole of
at 298K and 100 kPa (i.e. standard conditions)
formation (Hof,298)
a compound is formed from its elements
combustion (oc,298)
an element or compound is completely combusted in xs oxygen
Hess’s Law: the enthalpy change for a reaction is dependent only on the initial and final
states of the system and is independent of the route taken.
Reactants
Products
1
2
Hess’s Law states that 1 = 2
A MARK IN THE EXAM !!

Hor,298 = Hof,298 (Products) - Hof,298 (Reactants)
Elements
Reactants
s
Products
1
2
Hess’s Law states that 1 = 2
Hor,298 = Hoc,298 (Reactants) - Hoc,298 (Products)
Combustion
Products
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Enthalpy Definitions

note that once again in all these definitions you specify both the quantity and type of the
particles initially and finally
the enthalpy of
is the enthalpy change that occurs when one mole of
at 298K and 100 kPa (i.e. standard conditions)
atomisation

Hoat
note the differences between the enthalpies of formation, atomisation and bond dissociation
as these are often used incorrectly in calculations (think about the values for Iodine)
1st ionisation energy Hoi(1)


gaseous singly charged cations are formed from gaseous atoms
revise the trends (periods and groups) of first ionisation energies in addition to the changes in
magnitude (shells) and equations for successive ionisation energies of the same element
1st electron affinity Hoea(1)


gaseous atoms are formed from an element in its standard state
gaseous singly charged anions are formed from gaseous atoms
you should be able to write an equation for the gain of an electron by any given atom/ion
the first electron affinity is exothermic as the vacancy in the subshell is subject to an
attractive force from the nucleus that is not completely screened.
the second electron affinity is endothermic since an electron is being moved towards a
negative ion and mutual repulsion must be overcome (subsequent additions are increasingly
endothermic as the size of the negative charge increases)
EXOTHERMIC
lattice enthalpy HoL
an ionic lattice is formed from its gaseous ions
lattice dissociation enthalpy
ENDOTHERMIC
an ionic compound separates into gaseous ions

Note that these
have the same
numerical
value but
opposite signs
lattice enthalpy is higher if the charges are higher and/or the ions smaller as the electrostatic
force of attraction is greater – see 169 (note it’s the same as lattice energy for A-level use)
Lattice enthalpy and enthalpy of formation are not the same thing:
Lattice enthalpy starts from gaseous ions.
Enthalpy of formation starts from elements in their normal states.

lattice energy cannot be determined experimentally as it is impossible to react an exact
amount of gaseous anions and cations and measure the energy released.
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry

the value of the lattice energy can be derived theoretically using an electrostatic relationship
between charge sizes and distances but this makes two incorrect assumptions:
1
2



experimental values are higher than theoretically calculated values for HoL as there is in
reality a degree of covalency in all ionic compounds as the electrostatic attraction of the
cations will distort the valence electron cloud of the anions
in most cases (small ions that are singly charged) there is a reasonable agreement between
experimental and theoretical data (within 5%) i.e. the ionic model is adequate.
however, the deviation is more significant i.e. the degree of covalency increased if:
1
2
3



the ions are point sources
the ions do not mutually interact such as to distort their shape i.e that they are purely ionic
and exhibit no degree of covalency
the cations are smaller
the anions are larger
the charge on the ions is larger
a small cation with a charge of 2+ or higher will present a highly
polarising field to an anion the effect of which will be more
significant if that anion is relatively large
the anionic electron clouds will be distorted (polarised) towards the
polarising field of the cation (this is not dissimilar to the tenuous
outer atmosphere of a red giant star becoming distorted in
association with a strong gravitational field)
this can lead to an overlap of electron clouds i.e. the compound is in
effect covalent
Summary Questions
A2 Chemistry (Nelson Thornes) AQA
Chemguide
Page 166
Page 174
1–3
2
163 – 166, 169, 174
Lattice enthalpy, ionic structures
s-cool: Chemical Energetics
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4
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Born-Haber Cycles

a Born-Haber Cycle for a simple ionic solid can be constructed using Hess’s Law
(including ionisation energy, electron affinity, lattice energy, atomisation energy)
Note that, unlike the energy cycle diagrams met at AS-level, these are energy level diagrams in
that they are drawn with up = endothermic and down = exothermic change





elements are referenced to zero
NaCl and MgCl2 are shown in your textbook
for oxides/sulphides the second electron affinity will be endothermic (i.e. back up)
but why not MgCl without a second (endothermic) ionisation energy or MgCl3 where a 3+ ion
is forming the lattice giving a much higher (exothermic) lattice energy (see page 170)
ionisation energy and lattice energy are the two major contributory values
Enthalpies of Solution (Hosol)
the enthalpy of
is the enthalpy change that occurs when one mole of
at 298K and 100 kPa (i.e. standard conditions)
solution
Hosol
solute is completely dissolved in water
hydration
Hohyd
isolated gaseous ions are hydrated





hydration involves ion-dipole electrostatic forces hence ionic solids can dissolve in polar
solvents such as water
the payback of hydration enthalpy is not available in non-polar solvents rendering the process
too endothermic overall to be possible
solubility is essentially a trade off between overcoming the enthalpy of lattice dissociation
enthalpy (endothermic in this direction) and enthalpies of hydration (always exothermic)
solubility is feasible if this is an exothermic outcome overall (e.g. NaOH(aq))
although ENTROPY (see later) will also have a say if the enthalpy of solution is only slightly
endothermic (e.g. NH4NO3(aq))
Hess’s Law can be used to construct a solvation energy cycle (see 173)
Summary Questions
Page 172
1
Page 174
1
Pages 171 – 2 and Questions 1 - 4
Page 184
1
Page 248
5, 6 (read ‘entropy’ for d(i))
How Science Works
Exam Style Questions
A2 Chemistry (Nelson Thornes) AQA
Chemguide
164 – 173
Born Haber lattice
s-cool: Chemical Energetics
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Free Energy and Entropy








it seems reasonable that exothermic reactions can occur spontaneously, after all the products
are relatively more stable than the reactants (stronger bonds)
obviously the rate of the process will depend upon the kinetic barrier (EA) – i.e. kinetic
feasibility and they could be slow e.g. iron rusting or extremely slow e.g diamond into
graphite
spontaneously thus means it will happen under a given set of conditions but not necessarily in
an instant and it could be almost infinitely slow
a small proportion of spontaneous reactions however are endothermic e.g. solvating
NH4NO3(s)
at first sight this seems counter intuitive as endothermic reactions require a transfer of heat
from their surrounding resulting in products that are relatively less stable (higher stored
chemical energy) in terms of enthalpy
so why doesn’t NH4NO3(aq) spontaneously undissolve as this would be exothermic and also be
kinetically more feasible relative to the forward reaction (lower EA)
why does it only go one way (assuming the solubility limit has not been exceeded)
clearly there must be another factor in the viability of a chemical (and physical) process that
determines whether it will occur spontaneously or not
Entropy (S)


chemical and physical change are governed by the laws of probability
the degree of disorder (ways of distributing energy) within a system is called Entropy

the greater the degree of disorder in a system the greater the entropy e.g.
state changes
solvation
diffusion
complexity of molecules – increases way to distribute energy hence greater entropy
temperature – entropy is increased if it is raised and visa-versa

the second law of thermodynamics states that all viable chemical and physical changes result
in an increase in the TOTAL entropy in the Universe

hence all spontaneous reactions result in an increase in TOTAL entropy irrespective of
whether they are exothermic or endothermic

an endothermic reaction can thus be spontaneous IF it results in an increased degree of
randomness (of distribution of energy) in the universe
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
ΔS






heat energy flowing into a system will increase its entropy – more available energy levels –
more available ways in which energy can be distributed between the particles
for an endothermic reaction lowering the temperature reduces entropy but the change from
reactants to products could increase entropy (e.g. solvation).
IF there is a NET increase in entropy the reaction may be spontaneous
like enthalpy, So is specified for 298K and 100 kPa
unlike H it is possible to quantify S (units are JK-1mol-1) since it can be extrapolated from a
value of zero at 0K
Look
!
it is possible to determine the entropy change in the system (the chemicals reacting)
ΔS(system)
=
∑S(products)
-
∑S(reactants)
note that the overall entropy may only change slightly if all the chemicals are in the same phase



if entropy must increase – how can water condense spontaneously given that it achieves a
more ordered state i.e. a decrease in entropy?
We must also consider the entropy change in the surroundings to complete the picture –
remember it is the TOTAL entropy change that concerns us
the total entropy change is expressed by:
ΔSototal = ΔSosystem + ΔSosurroundings


it is very difficult to calculate the entropy change in the surroundings
however, it can be determined from the enthalpy transferred from/to the system at a given
temperature (since temperature effects the value of entropy)
T ΔS o surroundings
=
-ΔHosystem

energy is transferred from the steam (system) to the surface that it condenses on
(surroundings) increasing its entropy such that the overall entropy increases

the ligand displacement reactions that you will meet in transition metal chemistry can often be
attributable to favourable entropy changes where the total number of particles increases
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Gibbs free energy (G)



entropy and enthalpy can be combined in a quantity called Gibbs free energy (G)
G o is the standard free energy change (kjmol-1) and quantifies the enthalpy and entropy
changes at a given temperature
for a reaction to be feasible it must be zero or negative i.e such that ΔSototal is positive
ΔGo = - TΔSototal

some mathematical magic with the previous equations leads us to:
ΔGo = ΔHosystem - TΔSosystem
kjmol-1
Remember to
divide entropy
by a 1000
since its in J
not kJ !!!


G o effectively determines the feasibility of a reaction
reaction feasibility is thus determined by the relative magnitude of the enthalpy and entropy
changess at a given temperature


temperature has a significant effect on reaction feasibility
a reaction first becomes feasible when the temperature has been raised to a value where
G o = 0
this is particularly relevant to metal extraction where heating may be necessary to make the
reaction feasible not just to increase the rate



a feasible reaction is not necessarily spontaneous as there may be a high activation
energy involved
when G o = 0 an equilibrium will exist in a closed system

ΔS
Outcome
-
+
G always negative
Feasible at ALL temperatures
+
-
G always positive
NOT feasible at ANY
temperature
-
-
IF  > T ΔS then reaction
is feasible i.e. at LOWER
temperatures
+
+
IF T ΔS >  then reaction
is feasible i.e. at HIGHER
temperatures
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Change of State

a system is at equilibrium (e.g. ice/water at 0oC) when:
G = 0
 = T ΔS




if heat goes into the system then the entropy term must increase for G to remain zero
i.e. it melts to a more disordered state
visa-versa is true
the entropy change for evaporation is significantly higher than for melting as the increase in
randomness is more significant
Summary Questions
Exam Style Questions
A2 Chemistry (Nelson Thornes) AQA
Chemguide
Page 183
Page 185
Page 246
178 – 183
Entropy
LUND 7 May 2017
9
1-3
2–4
1, 2
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Periodicity
Elements of Period 3
Structure
Electrical
Conduction
Na
Mg
giant metallic
Al
good, delocalised (mobile) eincreases Na  Al as more
e- delocalised
Si
Giant
Covalent
P
S
Cl
Simple Molecular
Ar
Atoms
poor, electrons localised in covalent
bonds so not free to move through
structure
poor
Melting Point high, strong metallic
bonding, strong forces of
attraction throughout entire
structure increases Na  Al
as more e- delocalised
higher,
strong
covalent
bonding
between
atoms
low, strong covalent bonds
between atoms, but weak
intermolecular forces
between Molecules
very
low
Mpt/oC
1410
44
-189
98







651
660
113
-101
silicon has a similar structure to diamond – each silicon has 4 covalent bonds
silicon is important in the electronics industry
sodium is used in streetlights and as a coolant in nuclear reactors
magnesium and aluminium are important structural materials often used in alloys
chlorine is a product of the salt-alkali industry and has numerous uses
argon is used in lighting and lasers
sulphur is a raw material for the manufacture of sulphuric acid
Reaction of the elements of Period 3 with water
Na
vigorous reaction – darts
about the surface – dissolves
to form a very alkaline
solution
Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Mg
very slow with cold water, but
more readily with steam
Mg(s) + H2O(l)  MgO(s) + H2(g)
Cl
Unit 2
revision
only
chlorine reacts with water to produce a 0
-1
+1
mixture of two acids by disproportionation Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq)
(moist blue litmus is first turned red by
chlorine gas then bleached)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Reactions of the Elements of Period 3 with Oxygen



reactivity (generally) decreases across the period
argon does not react at all
chlorine-oxygen compounds do exist but are not formed by direct combination



the metals become coated in an oxide layer – hence appear dull
sodium must be stored under oil due to its reactivity
aluminium oxide forms a protective layer around aluminium protecting it from further
corrosion

red phosphorus needs heating to react with oxygen but its allotrope, white phosphorus, is
stored under water as when dry it burns spontaneously in oxygen

the oxide formed contains the element in its highest oxidation state except for sulphur
Na
burns very vigorously with a
yellow
flame
Mg
burns very vigorously with a white 2Mg(s) + O2(g)  2MgO(s)
flame – white MgO ash
Al
initially vigorous
2Al(s) + 3O2(g)  2Al2O3(s)
Si
slow
Si(s) + O2(g)  SiO2(s)
P
vigorous – white fumes of P4O10
P4(s) + 5O2(g)  P4O10(s) phosphorus(V) oxide
S
melts, blue flame – forms SO2 a S (s) + O2(g)  SO2(g)
colourless choking acidic gas
Summary Questions
A2 Chemistry AQA (Nelson Thornes)
Chemguide
4Na(s) + O2(g)  2Na2O(s)
Page 191
1–3
186 - 188
Period 3 elements
LUND 7 May 2017
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silicon(IV) oxide
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Periodicity of Oxides
Period 3
Oxides
Physical State
(298K)
Mpt/oC
Structure
Na2O
Mg
Al2O3
SiO2
2852
2072 – lower
higher – size since
and charge of different
cation
structure
Giant Ionic Lattice
(Al2O3 has a degree of covalency – small ion
+ large charge)
1275
Basic
Water
1610
Giant
Covalent
similar to
diamond
Amphoteric




Simple Molecular
Acidic
Acid
Alkali
Na2O + H2O  2NaOH
oxide ion is hydrolysed
very soluble pH 14
acid + metal oxide  salt +
MgO + H2O  Mg(OH)2
a weakly alkaline solution water
only slightly soluble pH 9
Al2O3 + 3H2O + 2OH 2[Al(OH)4]SiO2 + 2OH-  SiO32(silicate ion) + H2O
P4O10 + 12OH-  4PO43(phosphate(V) ion) + 6H2O
SO2 + 2OH-  SO32(sulphite ion) + H2O
SO3 + 2OH-  SO42(sulphate ion) + H2O
Si
S
580 ?
SO3
(SO2)
liquid
(gas)
17 (-73)
inc. IMF
Al
P
P4O10
Solid
Nature of
Oxide
Reaction
of Oxide
Na
MgO
P4O10 + 6H2O  4H3PO4
phosphoric(V) acid pH 0
SO2 + H2O  H2SO3
sulphuric(IV) acid pH 3
SO3 + H2O  H2SO4
sulphuric(VI) acid pH 0
aluminium oxide acts as an abrasive (as corundum) due to the strong bonding
aluminium/magnesium oxides are used as a refractory material (furnace)
melting point silicon(IV) oxide > carbon dioxide (strong covalent bonding throughout the
structure rather than simple molecular lattice held together by weak intermolecular forces)
you should also be able to write a 3 step equation for phosphoric acid and water
Summary Questions
Exam Style Questions
A2 Chemistry AQA (Nelson Thornes)
Chemguide
Page 193
Page 194
Page 249
1–3
1-6
7
189 - 193
Period 3 oxides
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Oxidation Numbers

oxidation is the LOSS of electrons, reduction is the gain of electrons OILRIG
oxidising agents (oxidants) accept electrons and are themselves reduced
reducing agents (reductants) donate electrons and are themselves oxidised
Oxidation State

this is a ‘book keeping’ method of the effective control of electrons used in bonding
elements
=
0
oxidation state of elements in simple ions
=
charge on ion
 oxidation state of elements in polyatomic ions
=
charge on ion
 oxidation state of elements of a compound
=
0
the relatively more electronegative element is assigned the negative oxidation state
hydrogen
=
+1 (except in metal hydrides where it is -1)
oxygen
=
-2 (except in peroxide O22-) where it = -1)
group 1 metals
=
+1
group 2 metals
=
+2
fluorine
=
-1 (even with oxygen, which is +2 in OF2)
Aluminium
=
+3
metals are always positive in a compound or polyatomic ion


maximum possible oxidation state
=
group number (note not always possible for
various reasons – see later)
oxidation numbers and nomenclature e.g. cobalt(II) nitrate(V), phosphorus(V) oxide
take care not to mix up charge and oxidation numbers in polyatomic species e.g. a
sulphate(IV) ion does NOT have a 4- charge (the IV refers to the oxidation state of the
sulphur)



changes in oxidation numbers can be used to identify redox reactions in inorganic and organic
reactions e.g. metal or halogen displacement reactions
we specifically refer to an element in a species (e.g. ‘the iron in Fe2O3 is reduced’)
assume that multiple instances of an element in a species have the same value (e.g. both
carbons in ethene are -2)
OXIDATION
REDUCTION
oxidation number becomes relatively more positive
oxidation number becomes relatively more negative
A2 Chemistry AQA (Nelson Thornes)
Chemguide
196 - 197
oxidation number
LUND 7 May 2017
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Balancing Redox Equations (Acidic conditions)

protocol for constructing half equations
1
2
3
4
5

try doing these important half equations (either way round):
(i)
(ii)
(iii)
(iv)
(v)

I have written
some answers. Try
to find them and
do the bonus
STAR question
get the species correct (and number e.g. 2Cr3+)
balance the oxygen by adding H2O(l)’s to the side with least O’s
balance the hydrogen’s using H+(aq)’s to the side with least H’s
balance the charge on each side by adding e- to relatively more positive side
add state symbols
Fe3+(aq)/ Fe2+(aq)
MnO4-(aq)/Mn2+(aq)
Cr2O72-(aq)/Cr3+(aq)
S4O62- (aq)/ S2O32-(aq) (tetrathionate/thiosulphate)
CO2(g)/C2O42-(aq) (ethanedioate ion)
HSW: This titration can be used to investigate
the % of copper in a coin – your teacher will
explain how if you ask
now try combining these half equations:
(i)
(ii)
(iii)
(iv)
(v)
iodine and thiosulphate ions
iron (II) ions and manganate(VII) ions
iron(II) ions and dichromate(VI) ions
manganate(VII) ions and ethanedioate ions
dichromate(VI) ions and ethanal
Summary Questions
Exam Style Questions
HSW: This titration can be used to standardise
solutions of potassium dichromate(VI) or
potassium manganate(VIII)
1 – 2 (Q2 should have equilibria arrows)
Page 199
Page
A2 Chemistry AQA (Nelson Thornes)
Chemguide
HSW: These titrations can be used to
investigate the % of iron in an iron tablet or
iron in tea (after some preparatory steps)
197 - 199
redox equation
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Redox equations in alkaline conditions
this is included to allow you to work out rather than rote learn redox reactions for some
organic (e.g. Fehling’s and Tollen’s), transition metal reactions (e.g. chromium) in
alkaline conditions and certain reactions associated with batteries


writing half equations for redox reactions under alkaline conditions is a little more
tricky than under acidic conditions – but here is a useful ‘cheat’
do it exactly as if it was in acidic conditions then cancel out the hydrogen ions by adding
hydroxide ions equally to both sides

you must know and use the formula of the species containing the atom(s) to be
oxidised/reduced as it exists in alkaline conditions

lets consider Fehling’s (see Module 4) in which copper(II) ions are reduced to copper(I)
which exists as copper(I) oxide and an aldehyde is oxidised to a carboxylic acid (why would
you not smell this?)
using the correct species for alkaline conditions but as if in an acidic solution you should
arrive at:

RCHO(aq) + 2Cu2+(aq) + 2H2O(l)




RCOOH(aq) + Cu2O(s) + 4H+(aq)
now the modification
add OH-(aq) to BOTH sides in order to cancel 2H+
subsequently cancel out H2O’s to arrive at:
RCHO(aq) + 2Cu2+(aq) + 4OH-(aq)

RCOOH(aq) + Cu2O(s) + 2H2O(l)

this method also works perfectly well if applied to half equations

try using it to derive the equations met in Periodicity of oxides reacting with alkali from the
known reactant and product
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Electrode potentials and the Electrochemical Series

dynamic equilibrium exists between a metal and its ions in solution which results in a
potential on the metal
Mn+(aq) + ne-




M(s)
this potential depends upon the position of the equilibrium and cannot be measured absolutely
it will be different for a more or less reactive metal so there will be a potential difference,
which is measurable, between the two pieces of metal (try chewing a piece of aluminium if
you have a metal filling – you’ll get the idea)
the relatively more reactive a metal, the relatively more biased the above equilibria will be to
the LHS, the relatively more negative the potential on that metal will be
you should know the practical arrangement of a single cell made up of two half cells (see
figure 3 on page 200)
Zn2+(aq)|Zn(s) and
Cu2+(aq)|Cu(s)
Note: the relatively more oxidised form is shown on the LHS by convention (in practice both
reactions can go either way depending upon what they are combined with)

the pieces of metal are called electrodes
(in this case the cathode is the +ve terminal unlike in electrolysis where it’s the –ve
terminal since the cathode is defined as the electrode where reduction takes place)


the wire allows electron flow between the two electrodes
the salt bridge provides mobile ions to complete the circuit
potassium nitrate (or similar) used in the salt bridge as the ions are unlikely to partake
in electron exchange or precipitation
oxidation takes place at the negative electrode (zinc in this case)
reduction takes place at the positive electrode (copper in this case)

given that the copper is the +ve electrode (by + 1.10V in theory)
electrons flow to it so the reactions that are taking place in the two
half cells are:
Cu2+(aq) + 2e- → Cu(s)

Zn(s) → Zn2+(aq) + 2e-
and
in effect the more reactive metal (better reducing agent) has displaced
the less reactive metal from its solution
A2 Chemistry AQA (Nelson Thornes)
Chemguide
200 - 201
redox equilibria, half cells
LUND 7 May 2017
16
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Standard Hydrogen Electrode




there are a huge number of different half cells that can be constructed (some of which don’t
involve a metal in the redox reaction thus requiring an inert platinum electrode)
to have comparisons between them all would just be silly so we use an electrode potential
equivalent of sea level to which all are then compared
this works exactly like in geography where relative heights of mountains and depths of ocean
trenches are given allowing differences to be calculated
similarly, the standard electrode (reduction) potential Eo for each half cell is measured
relative to the standard hydrogen electrode (SHE), which is defined as 0V under standard
conditions:
25oC
100Kpa
1.0 moldm-3 H+(aq) CARE – must be strong acid, and 0.5 moldm-3 if H2SO4



a diagram of the standard hydrogen electrode (PRIMARY STANDARD) and how it is used
is shown in figure 6 on page 201 which you must be able to draw
a platinum electrode is used as this will not undergo a redox reaction thus will not alter the
potential of the hydrogen electrode
it is covered with finely divided platinum (termed ‘platinized platinum’) to increase surface
area and increase rate
all solutions will be 1.0 moldm-3 and 25oC as other values will alter the relative position
of the equilibria and therefore the absolute and hence measured potential difference
in practice the Calomel electrode A SECONDARY STANDARD is used for convenience
and is initially calibrated against the SHE at a value of +0.27V
Hg2Cl2(aq)


+
-
2e
2Hg(l) +
-
2Cl (aq)
the voltmeter used has a high resistance hence a low current is drawn so a true reading of the
potential difference is made
if a large current is drawn the concentrations of the ions in solution will change thus upsetting
the equilibria and consequentially changing the value being measured
A2 Chemistry AQA (Nelson Thornes)
Chemguide
200 - 1
hydrogen electrode,
non-metal systems, calomel
LUND 7 May 2017
17
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Electrochemical Series



the electrochemical (reactivity) series is related to standard electrode potentials.
standard electrode potentials can be used to predict the relative feasibility of a reaction
by convention the standard electrode potentials are always quoted for the reduction process
the more +ve its value the relatively more likely the forward reaction is going to occur
i.e. the stronger the oxidant (on the left)
the more -ve its value the relatively more likely the reverse reaction is going to occur i.e.
the stronger the reductant (on the right)
It thus follows that…
any species on the LHS can potentially oxidise any species on the RHS with a relatively
more negative electrode potential
any species on the RHS can potentially reduce any species on the LHS with a relatively
more positive electrode potential



thus relatively more reactive metals will have a relatively more –ve electrode potential
lets use a familiar idea to reinforce this – ‘a more reactive metal can displace a less reactive
metal’ e.g. iron nails in copper(II) sulphate solution (not the other way around)
Cu2+(aq) + 2e-
Cu(s)
Eo = + 0.34V
Fe2+(aq) + 2e-
Fe(s)
Eo = - 0.44V
which of the above metals will dissolve in 0.5M sulphuric acid?
NOTES:
1.
2
the number of electrons lost or gained is not a factor in their relative availability
other factors, e.g. kinetics, may prevent a feasible reaction from occurring i.e electrode
potentials do not indicate the rate of a reaction, just its feasibility
(for example where a solid is involved, or where two ions of the same charge are the
reactants)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Representing Cells

convention for representing cells (with the relatively +ve half cell on the RHS):
R  O
O  R
eZn(s)| Zn2+(aq)||Cu2+(aq)|Cu(s)




single line is boundary between phases
commas are used between same phase (e.g. Fe3+(aq),Fe2+(aq))
double line = salt bridge (sometimes shown as a single dashed lines in some sources)
the emf quoted is the right hand side relative to the left hand side
Eocell


=
EoRHS
the hydrogen half cell can be represented in two different ways
life is easier if you put it on the side that provides a positive value for the cell overall as it is
written
Pt(s)| H2(g)|H+(aq)||other half cell

other half cell||H+(aq)|H2(g)|Pt(s)
e.g. for Zn2+(aq)|Zn(s) the electrons flow from the zinc half cell to the hydrogen half cell (Zn is
higher in the reactivity series than hydrogen):
Zn(s)| Zn2+(aq)||H+(aq)|H2(g)|Pt(s)
Eocell =
+0.76V

- EoLHS
zinc is –0.76V
i.e its – ( - 0.76)
EoRHS - EoLHS
0.00V - Zn half cell
the standard electrode potentials of ions of the same element in different oxidation states
can be measured e.g. Eo Fe3+/Fe2+ = + 0.77V
1
2.
3.
What will be used as the electrode?
Show the cell diagram
What concentrations of iron sulphate solutions will you mix together?
Summary Questions
Exam Style Questions
1–2
4
Page 203
Page 213
A2 Chemistry AQA (Nelson Thornes)
Chemguide
203
Electrochemical
LUND 7 May 2017
19
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Electrode Potentials and Reaction Direction

the feasibility of a reaction requires electrons to flow in the right direction – the reverse
reaction is NOT feasible under standard conditions
Reducing agent provides
e- and is itself oxidised
R  O
O  R
Oxidising agent accepts eand is itself reduced
ethe oxidising agent must have a relatively more positive electrode potential
the reducing agent must have a relatively more negative electrode potential

remember that the electrode potentials are quoted for unimolar concentrations under standard
conditions and will be modified through changes in:
concentration

pH
temperature
the effect of changes can be determined by applying LCP to the feasibility of a reaction
e.g. increasing the concentration of one of the reactants involved in the oxidising half will
increase its electrode potential thus feasibility i.e. the forward reaction is more favoured
similarly increasing the concentration of one of the reactants involved in the reducing half
will decrease its electrode potential thus the reverse reaction is favoured
our technician wishes to prepare chlorine gas based on the electrode potentials given below
and you need to suggest how this can be achieved



Cl2 + 2e-
2Cl-
Eo
=
+1.36 V
MnO2 + 4H+ + 2e-
Mn2+ + 2H2O
Eo
=
+1.23 V
when you study the variable oxidation states of chromium later on you will note that alkaline
conditions are used during oxidation – once again this modifies the associated electrode
potentials and hence reaction feasibility
consider the feasibility of coating a copper/nickel coin with zinc as a first step to making a
‘gold’ coin based on standard electrode potentials
during the operation of a battery the conditions change such that it goes ‘flat’
Summary Questions
Exam Style Questions
1–4
1–3
3
Page 207
Page 212
Page 246
A2 Chemistry AQA (Nelson Thornes)
Chemguide
204 - 7
Electrochemical, vanadium, reaction
feasibility
LUND 7 May 2017
20
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Learn the Daniel cell, zinc-carbon battery, lead acid battery
and
hydrogen fuel cell. Develop an understanding of the rest.
Electrochemical Cells

1800 – Alessandro Volta invents a primitive battery called a
voltaic pile consisting of alternative layers of silver (or copper)
and zinc discs separated by cardboard soaked in salt water
Non-rechargeable Cells (Primary Cells)
The electrochemical reaction is not reversible (cells can be used only once) as when discharging the
cell the chemicals are permanently changed.
Daniell Cell





a ‘wet cell’ invented by John Daniell in 1836
the porous pot was used (in place of the salt bridge shown
below) to allow the ions to migrate when the battery was
operating whilst preventing the solutions from mixing
or
without this barrier, even when no current was drawn, the
ZnSO4
copper ions would be reduced at the zinc anode thus
shortening the battery's life
negative ions will migrate in the same direction as the electrons around the circuit and
positive ions visa-versa
it was widely used at telegraph stations in the 19th century, however, its portability was
limited due to the liquid electrolytes that it contained
THE CATHODE
THE ANODE
Reduction takes place as the
positive pole of the battery
(half-cell with the highest
electrode potential)
accepts electrons from the
external circuit
Copper metal deposits on the
cathode – hence its mass
increases
Oxidation takes place as the
negative pole of the battery
(half-cell with the lowest
electrode potential)
releases electrons to the
external circuit
Zinc goes into aqueous solution
(hence the zinc plate loses
mass)
The solution becomes
more dilute hence the blue
colour fades
eYou could write
lots of notes but
ideally you
should be able to
work it all out
from this
schematic
LUND 7 May 2017
Zn(s)| Zn2+(aq)||Cu2+(aq)|Cu(s)
Eocell = EoRHS 0.34 -
EoLHS
- 0.76
Note that, unlike in
electrolysis, the cathode
is the positive terminal
in a battery. This is
because the cathode is
specified as the electrode
at which reduction takes
place.
=
+1.10V
Forgetting the double minus is a common
source of error
21
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
STANDARD Zinc – Carbon ‘Dry Cell’






1887 - Carl Gassner patented a dry cell variant
of the 1866 Leclanché wet cell
the electrodes are zinc and carbon, with an
acidic paste between them
zinc serves as both the anode and the
container, allowing the battery to be
completely self-contained and in effect more
portable and practical than wet cells
it could be used in any position as well rather
than on a flat surface and the risk of leaking
was greatly reduced
the cathode is a mixture of powdered (surface area contact) manganese dioxide and
graphite surrounding a solid graphite rod
the electrolyte is a paste of ammonium chloride inside the zinc can
+ CATHODE: Eo = +0.7V
- ANODE: Eo = -0.8V
2NH4+(aq) + 2e- → H2(g) + 2NH3(aq)
Zn(s) → Zn2+(aq) + 2e-
This means
the casing
gets thinner
in use
2NH3(aq) + Zn2+(aq) → [Zn(NH3)2]2+(aq)
prevents ammonia leaking (what shape is the complex ion)
+4
+3
2MnO2(s) + H2(g)→ Mn2O3(s) + H2O(l)
prevents a pressure build up from H2(g)
overall reaction in a STANDARD zinc-carbon cell is:
Note that overall
it’s the manganese
that is reduced
Zn(s) + 2MnO2(s) + 2NH4+(aq) → Mn2O3(s) + Zn(NH3)22+(aq) + H2O(l)

whilst zinc-carbon batteries are inexpensive they have very low power density so are only
useful in devices that draw very little current


the zinc container also becomes thinner when used as the zinc is oxidised
it also thins when not used as ammonium
NH4+(aq) + H2O(l) ⇋ H3O+(aq) +
chloride is acidic and slowly reacts with the zinc
NH3(aq)
which may lead to leakage
hence the service/shelf life of the battery is relatively short



the terminals of the battery are made of tin plated steel or brass to prevent the exposure the
zinc, not allowing it to corrode as quickly, thus adding to the total battery life
the seal usually is made of asphalt pitch, wax, or plastic to allow the cathode mix (when the
battery gets warm) to expand without rupturing the casing
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22
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
‘HEAVY DUTY’ Zinc Chloride Cell

zinc chloride cells are an improvement on the original zinccarbon cell giving a longer life (about 50% due to a variation in
the chemical mix and greater mass)

electrolyte ZnCl2 paste (cf NH4Cl in standard Zn-C)
- ANODE:
+ CATHODE:
+4
+3
2MnO2(s) + H2O(l) + 2e- → Mn2O3(s) + 2OH-(aq) Zn(s) → Zn2+(aq) + 2eoverall reaction in a HEAVY DUTY zinc-chloride cell is:
Zn(s) + 2MnO2(s) + H2O(l) → Mn2O3(s) + Zn2+(aq) + 2OH-(aq)

these were originally marketed around 50 years ago as "Heavy Duty” batteries, but since this
term is not standardised it is a misleading as they are much inferior to alkaline batteries
which have since been introduced and are around 300% better
Alkaline Cell Battery





sold under brand names such as ‘Duracell’ ‘Energizer’
these are more expensive than, but last considerably longer than, ‘ordinary zinc-carbon cells’
the cathode is manganese(IV) oxide powder
the anode is zinc powder (more surface area for increased rate of reaction therefore increased
electron flow to allow for heavy duty usage
the electrolyte is potassium hydroxide paste (hence ‘alkaline’)
+ CATHODE:
- ANODE:
+4
+3
2MnO2(s) + H2O(l) + 2e → Mn2O3(s) + 2OH-(aq)
Zn(s) + 2OH- (aq) → ZnO(s) + H2O(l) + 2e-
Regenerated
overall reaction is:
Zn(s) + 2MnO2(s) → Mn2O3(s) + ZnO(s)

no gases (which insulate the electrodes) are produced which is one reason why they don’t
suffer from a voltage drop as do zinc-carbon batteries when worked hard
NOTE: ‘Button’ batteries consist of a range of different types of system e.g. silver oxide,
alkaline and are a structural form, for a specific use (compact size/long life), rather than a
particular chemistry so are not described in this guide. At one time mercuric oxide was used
but this has ceased for obvious reasons.
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23
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Rechargeable Cells (Secondary Cells)
The electrochemical reaction is not reversible (cells can be used only once) as when discharging the
cell the chemicals are permanently changed
Lead-acid battery

1859 - Gaston Planté - the lead-acid cell: the first rechargeable battery

the cathode is a lead-antimony alloy grid coated with lead dioxide

the anode is a lead-antimony alloy grid coated in spongy lead (its porous to increase
surface area)

the electrolyte is ~ 6M sulfuric acid in which the plates are immersed

at the anode lead combines with sulphate ions to
create lead sulfate and release electrons
as the battery discharges, both plates build up lead
sulphate and water builds up in the acid thus
diluting it
the voltage is about 2 volts per cell, so by
combining six cells in series you get a 12-volt
battery
upon discharging the following reactions take
place:



+ CATHODE:
Eo = +1.68V
PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- → PbSO4(s) + 2H2O(l)
- ANODE:
Eo = -0.36V
Pb (s) + SO42-(aq) → PbSO4(s) + + 2eOverall:
Eo = 2.04V
PbO2(s) + 4H+(aq) + 2SO42-(aq) + Pb(s) → 2PbSO4(s) + 2H2O(l)






upon charging the reactions are reversed and this takes place when the car is running (using
the alternator) so that lead sulphate does not build up hence giving a ‘flat battery’
even when not in use, leakage of current takes place so there is a net usage of sulphuric acid
hence the need for it to be checked and topped up when the vehicle is serviced
given the emphasis on fuel economy the battery must not add too much weight to the vehicle
but must provide enough power to start the car even in cold weather
for most of the last century these batteries have become standard for starting cars as they can
produce short bursts of a high current for many decades and are easy and cheap to
manufacture
however they cannot be used to power a vehicle for longer journeys as they provide relatively
little energy per kilogram and suffer power loss as insulating lead sulphate builds up
why do you think a car battery could explode if it is overcharged
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24
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Nickel-Cadmium Cells ‘Nicads’








1950s - Nickel-Cadmium (NiCd) first appeared
these are much more portable than lead-acid batteries but are more expensive
they are manufactured in sizes/voltages to act as direct replacements for cheaper zinc-carbon
batteries
the fact that they can be recharged make them economical in the longer term
the cathode is made from nickel(III) oxyhydroxide
the anode is made from cadmium
the electrolyte is potassium hydroxide
when discharging the reactions are (reversed when charging):
+ CATHODE:
Eo = +0.52V
NiO(OH)(s) + H2O(l) + e- → Ni(OH)2 (s) + OH-(aq)
Note some sources state Ni(OH)3
here and omit H2O’s to balance
Overall:
- ANODE:
Eo = -0.81V
Cd(s) + 2OH-(aq) → Cd(OH)2(s) + 2e-
Eo = 1.33V
2NiO(OH)(s) + 2H2O(l) + Cd(s) → 2Ni(OH)2 (s) + Cd(OH)2(s)

cadmium is toxic so there are environmental issues regarding disposal
Nickel Metal Hydride Cells






Comparison of the discharge
1986 - NiMH battery was patented as
a
voltage of an alkaline battery (red)
bi-product from research on the
and a NiMH battery (blue). The
green line is the voltage at which
storage of hydrogen for use as an
the battery is considered dead
alternative energy source in the
1970s
some metallic alloys were observed
to
form hydrides that could capture (and
release) Hydrogen in volumes up to
nearly a thousand times their own
volume
compared to lead-acid and NiCd,
NiMH batteries have a higher storage capacity
they are more expensive than lead-acid and NiCd, but they are considered better for the
environment (lead and cadmium are toxic) – prices are falling but lithium ion batteries are
starting to gain some of the market
however, a NiCd battery has a lower self-discharge rate i.e. they hold their charge better
as with NiCd’s they offer slightly under 1.5V so may not work with some devices designed to
operate off the 1.5V of alkaline or zinc-carbon batteries
LUND 7 May 2017
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry



the cathode is made from nickel(III) oxyhydroxide (the same as NiCd’s)
the anode is made from an alloy of rare earth metals that can soak up hydrogen atoms with
the general formula AB5 where A can be lanthanum and B nickel
the electrolyte is potassium hydroxide
+ CATHODE:
Eo = +0.52V
NiO(OH)(s) + H2O(l) + e- → Ni(OH)2 (s) + OH-(aq)
Eo = -0.83V
- ANODE:
AB5H(s) + OH-(aq) → AB5(s) + H2O(l) + eOverall:
Eo = 1.35V
NiO(OH)(s) + AB5H(s) → Ni(OH)2 (s) + AB5
Lithium ion Cells

compared with NiMH and NiCd batteries of the same sizes or weights, Lithium Ion batteries
are designed to deliver the highest energy output

a single cell voltage is 3.7V, 3 times that of NiMH batteries so a simpler battery configuration
and better space utilization is achievable in devices such as cameras
they are relatively expensive as a computer chip is required to control charging and
discharging but do offer a high capacity (hence reducing mass/size)
Li Ion batteries contain no toxic heavy metals, such as mercury, cadmium or lead




the cathode is made from lithium cobalt oxide powder
the anode is lithium/graphite formulation (a lot of technological development was
required to prevent lithium oxidising and costing the electrode with insulating lithium
oxide



1996, the lithium ion polymer battery was developed from the lithium ion battery
these batteries hold their electrolyte in a solid polymer composite which can’t leak
the electrodes and separators are laminated to each other with the whole devices encased in a
flexible wrapping instead of a rigid metal casing, which means such batteries can be
specifically shaped to fit a particular device
Fuel Cells
The Hydrogen Economy


fossil fuels are at present the most economical way to power transportation
however, price rises commensurate with supply and demand, plus pollution issues such as the
greenhouse gas CO2 and acidic nitrous oxides (from atmospheric N2 + O2 in a hot engine) etc
are driving the need for an alternative
LUND 7 May 2017
26
A2 Unit 5 Energetics, Redox and Inorganic Chemistry






an alternative fuel is hydrogen which if combusted does not give the pollution problems
associated with hydrocarbons, the product is water
hydrogen can be obtained from the electrolysis of sea water in the longer term but at present
most hydrogen is still obtained from fossil fuels by steam methane reforming
this reacts steam with methane (natural gas) over a heated nickel catalyst to produce hydrogen
and carbon monoxide
energy is obviously required to obtain the hydrogen so fuel cells are not the ‘free’ energy from
water that is often suggested
however, given that nuclear reactors can’t be turned off, off peak generation could be one of
the means of generating less expensive hydrogen along with wind, hydro, solar and tidal
whichever, there may still be some pollution associated with hydrogen production
Hydrogen Fuel Cell





in a battery the chemical energy is stored within the electrodes and the solution
in a fuel cell hydrogen (fuel) and air (oxygen) are fed into the cell in a similar way that petrol
and air are fed into an internal combustion engine
the difference is that the chemicals are not combusted but react to produce electricity directly
this is more efficient than combusting the hydrogen (chemical → heat → kinetic → electrical
energy)
since a continuous supply of hydrogen is provided the voltage output remains constant



a typical fuel cell consists of two
platinum electrodes
these also act as catalysts to assist the
decomposition of hydrogen molecules
the electrodes are separated by a
polymer electrolyte (proton exchange
membrane) through which hydrogen
ions can migrate whilst the gases are
kept apart
+ CATHODE: Eo = +1.2V
- ANODE: Eo = 0.0V
4H+(aq) + O2(g) + 4e- → 2H2O(l)
H2(g) → 2H+(aq) + 2eZn(s) → Zn2+(aq) + 2e-
Overall:
Eo = 1.2V
O2(g) + 2 H2(g) → 2H2O(l)





a major issue is the storage and transportation of liquid hydrogen
research is currently being undertaken to develop hydrocarbon fuel cells so that car
manufacturers can rely of normal fuel tanks
this is more complex as it requires preliminary reforming of the fuel within the vehicle
many other alternatives are currently under investigation
it is likely in the interim that fuel cell/battery/petrol hybrids will be employed to a greater
extent
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27
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Overview




batteries reduce the need for expensive cabling and can provide power supplies to remote
places
non-rechargeable batteries are cheap and can be manufactured in all sizes e.g. button batteries
for watches, larger batteries for torches etc
however, they are thrown away afterwards along with energy and resources used in their
manufacture
rechargeable batteries are more expensive initially but less resources are wasted in the longer
term

they are vital in solar powered devices (and similar devices intent on
storing power for future use)

the lead and nicad’s contain toxic chemicals thus there are disposal
issues – they must not go to landfill sites and must be recycled

neither are suitable for vehicles as they add too much mass and
alternatives are being investigated

NiMH and lithium ion batteries are more environmentally friendly as
they do not contain toxic heavy metals

sodium-sulphur batteries do offer a better power per kg output but have
to operate at 300oC

another possibility is metal-air batteries (possible metals include
aluminium and zinc)

a major issue is power density i.e. how
much energy can be stored per kilogram of
battery, particularly where small size (e.g.
MP3 players) or small mass (transport) is
required

most batteries have performance that varies
with temperature (either way)

fuel cells only produce water (spacecraft can use fuel cells to provide drinking water) and will
eventually become the standard power source for vehicles which will reduce CO2 emissions
IF the hydrogen can be produced cheaply without fossil fuels
they provide a more efficient means of converting chemical energy into electrical energy
since it is direct rather than by a turbine
there are however problems associated with the transportation and storage of hydrogen and
the means of refuelling the vehicle



you might be asked to calculate a voltage from electrode potentials but should be aware
that the actual value will be unlikely to be this as it will not be operating under standard
conditions e.g variation in temperature
Exam Style Questions
Page 246
A2 Chemistry AQA (Nelson Thornes)
3
200 – 201, 208 - 211
tbd
LUND 7 May 2017
28
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Transition Metals
transition metal - d block elements that are able to form ions with a partially filled d sub shell


Sc is not a transition metal because its ion, Sc3+, is iso-electronic with Ar i.e. no d electrons
Zn is not a transition metal because its ion, Zn2+, has full d sub-shell

the electronic configuration for atoms and ions (remember to write 3d then 4s !) are written
left to right in order of increasing energy
whilst the 4s subshell is initially of a lower energy than an unoccupied 3d, hence filled first,
adding electrons to the 3d pushes the 4s electrons away from the nucleus thus raising their
energy
don’t forget that copper and chromium are not systematic
note that 4s electrons are always lost first when ions are formed and so first series transition
metal ions never have any 4s electrons present – so don’t even show 4s (4s0 is incorrect)
if you are showing ions using the electrons in boxes nomenclature then note that paired d
electrons are lost first (check Fe3+ and Co2+ for example) as mutual repulsion makes these
easier to remove



Physical Properties


typical metals i.e. malleable; ductile; good electrical and thermal conductors; all explained by
the same ideas taught in Foundation Chemistry – remind yourself of these.
melting point is higher than s block metals since there are more delocalised electrons holding
structure together – which also explains their greater mechanical strength
Chemical Properties




variable oxidation states
coloured ions (coloured rocks e.g. hæmatite, malachite,
typically include transition metal compounds)
catalysts
complex ions
Summary Questions
Page 215
A2 Chemistry AQA (Nelson Thornes)
Chemguide
Q
1, 2
214 - 5
Transition
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Complex Ions

a complex ion is a metal ion surrounded by ligands

ligands are molecules or ions which form dative (co-ordinate bonds) by donating electron
pair(s) - lone pair donor - to a central metal ion – lone pair acceptor

the coordination number is the number of coordinate bonds formed with the central metal ion
NOT necessarily the number of ligands


ligands (electron pair donors) are Lewis bases
transition metal ions (electron pair acceptors) are Lewis acids
Unidentate ligands

form only one co-ordinate bond with the TM ion
Neutral
NH3
Negative
OHSCNPr-
ammine
H2O
aqua
Cl-
hydroxo
thiocyanato
chloro
CN-
cyano
pavarotto (a rather large ligand not on the syllabus)
alphabetical order of name (prefixes i.e. di, tri etc ignored)
oxidation state of transition metal is given by roman numerals and this will only be the same as the
charge if all the ligands are neutral
name of metal is in Latin if complex ion is –ve
copper
=
cuprate
lead
iron
=
ferrate
vanadium
manganese
=
manganate
chromium
zinc
=
zincate
aluminium
e.g.
=
=
=
=
plumbate
vanadate
chromate
aluminate
tetrachlorocuprate(II) ion [CuCl4]2- (ends in ‘ate’ because it is anionic)
hexaaquacopper(II) ion [Cu(H2O)6]2+
NOT
3+
tetraamminedicyanocobalt(III) ion
number of each
[Co(NH3)4(CN)2]+
ligand
central metal ion and its oxidation
type of ligands
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30
number
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Bi-dentate ligands





two coordinate bonds per ligand
Neutral
1,2-diaminoethane (‘en’)
benzene-1,2-diol
Negative
ethandioate (oxalate) ions (C2O42-)
benzene-1,2-dicarboxylate
two co-ordinate bonds per molecule leading to chelated (crab) complexes
this is also called chelation
the possibility of a bidentate ligand acting as a bridge between two separate metal ions exists
the replacement of unidentate with bidentate ligands is favoured by entropy since the total
number of particles increases (see the section on entropy)
Multidentate ligands

more than two coordinate bonds per ligand
e.g.


ethylenediamminetetraacetate ion - EDTA4-
EDTA complexes are very stable – in effect a protective cage is formed around the transition
metal ion thus isolating it from a biological system
the replacement of unidentate with multidentate ligands is favoured by entropy since the total
number of particles increases (see the section on entropy)
Uses of
EDTA
1.
2.
3.
4.
Summary Questions
antidote to Hg/Pb poisoning (traps metal ions)
Ca2+ trap in blood transfusions – prevents clotting
removal of Ca2+ from hard water (e.g. in shampoo)
titimetric determination of metal ion concentration
Page 219
A2 Chemistry AQA (Nelson Thornes)
Chemguide
2
216 - 8
complex
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Shapes of Complex Ions
octahedral

whilst the coordination number is 6 there are 8 faces which
determines the allocated name based on an octahedral unit
tetrahedral

e.g. with H2O, NH3, EDTA, en
e.g. with Cl-
note that only 4 chloro ligands can fit around the
central metal ion due to their relatively large size
hence the complex ion adopts a tetrahedral geometry
linear
(often in silver(1)
and copper(I)
complexes
e.g.
[Ag(NH3)2]+
Tollens’ reagent
[Cu(NH3)2]+
square planar
e.g.
cis-platin [Pt(NH3)2Cl2]
(more on this later)
Extra info:
Summary Questions
Exam Style Questions
geometrical and optical isomerism are possible
cis [Co(NH3)4Cl2] +(aq) is violet
trans [Co(NH3)4Cl2] +(aq) is green
Page 219
Page 233
A2 Chemistry AQA (Nelson Thornes)
Chemguide
1
3
217 - 8
Complex shape
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Colours of Complex Ions






light absorbed depends on the E of electronic transitions to the next vacant energy level
E = hh is Planck’s constant
however, E(3d4s) in Ti3+ cannot account for lilac colour as E is too high (i.e. uv frequency)
in an isolated transition metal ion the d orbitals all have the same energy i.e. they are
degenerate
however, ligands split the 3d energy level so that E is of a lower value corresponding to the
energy of visible light
the reason for this is that the electrons donated by the ligand change the electronic
environment to different extents for different d-orbitals in different geometrical positions i.e
the are all raised in energy but to differing degrees

white light incident upon a transition metal solution or solid will have certain wavelengths
absorbed in accordance with the value of E when exciting an electron thus removing this
colour from the spectrum


colour observed is the complementary colour of light absorbed
hexaaquacopper(II) ions are blue as red is absorbed (see colour wheel)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry

Sc3+ and Zn2+ are not coloured as there are no partially filled 3d sub-shells which is
necessary for this to work i.e. to allow the promotion of d-electrons between d-ortbitals

colour depends on:
central metal ion
this is a major factor - obviously
oxidation state
e.g.
Fe(OH)2(s)  Fe(OH)3(s)
Cr2O72-(aq)  [Cr(H2O)6]3+(aq)
MnO4-(aq)  Mn2+(aq)
[Co(NH3)6]2+(aq)  [Co(NH3)6]3+(aq)
co-ordination number
has a significant effect on d-d splitting hence colour change
varies size of E and the type of d-d splitting
e.g.
[Co(H2O)6]2+(aq)  [Co(Cl)4]2-(aq)
[Cu(H2O)6]2+(aq)  [Cu(Cl)4]2-(aq)
octahedral geometry yields 2 higher 3 lower
tetrahedral geometry yields 3 higher 2 lower
type of ligand
stronger bonding causes greater d-d spitting hence shorter
wavelength absorbed
(spectrochemical series (Cl- < H2O < NH3 < en < CN-)
e.g.
[Cu(H2O)6]2+(aq)  [Cu(NH3)4(H2O)2]2+(aq)
[Co(H2O)6]2+(aq)  [Co(NH3)6]2+(aq)
YOU MUST LEARN THESE COLOURS
[Cu(H2O)6]2+(aq)
Blue
[CuCl4]2-(aq)
Yellow
2+
Adding cHCl to [Cu(H2O)6] (aq) gives green!
[Cu(NH3)4(H2O)2]2+(aq)
Deep
Blue
NH3(aq) definitive test for Cu2+(aq)
[Co(H2O)6]2+(aq)
Pink
[CoCl4]2-(aq)
Blue
cobalt chloride paper is a test for water
[Co(NH3)6]2+(aq)
Yellow
[Co(NH3)6]3+(aq)
[Fe(SCN)(H2O)5]2+(aq)
sensitive test for the presence of Fe3+(aq)
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Brown
Blood
Red
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Colorimetry


you should re-familiarise yourself with the nature of light and why we see colour
two ways in which chemicals can interact with light result in:
absorption spectra (star light through a
a planetary atmosphere, chlorophyll)
emission spectra (e.g. flame tests, street lights)








absorption in the visible light varies according to the complex ion present, path length and
concentration
absorption of aqua complexes is relatively weak so colours are not very intense
certain complexing agents (e.g. EDTA) increase colour intensity to aid detection and
determination
for example complexing [Fe(H2O)6]3+(aq) ions with colourless thiocyanate ions (SCN-) to
produce the more deeply coloured [Fe(SCN)(H2O)5]2+(aq) complex ion which can be used to
detect low concentrations of iron in substances like tea by comparing absorbance against a
calibration curve of known concentrations.
in a colorimeter interchangeable filters are used to illuminate the sample with its
complementary colour where absorption is greatest hence sensitivity optimised
it is also possible to determine the formula of a complex ion
as the complexing agent is added to separate batches of the transition metal sample the
intensity of the colour will increase until there is no more transition metal ions for it to
combine with (the volume would be kept constant using water)
this allows us to determine the number of complex ions that combine with a transition metal
where both concentrations are known (ideally the same value)
Summary Questions
Exam Style Questions
Page 222
Page 233
A2 Chemistry AQA (Nelson Thornes)
Chemguide
1, 2
3 (if not already done)
220 - 222
Colour
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Variable Oxidation States


across the d block the effective nuclear charge increases
hence relative stability of 2+ oxidation state cf 3+ increases as the e- are better held

as the number of 4s/3d electrons increases from Ti to Mn so does the maximum oxidation
state (Sc=3, Ti = 4, V = 5, Cr = 6, Mn = 7).
thereafter the maximum declines as effective nuclear charge increases suggesting that the 4s
and unpaired 3d electrons only are involved






transition metals with charges > +3 cannot exist in aqueous solution where they exist as
oxoanions instead e.g. MnO4-, Cr2O72-(aq), CrO42-(aq) with covalent bonding between the
oxygen and the transition metal (can you suggest their shape?)
this can be explained in two ways – a lot of energy would be required to form a 4+ ion, and
if it existed it would have a large charge density (thus be highly polarising) so would react
with water molecules and decompose them
+2 state tends to be reducing, as exemplified by Fe2+ in the manganate(VII) titration
some +2 ions are unstable in air due to aerial oxidation (where they are themselves reducing
agents)
this can be pH dependent and occurs more readily in alkaline conditions e.g. keeping
[Fe(H2O)6]2+(aq) in acidic solution helps it resist aerial oxidation to [Fe(H2O)6]3+(aq)
Redox Titrations








higher oxidation states (typically +4 and higher) are good oxidising agents
MnO4-(aq) and Cr2O72-(aq) are particularly good as oxidising agents in redox titrations
you will need to balance redox equations – some revision may be necessary here
the titration procedure is pretty much the same as with acid-base titrations
acidic conditions are employed and you should be able to carry out a calculation to
determine the minimum amount of sulphuric acid required
sulphuric acid is preferred (can you explain why each of the following: hydrochloric,
nitric and ethanoic might not be suitable?)
Manganate(VII) shows a distinct colour change whilst dichromate(VI)
require an indicator since both Cr2O72-(aq) and [Cr(H2O)6]3+ ions are coloured
the indicator used is sodium N-phenylamine-4-sulphonate which turns from colourless to
purple at the end point
in case you wondered it works by changing colour at a particular electrode potential, in this
case + 0.84V
Summary Questions
Exam Style Questions
Page 228
Page 233
A2 Chemistry AQA (Nelson Thornes)
Chemguide
2, 3
1(some extra reading will be needed)
223 - 226
Variable oxidation state, redox titration
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Chromium
The oxidation of cobalt
is covered later on
Reducing Chromium (VI) to Chromium(III)

chromium in chromate(VI) can be reduced by reacting it with Zn in the presence of conc. HCl
(this also releases H2 - a reducing atmosphere)

the feasibility can be demonstrated using electrode potentials
Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O(l)
orange
green
3+
Cr (aq)
+
e

Cr2+(aq)
green
blue
Zn2+(aq)
+
2e
Zn(s)

Eo
+1.33V
-0.41V
-0.76V
the +2 state is readily oxidised back to the +3 state by air – unless preserved in a reducing
(e.g. H2) atmosphere
Ox State
+6
+3
+2
Chromium
CrrO7
orange
CrO42- (shape?)
yellow
3+
[Cr(H2O)6]
Green
2+
[Cr(H2O)6]
Blue
2-
Oxidising Chromium (III) to Chromium(VI)



oxidation of transition metals tends to occur more readily in alkaline conditions
Iron(II) sulphate for example is kept in acidic conditions to prevent aerial oxidation
a plausible reason is that it is harder to remove electrons from the positively charged complex
present in acidic solutions

Cr3+ can be oxidised to chromate(VI) by H2O2 in strongly alkaline conditions

initially further deprotonation in xs OH- produces deep green [Cr(OH)6]3-(aq)
[Cr(H2O)6]3+(aq)

+

6OH(aq)-
[Cr(OH)6]3-(aq)
+
subsequent oxidation with H2O2 yields yellow chromate(VI) CrO42-(aq)
tetraoxochromate(VI)) ions)
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6H2O(l)
ions (aka
A2 Unit 5 Energetics, Redox and Inorganic Chemistry


writing half equations for redox reactions under alkaline conditions is a little more
tricky than under acidic conditions – but here is a useful ‘cheat’
do it exactly as if it was in acidic conditions then cancel out the hydrogen ions by adding
hydroxide ions equally to both sides
i.e.
H2O2(aq)
+
2H+(aq)

+
2e-

2OH-(aq)
2H2O(l)
cancelling 2H+, and subsequently H2O
[Cr(OH)6]3-(aq) 
H2O2(aq)
+
2e-
CrO42-(aq)
+
2H2O(l)
+
2H+(aq) + 3e-
CrO42-(aq)
+
4H2O(l)
cancelling 2H+, and subsequently H2O
[Cr(OH)6]3-(aq)
+
2OH-(aq)

+3e-
now balancing for electrons and combining:
2[Cr(OH)6]3-(aq)


+

3H2O2(aq)
2CrO42-(aq)
+ 2OH-(aq) + 8H2O(l)
upon acidification orange dichromate(VI) Cr2O72-(aq) is formed – this is an acid-base
equilibrium NOT a redox – check the oxidation state of chromium (you should be able to
write the equation)
see if you can write half equations and then a full balanced equation for other oxidations
carried out in alkaline conditions e.g.:
[Co(NH3)6]2+(aq) to [Co(NH3)6]3+(aq) by aerial oxygen
Co(OH)2(s) oxidised by H2O2 to Co(OH)3(s).
Fe(OH)2(s) to Fe(OH)3(s) by aerial oxygen
Summary Questions
Page 228
A2 Chemistry AQA (Nelson Thornes)
Chemguide
1, 3
204 – 7, 226 – 228, 238
Chromium
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Catalysis by Transition Metals


catalysts enable a different mechanism with a different activation energy (hence different rate)
whilst being chemically unchanged at the end of a reaction
NOTE CATALYSTS CHANGE THE VALUE EA NOT THE SHAPE OF THE CURVE.
Heterogeneous Catalysis


heterogeneous catalysis – the catalyst is in a different phase to the reactants
typically transition metals or their compounds are used e.g.
manufacture of ammonia
Haber Process
Fe
catalytic converters
Pt and Rh
hardening fats (making margarine) Hydrogenation
Ni
(adsorption onto the surface of the solid nickel catalyst weakens π bonds)
manufacture of nitric acid
Ostwald Process
Pt and Rh
manufacture of sulphuric acid
Contact Process
V2O5

adsorption occurs onto active sites and consequently:
weakens the bonds in the reactants hence lowers the activation energy
improves the stereochemistry for collisions by orienting molecules favourably
provides a localised relatively high concentration of reactants
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry


adsorption must be strong enough to hold the reactant for long enough to promote a reaction
but must not be too strong (e.g. as with tungsten) otherwise regeneration of active sites is too
slow as the product undergoes desorption from the surface
in general the strength of adsorption decreases from left to right in the transition metals and
the desorption occurs more readily in the same direction hence Fe (Haber process) Co and Ni
(Hydrogenation) are commonly used since they offer a compromise
Catalytic Converters




platinum and rhodium are coated onto a
honeycomb ceramic material (minimises
costs whilst providing a large surface
area = increased rate) since adsorption
only occurs at the surface (expensive
metal underneath would be wasted)
the reactant gases form weak bonds with
the surface of the catalyst (adsorption)
this weakens their bonds thus lowering the activation energy (additionally the catalyst also
helps promote more favourable molecular orientation)
this is followed by desorption in which the products depart
the catalyst selected provides bonding strong enough to hold the reactant gases on the surface
whilst not preventing the products from leaving thus blocking an active site
Write equations
for these



CO and NO react to form CO2 and N2
NO also reacts with uncombusted hydrocarbons to produce CO2, H2O and N2
poisoning can occur if impurities contaminate the active sites e.g. sulphur dioxide and lead
poisons catalytic converters hence unleaded low sulphur fuels must be used
in addition the finely coated Pt/Rh can be lost from the surface
this reduces efficiency and can result in an MOT failure and a large bill
Haber process



Pea sized Fe lumps are the catalyst - large surface
area (enhanced by an aluminium oxide promoter) to
increase rate without requiring an even higher
temperature (energy cost plus unfavourable for yield)
the iron catalyst does not effect the equilibrium
position as both forward and backward reactions are
favoured equally
sulphur impurities (present in natural gas) can poison the iron so ‘scrubbing’ is carried out to
remove the sulphur compounds (carbon monoxide can also be a problem) however it
eventually has to be replaced
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Methanol Production



methanol is used as a chemical feedstock and as an additive to petrol
it can be manufactured by the reversible reaction between carbon monoxide and hydrogen in
the presence of a copper catalyst or alternatively Cr2O3
the reactants (‘synthesis gas’) are manufactured from the reaction of methane or propane with
steam
Contact process


V2O5 used rather than faster Pt as lowers costs and less prone to poisoning
specific use of the variable oxidation states of transition metals is made
+5
V2O5(s)
+4
+
SO2(g)
+
+4
V2O4(s)

1 O2(g)
2
+
V2O4(s)
+
SO3(g)
+5

V2O5(s)
Reactants
Product
Homogeneous Catalysis



homogeneous catalysis – the catalyst is in the same phase to the reactants
in this case the reaction proceeds via an intermediate species and will typically have a two
step reaction profile with two activation energies both less than that for the uncatalysed
reaction
same phase as reactants (e.g. all in solution):
acid catalysed esterification
enzymes in biological systems
chlorine free radicals (formed by the action of UV light on CFC’s) and ozone (O3) depletion
Peroxodisulphate and iodide ions


redox reaction between peroxodisulphate (S2O82-) and iodide ions is slow as both are
negatively charged
catalysed by iron(II) (or iron(III) – either will do) – note oppositely charged ions now react
+2
Fe2+(aq)
+3
+
S2O82-(aq)
+
2I-(aq)

2SO42-(aq)
+
+3
2Fe3+(aq)
+2

I2(aq)
Reactants

Fe3+(aq)
Products
this could be followed experimentally using a colorimeter
LUND 7 May 2017
41
+
Fe2+(aq)
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Autocatalysis


some reactions can speed up rather than slow down
relative to the initial rate
there can be several reasons for this:
an oxide layer on a surface is being removed before
the acid gets at the metal
the reaction might be exothermic
the product is itself a catalyst for the reaction –
this is autocatalysis
Conc.
Slow reduction in conc.
initially
Faster reduction in conc. as
autocatalysis begins
Reaction slows down
as reagents run out
time

in all cases the reaction will eventually start to slow down as reactants are used up

e.g. manganate(VII) initially reacts slowly with ethanedioate ions (from oxalic acid)
2MnO4-(aq) + 5C2O42-(aq)


2Mn2+(aq) + 10CO2(g) + 8H2O(l)
the Mn2+ ions produced autocatalyse the reaction, hence it actually speeds up once started
they change to Mn3+ initially but are changed back in the next step:
MnO4-(aq)


+ 16H+(aq)
+ 4Mn2+(aq)
+ 8H+(aq)

5Mn3+(aq)
+
4H2O(l)
2Mn3+(aq)
+ C2O42-(aq)

2Mn2+(aq)
+
2CO2(g)
this could be followed experimentally using a colorimeter
Summary Questions
Exam Style Questions
Exam Style Questions
Page 232
Page 233
Page 247
A2 Chemistry AQA (Nelson Thornes)
Chemguide
1-4
1, 2, 4
4, 8
229 - 230
Heterogeneous catalysis
LUND 7 May 2017
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Applications of Transition Metal Complexes
EDTA

this is an ethylenediamminetetraacetate ion or EDTA for short

EDTA complexes are very stable – in effect a protective cage is formed around the transition
metal ion thus isolating it from a biological system
the replacement of unidentate with multidentate ligands is favoured by entropy since the total
number of particles increases (see the section on entropy)

Uses of
EDTA
1.
2.
3.
4.
antidote to Hg/Pb poisoning (traps metal ions)
Ca2+ trap in blood transfusions – prevents clotting
removal of Ca2+ from hard water (e.g. in shampoo)
titimetric determination of metal ion concentration
Haemoglobin









haem – forms four co-ordinate bonds (tetradentate) with Fe2+ (a porphyrin structure)
N in globin – a protein forms a fifth to form haemoglobin
O2 or H2O form the sixth bond in oxyhaemoglobin or deoxyhaemoglobin
as oxygen is a poor ligand it is easily released in cells
lack of iron in the blood can cause anaemia as insufficient oxygen is transported resulting in
tiredness and fatigue (or is that homework)
taking iron tablets which contain soluble iron(II) sulphate counteracts this
CO (which is a better ligand than oxygen) bonds with haemoglobin more strongly to form the
relatively stable carboxyhaemoglobin thus reducing the bloods capacity to transport oxygen
cyanide ions act in a similar way
similar structures are found in a range of biologically important substances such as vitamin
B12 (cobalt), and chlorophyll (magnesium)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Cis-platin



cis-platin - [Pt(NH3)2Cl2] - is used in treating certain cancers (note
the DNA complexing explanation on page 219 is wrong as it’s the
chlorines that are displaced)
its full name is cis-diamminedichloroplatinum(II) in case you
wondered
it forms DNA cross links via the platinum which damage the
cancer cells
http://www.youtube.com/watch?v=Wq_up2uQRDo&feature=related


however, it does have side effects as it also effects normal cells e.g. renal toxicity, bone
marrow suppression (loss of white blood cells increases the risk of other infections), fatigue
and hearing loss and can also induce nausea and vomiting.
testing renal function, blood and hearing is recommended before each cycle of therapy. so a
cautious approach to dosage is necessary
http://www.cancerhelp.org.uk/about-cancer/treatment/cancer-drugs/cisplatin

geometrical isomerism possible the other form being trans-platin (which has no effect on
cancer for stereochemical reasons
Tollens’ Reagent







diamminesilver(I) ion
[Ag(NH3)2]+(aq)
formed in ammonical silver nitrate (Tollens’ reagent)
used in silver mirror test for aldehydes and distinguish them from ketones
[Ag(NH3)2]+(aq) is reduced to Ag – the silver mirror - and NH3 is displaced
see if you can write a balanced redox equation under alkaline conditions
complexing prevents the precipitation of Ag2O in alkaline conditions which would otherwise
mask the test

[Ag(NH3)2]+(aq) is also formed when testing for silver chloride and silver bromide with the
addition of ammonia following the silver nitrate test
the ligand displacement allows the precipitate to be re-solvated
A2 Chemistry AQA (Nelson Thornes)
Chemguide
69 – 70, 217 – 219, 241
Cis-platin, haemoglobin, edta, Tollens’
LUND 7 May 2017
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Metal-aqua Ions





in aqueous solution tri-positive and di-positive TM ions form hexaaqua- complexes
ligands (electron pair donors) are Lewis bases and transition metal ions (electron pair
acceptors) are Lewis acids
the presence of the ligand creates the familiar colour of transition metal solutions
this can also be locked into the crystalline structure – again imparting colour (.xH2O)
e.g. white anhydrous copper(II) sulphate dissolves (very) exothermically in water to form
blue hexaaquacopper(II) ions
[Cu(H2O)6]2+(aq)
[Co(H2O)6]2+(aq)
[Fe(H2O)6]2+(aq)
[Cr(H2O)6]3+(aq)
[Fe(H2O)6]3+(aq)
[Al(H2O)6]3+(aq)
Blue
Pink
Green
Green
Yellow
Colourless
Actually they are lilac but the presence of a
small amount of orange [Fe(H2O)5OH]2+(aq)
makes it appear yellow (see later)
Hydrolysis of Metal-aqua Ions

water ligands have increased Oδ-H δ+ bond polarity which promotes the abstraction of a
hydrogen ion by another water molecule compared to that which takes place in the autoionisation of water (see Kw)
Note the charge !
[Fe(H2O)6]3+(aq)
Pale Lilac
ACID






+
[Fe(H2O)5OH]2+(aq) +
H3O+(aq)
Orange
oxonium ion
BASE
ACID
H2O(l)
BASE
this is hydrolysis (reaction with water) and makes the solution pH around 2 for a 1M solution
it would be wrong to assume that most populous species is the complex ion on the RHS as the
equilibria still lies strongly to the LHS, however it does result in an increased hydrogen ion
concentration
further deprotonation very limited as water is a relatively weak base and the charge on the
complex ion is less positive and so the Oδ-H δ+ bond polarity is less pronounced.
the solution appears yellow as the orange colour is more intense than the pale lilac
relative acidity of M3+ cf M2+ reflects the relative polarising power of the central transition
metal on the polarity of the O-H bond of the ligand
the chemistry of Al3+(aq) is similar to tri-positive transition metal ions
Summary Questions
How science works
Page 238
Page 243
Page 235
A2 Chemistry AQA (Nelson Thornes)
Chemguide
1, 2
1
Theories of acids
234 - 7
Acidity of hexaaqua
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45
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Reactions of Transition Metal-aqua ions with OH-(aq)

upon the addition of sodium hydroxide a metal hydroxide precipitation occurs
Cr(OH)3(s)
Fe(OH)3(s)
Al(OH)3(s)
green
orange
white
Cu(OH)2(s)
Co(OH)2(s)
Fe(OH)2(s)
pale blue
blue
green
goes brown on standing
turns orange on standing
(some texts will show the water ligands as well)


deprotonation reactions occur with the addition of basic hydroxide ions to a greater extent
than in aqueous solution and this eventually presents a neutral complex ion and thus
precipitation (hydroxide ions are a better base than water)
alternatively this reaction can be depicted as a reaction between a hydroxide ion and the
oxonium ion (based on the equation on page 45) with a consequential shift in equilibria
– either is acceptable as the outcome is essentially the same, but direct abstraction by
OH-(aq) is easier to produce equations for.
[Cr(H2O)6]3+(aq)



+
OH-(aq)
[Cr(H2O)5OH]2+(aq) + H2O(l)
[Cr(H2O)5OH]2+(aq) +
OH-(aq)
[Cr(H2O)4(OH)2]+(aq) + H2O(l)
[Cr(H2O)4(OH)2]+(aq) +
OH-(aq)
[Cr(H2O)3(OH)3](s) + H2O(l)
here there is no repulsion since there is no charge and so the complexes can hydrogen bond
together producing a gelatinous precipitate.
all the hydroxide precipitates are solvated by the addition of acid which reverses the equilibria
(equations must be known – acid + base  salt + water).
precipitates insoluble in XS sodium hydroxide solution, but soluble in acid, are basic
Fe(OH)3(s)
Cu(OH)2(s)
Fe(OH)2(s)
Co(OH)2(s)
XS NaOH(aq)

XS sodium hydroxide can cause the precipitate to re-dissolve for amphoteric hydroxides due
to further deprotonation for:
Cr(OH)3(s)
Al(OH)3(s)
[Cr(H2O)3(OH)3](s)

+
OH-(aq)
[Cr(H2O)2(OH)4]-(aq) + H2O(l)
the charged particle created can now be solvated and if the sodium hydroxide solution is
concentrated enough then the hexahydroxo- complex can eventually be formed.
[Cr(H2O)2(OH)4]-(aq) +
OH-(aq)
[Cr(H2O)(OH)5]2-(aq) + H2O(l)
[Cr(H2O)(OH)5]2-(aq) +
OH-(aq)
[Cr(OH)6]3-(aq)
LUND 7 May 2017
46
+ H2O(l)
Any one of these ions would be
credited in the exam, but this one
is easiest to remember.
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Reactions of Transition Metal-aqua Ions in Solution with CO32-(aq)

carbonate ions react with oxonium ions to yield carbon dioxide gas
CO32-(aq)

+
2H3O+(aq)
CO2(g) +
3H2O(l)
the relatively high acidity of tri-positive transition metal results in their metal carbonates
being unstable
[Fe(H2O)6]3+(aq)
+
2H3O+(aq)
CO32-(aq)
+
[Fe(H2O)5OH]2+(aq) +
H2O(l)

CO2(g) +
H3O+(aq)
3H2O(l)

thus carbonate ions react with tri-positive transition metal ions to produce carbon dioxide gas
HENCE FIZZING in addition to a hydroxide precipitate

the deprotonation equilibrium is shifted to the RHS as the carbonate ion removes the
oxonium ion until the neutral triaquatrihydroxo- complex (the precipitate) is obtained
the overall equation should be known but can be derived on the basis of a shift in equilibrium
as the carbonate ion reacts with oxonium ion (it is slightly harder to work it out starting with
the carbonate ion abstracting a hydrogen ion directly, but is also acceptable)

2[M(H2O)6]3+(aq) +
Cr(OH)3(s)
green
3CO32-(aq)
2[M(H2O)3(OH)3](s) + 3CO2(g) + 3H2O(l)
Fe(OH)3(s)
orange
Al(OH)3(s)

as previously it is soluble in acid since hydroxides are bases

it is not soluble in xs sodium carbonate solution as the concentration of hydroxide ions is
relatively low so no further deprotonation occurs.

di-positive transition metal carbonates are stable as the oxonium ion concentration is
relatively low
hence metal carbonate precipitates are produced and no CO2(g) is evolved

CoCO3(s)
mauve
CuCO3(s)
A2 Chemistry AQA (Nelson Thornes)
Chemguide
blue-green
237 - 8
aqua ions hydroxide,
aqua ions carbonate
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FeCO3(s)
green
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Reactions of Transition Metal-aqua Ions in Solution with NH3(aq)


ammonia undergoes hydrolysis with water although it is only a weak alkali as the equilibria is
biased to the LHS
the hydroxide ion concentration is thus much lower than with fully ionised NaOH(aq)
NH3(aq)
BASE


+
H2O(l)
ACID
NH4+(aq)
ACID
+
OH-(aq)
BASE
as with sodium hydroxide a metal hydroxide precipitation can be assumed to occur by direct
deprotonation by the hydroxide ion.
alternatively this reaction can be depicted as a reaction between a hydroxide ion and the
oxonium ion (based on the equation on page 45) with a consequential shift in equilibria
or abstraction of a proton by an ammonia molecule – either is acceptable as the outcome
is essentially the same, but direct abstraction by OH-(aq) is easier to produce equations
for and offers commonality with sodium hydroxide – so less to remember.
Cr(OH)3(s)
Fe(OH)3(s)
Al(OH)3(s)
green
orange
white
Cu(OH)2(s)
Co(OH)2(s)
Fe(OH)2(s)
pale blue
blue
green
goes brown on standing
turns orange on standing
(some texts will show the water ligands as well)

XS ammonia causes to [Cu(H2O)4(OH)2](s) to re-dissolve
[Cu(H2O)4(OH)2](s) + 4NH3(aq)
pale blue
[Cu(NH3)4(H2O)2]2+(aq)
deep blue
Note that two water
molecules remain
+ 2H2O(l)
+ 2OH-(aq)

unlike adding xs NaOH further deprotonation is NOT an explanation of why solvation
occurs since the concentration of hydroxide ions in the weakly alkaline ammonia
solution is inadequate to further deprotonate the uncharged species with the far less
polar Oδ-H δ+ bond polarity

however, ammonia is a better ligand than hydroxide ions and water as its lone pair are
less well held than for the relatively more electronegative oxygen hence ligand
displacement can occur once the concentration of ammonia is high enough to bias the
equilibria adequately

the geometry is still octahedral as the ammonia ligand is roughly the same size as the
water ligand – but is slightly distorted as the Cu-O bonds are weaker (water is a poorer
ligand) therefore longer
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry

similarly:
[Co(H2O)4(OH)2](s) + 6NH3(aq)
blue
[Co(NH3)6]2+(aq)
+
4H2O(l)
+ 2OH-(aq)
pale yellow
darkens on standing due to aerial oxidation
to brown [Co(NH3)6]3+(aq)
[Cr(H2O)3(OH)3](s)
green
[Cr(NH3)6]3+(aq)
purple

+ 6NH3(aq)
+
3H2O(l)
+ 3OH-(aq)
aluminium, iron(II) and (III) DO NOT undergo these ligand displacement reactions
hence their hydroxide precipitates are insoluble in xs ammonia
Summary Questions
Exam Style Questions
Page 243
Page 244-5
A2 Chemistry AQA (Nelson Thornes)
Chemguide
3, 4
2, 4, 5
239 – 240, 242 - 243
aqua ions ammonia
Ligand Displacement Reactions (summary)

replacement is stepwise and is reversible – this also means that it may be incomplete e.g. in
the example below whilst [Cu(Cl)4]2-(aq) is yellow what will be seen is a green solution,
implying that complete substitution has not occurred

some ligands form more stable complexes than others which will effect the equilibrium
position for a given set of conditions
there is a relationship between effectiveness as a ligand, base and nucleophile as all three are
essentially related to the availability of a lone pair e.g. cyanide ions are better than water/O2 –
see haemoglobin section


substitution by bidentate and multidentate ligands are also favoured by an increase in entropy
as there will be a consequential increase in the number of independent particles

moles of EDTA used = moles of T.M. ion and is useful for the analysis of the concentration
of a transition metal (by titration or colorimetry)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry

concentration and temperature will effect the position of equilibrium and hence the colour e.g.
[Co(H2O)6]2+(aq)
pink
+ 4Cl-(aq)
[Co(Cl)4]2-(aq) +
blue
6H2O(l)
[Cu(H2O)6]2+(aq)
blue
+ 4Cl-(aq)
[Cu(Cl)4]2-(aq) +
yellow
6H2O(l)



adding concentrated HCl(aq) will shift the equilibrium to the RHS
the forward reaction is endothermic and hence a temperature change will effect the colour
the above equilibrium is the basis of cobalt chloride paper used as a chemical test for water
(e.g. in the combustion of a hydrocarbon demonstration)

the co-ordination number is changed as water is replaced by the far larger chloride ion –
usually associated with a significant colour change
when water ligands are replaced by the similar sized ligands, e.g. ammonia, the octahedral
geometry is retained


the stability of the +II oxidation state can be altered by the ligand present e.g. yellow
[Co(NH3)6]2+(aq) is readily oxidised to brown [Co(NH3)6]3+(aq) by oxygen from the air
whereas [Co(H2O)6]2+ (aq) is not.

similarly alkaline conditions/hydroxyl ligands can also make oxidation easier e.g.


Co(OH)2(s) is readily oxidised by H2O2 to brown Co(OH)3(s).
Fe(OH)2(s) is readily oxidised by oxygen from air to Fe(OH)3(s)

in case you wondered this is because the electrode potentials in your text book relate to the
hexaaqua complex and changing the ligand will modify its value
a similar change in relative electrode potentials occurs when alkaline rather than acidic
conditions are employed – see the oxidation of chromium

Summary Questions
Exam Style Questions
Page 241
Page 244
A2 Chemistry AQA (Nelson Thornes)
Chemguide
1-4
1, 3
239 – 241, 242 – 243, 228
Ligand displacement, cobalt, iron hydroxide
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
THIS IS AVAILABLE IN COLOUR – See ‘Mr Lund’s Classes’ at READ CHEMISTRY
Example
Solution
NaOH(aq)
XS NaOH(aq)
NH3(aq)
XS NH3(aq)
Na2CO3(aq)
Other test
Result
Fe3+
Fe2(SO4)3
FeCl3
Yellow
Orange-brown
gelatinous ppt.
Insoluble
Orange-brown
gelatinous ppt.
Insoluble
Orange-brown
gelatinous ppt.
KMnO4
(acidified in
dilute H2SO4)
No reaction
occurs.
Fe2+
FeSO4
Green
Green ppt., then,
slowly turning
orange due to
aerial [O]
Insoluble
Insoluble
Dark green ppt.
KMnO4
(acidified in
dilute H2SO4)
Purple
MnO4decolourise
dFe2+  Fe3+
Cu2+
CuSO4
Pale blue
Pale blue ppt.
Insoluble
Green ppt.,
then, slowly
turning orange
due to aerial
[O]
Pale blue ppt.
Soluble (deep
blue solution)
Blue-green ppt.
Co2+
CoCl2
Pink
Blue ppt. then,
slowly turning
brown
Insoluble
Blue ppt.
Mauve ppt.
Al3+
Al2(SO4)3
Colourless
White ppt.
Soluble
(colourless
solution)
White
gelatinous ppt.
Soluble (yellow
solution which
darkens on
standing due to
aerial [O])
Insoluble
Cr3+
CrCl3
Green
Grey-green ppt.
Soluble (green
solution)
Grey-green to
grey-blue
gelatinous ppt.
Soluble
(slightly) -in
excess a violet
or pink solution
formed.
Green ppt.
ION
White ppt.
Please note that the above colours are illustrative (it should be obvious that variables beyond control are concentration, colour on computer screen, colour of printer !!!)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
THIS IS AVAILABLE IN COLOUR – See ‘Mr Lund’s Classes’ at READ CHEMISTRY
Colour
Cr
Mn
Fe
Co
Cu
(all hexaaqua
unless
otherwise)
Mn2+ (almost)
Colourless
Lilac/Purple
Pink
Blue
Green
Yellow
Orange/Brown
[Cr(NH3)6]2+
MnO4-
CoCO3
[CoCl4]2-
Cr2+
Co(OH)2
Cu2+
Cu(OH)2
[Cu(NH3)4]2+
CuCO3
[Co(NH3)6]2+
[CuCl4]2-
Cr3+
Cr(OH)3
[Cr(OH)6]3CrO42_
Fe2+
Fe(OH)2
FeCO3
Fe3+
Cr2O72_
Fe(OH)3
Co(OH)3
[Co(NH3)6]3+
[Fe(SCN)]2+
Co2+
Red
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Cu2O
A2 Unit 5 Energetics, Redox and Inorganic Chemistry
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
A2 Module 5 Thermodynamics and Further Inorganic Chemistry
Thermodynamics
Enthalpy change (ΔH)
be able to define and apply the terms enthalpy of formation, ionisation enthalpy, enthalpy of atomisation of an element and of a compound, bond
dissociation enthalpy, electron affinity, lattice enthalpy (defined as either lattice dissociation or lattice formation), enthalpy of hydration and enthalpy
of solution.
be able to construct a Born.Haber cycle to calculate lattice enthalpies from experimental data
be able to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration
be able to use mean bond enthalpies to calculate an approximate value of ΔH for other reactions
be able to explain why values from mean bond enthalpy calculations differ from those determined from enthalpy cycles
Free-energy change (ΔG) and entropy change (ΔS)
understand that ΔH, whilst important, is not sufficient to explain spontaneous change (e.g. spontaneous endothermic reactions)
understand that the concept of increasing disorder (entropy change ΔS) accounts for the above deficiency, illustrated by physical change (e.g. melting,
evaporation) and chemical change (e.g. dissolution, evolution of CO2 from hydrogencarbonates with acid).
be able to calculate entropy changes from absolute entropy values
understand that the balance between entropy and enthalpy determines the feasibility of a reaction; know that this is given by the relationship
ΔG = ΔH . TΔS
be able to use this equation to determine how ΔG varies with temperature
be able to use this relationship to determine the temperature at which a reaction is feasible
Periodicity
Study of the reactions of Period 3 elements Na . Ar to illustrate periodic trends
be able to describe trends in the reactions of the elements with water, limited to Na and Mg.
be able to describe the trends in the reactions of the elements Na, Mg, Al, Si, P and S with oxygen, limited to the formation of Na2O, MgO, Al2O3,
SiO2, P4O10 and SO2.
A survey of the acid-base properties of the oxides of Period 3 elements
be able to understand the link between the physical properties of the highest oxides of the elements Na - S in terms of their structure and bonding.
be able to describe the reactions of the oxides of the elements Na - S with water, limited to Na2O, MgO, Al2O3, SiO2, P4O10, SO2.and SO3.
know the change in pH of the resulting solutions across the Period.
be able to explain the trends in these properties in terms of the type of bonding present.
be able to write equations for the reactions which occur between these oxides and given simple acids and bases.
Redox Equilibria
Redox equations
be able to apply the electron transfer model of redox, including oxidation states and half equations to d-block elements
Electrode potentials
know the IUPAC convention for writing half-equations for electrode reactions.
know and be able to use the conventional representation of cells.
understand how cells are used to measure electrode potentials by reference to the standard hydrogen electrode
know the importance of the conditions when measuring the electrode potential, Eo (Nernst equation not required).
know that standard electrode potential, Eo , refers to conditions of 298 K, 100 kPa and 1.00 mol dm−3 solution of ions.
Electrochemical series
know that standard electrode potentials can be listed as an electrochemical series.
be able to use Eo values to predict the direction of simple redox reactions and to calculate the e.m.f of a cell.
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Electrochemical cells
appreciate that electrochemical cells can be used as a commercial source of electrical energy
appreciate that cells can be non-rechargeable (irreversible), rechargeable and fuel cells
be able to use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells and to
deduce the e.m.f. of a cell
understand the electrode reactions of a hydrogen-oxygen fuel cell and appreciate that a fuel cell does not need to be
electrically recharged
appreciate the benefits and risks to society associated with the use of these cells
Transition Metals
General properties of transition metals
know that transition metal characteristics of elements Ti - Cu arise from an incomplete d sub-level in atoms or ions.
know that these characteristics include complex formation, formation of coloured ions, variable oxidation state and catalytic activity.
Complex formation be able to define the term ligand
be able to define the term ligand
know that co-ordinate bonding is involved in complex formation.
understand that a complex is a central metal ion surrounded by ligands.
know the meaning of co-ordination number.
understand that ligands can be unidentate (e.g. H 2O, NH3 and Cl−) or bidentate (e.g. NH2CH2CH2NH2 and C2O42-) or multidentate (e.g.
EDTA4- ).
know that haem is an iron(II) complex with a multidentate ligand.
Shapes of complex ions
know that transition metal ions commonly form octahedral complexes with small ligands (e.g. H 2O and NH3).
know that transition metal ions commonly form tetrahedral complexes with larger ligands (e.g. Cl− ).
Know that square planar complexes are also formed e.g. cisplatin
know that Ag+ commonly forms linear complexes, (e.g. [Ag(NH3)2]+, [Ag(S2O3)2]3− and [Ag(CN)2]−).
Formation of coloured ions
know that transition metal ions can be identified by their colour, limited to the complexes in this module.
know that colour changes arise from changes in oxidation state, co-ordination number and ligand.
know that colour arises from electronic transitions from the ground state to excited states: ΔE = hv.
appreciate that this absorption of visible light is used in spectrometry to determine the concentration of coloured ions
Variable oxidation states
know that transition elements show variable oxidation states.
know that Cr3+ and Cr2+ are formed by reduction of Cr2O72- by zinc in acid solution.
know the redox titrations of Fe2+ with MnO4- and Cr2O72− in acid solution.
be able to perform calculations for these titrations and for others when the reductant and its oxidation product are given.
know the oxidation of Co2+ by air in ammoniacal solution.
know the oxidations in alkaline solution of Co2+ and Cr3+ by H2O2.
Catalysis
know that transition metals and their compounds can act as heterogeneous and homogeneous catalysts.
Heterogeneous
know that a heterogeneous catalyst is in a different phase from the reactants and that the reaction occurs at the surface.
understand the use of a support medium to maximize the surface area and minimize the costs and that the reaction occurs at the surface
understand how V2O5 actsas a catalyst in the Contact Process.
know that a Cr2O3 catalyst is used in the manufacture of methanol from carbon monoxide and hydrogen.
Know that Fe is used as a catalyst in the Haber process
know that catalysts can become poisoned by impurities and consequently have reduced efficiency; know that this has a cost implication (e.g. poisoning
by sulphur in the Haber Process and by lead in catalytic converters in cars)
Homogeneous
know that when catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species (e.g. the reaction
between I− and S2O82- catalysed by Fe2+ and autocatalysis by Mn2+ in titrations of C2O42- with MnO4−)
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A2 Unit 5 Energetics, Redox and Inorganic Chemistry
Other applications of transition metal complexes
understand the importance of variable oxidation states in catalysis; both heterogeneous and homogeneous catalysts
understand that Fe(II) in haemoglobin enables oxygen to be transported in the blood, and why CO is toxic
know that the Pt(II) complex cisplatin is used as an anticancer drug
understand that [Ag(NH3)2] + is used in Tollens’ reagent to distinguish between aldehydes and ketones
Reactions of Inorganic Compounds in Aqueous Solution
Lewis acids and bases
know the definitions of a Lewis acid and Lewis base;
understand the importance of lone pair electrons in co-ordinate bond formation.
Metal-aqua ions
know that metal.aqua ions are formed in aqueous solution:
[M(H2O)6]2+, limited to M = Fe, Co and Cu
[M(H2O)6]3+, limited to M = Al, Cr and Fe
know that these aqua ions can be present in the solid state (e.g. FeSO 4.7H2O and Co(NO3)2.6H2O).
Acidity or hydrolysis reactions
understand the equilibria
[M(H2O)6]2+ + H2O  [M(H2O)5(OH)]+ + H3O+
and
[M(H2O)6]3+ + H2O  [M(H2O)5(OH)]2+ + H3O+
to show generation of acidic solutions with M3+, and very weakly acidic solutions with M2+.
understand that the acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+ in terms of the polarising power (charge/size ratio) of the metal ion
be able to describe and explain the simple test-tube reactions of M2+ (aq) ions, limited to M = Fe, Co and Cu, and of M 3+ (aq) ions, limited to M = Al,
Cr and Fe, with the bases OH¯, NH3 and CO32−.
know that MCO3 is formed but that M2(CO3)3 is not formed.
know that some metal hydroxides show amphoteric character by dissolving in both acids and bases (e.g. hydroxides of Al 3+ and Cr3+).
know the equilibrium reaction 2CrO42- + 2H+ ⇋ Cr2O72− + H2O
Substitution reactions
understand that the ligands NH3 and H2O are similar in size and are uncharged, and that ligand exchange occurs without change of co-ordination
number (e.g. Co2+ and Cr3+).
know that substitution may be incomplete (e.g. the formation of [Cu(NH3)4(H2O)2]2+).
understand that the Cl− ligand is larger than these uncharged ligands and that ligand exchange can involve a change of co-ordination
number (e.g. Co2+ and Cu2+).
know that substitution of unidentate ligand with a bidentate or a multidentate ligand leads to a more stable complex.
understand this chelate effect in terms of a positive entropy change in these reactions
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AS Chemistry Unit 2
57
AS Chemistry Unit 2
58