Redox and standard electrode potential
Specification
(a) describe redox in terms of electron transfer, use oxidation states (numbers) to identify
redox reactions and decide which species have been oxidised and which reduced;
(b) write ion-electron half equations for redox reactions for which stoichiometric information
is supplied, and use titration and other data to carry out appropriate calculations;
(c) show awareness that electrode processes represent oxidations and reductions;
(d) recall and use the redox systems specified below, including the appropriate colour change
and ion/electron half-equations: Cu2+(aq)|Cu(s); Zn2+(aq)|Zn(s); H+(aq)|H2(g) Pt; Fe3+(aq),
Fe2+(aq)|Pt; MnO4– (aq), Mn2+(aq)|Pt; X2(g)|2X (aq) (X = Cl , Br , I );
(e) use redox systems in addition to those in (d), for which all relevant information is
supplied;
(f) describe simple electrochemical cells involving;
(i) metal/metal ion electrodes, and
(ii) electrodes based on different oxidation states of the same element.
(g) explain and use the term standard electrode potential especially (i) the use of the standard
hydrogen electrode in determining standard electrode potential; (ii) to calculate standard
potentials of cells formed by combining different electrodes and; (iii) to predict the feasibility
of specified reactions.
Redox reactions
Redox (reduction-oxidation) reaction = chemical reaction in which electrons are
transferred from one species (the reducing agent), which is oxidised by losing an
electron, to another species (the oxidising agent), which is reduced by receiving the
electron.
Reducing agent = reduces another species by giving it electrons and so is oxidised by
its loss of electrons. A reducing agent contains at least one element which is oxidised.
Oxidising agent = oxidises another species by removing electrons from it and so is
reduced by its gain of electrons. An oxidising agent contains at least one element
which is reduced (e.g. although Cr is reduced, potassium dichromate is the oxidising
agent).
Oxidation states
Oxidation states measure how much an atom has been oxidised compared with the
element i.e. it shows how many electrons must be added (if negative) or lost (if
positive) from the element to get to the atom in its current state.
The point of oxidation numbers is to allow us to track electrons in redox reactions as
it is normally hard to track changes in electrons just with a balanced equation.
Oxidation states allow us to measure oxidation and reduction in covalent compounds
as well as ionic compounds.
Oxidation = loss of electrons or oxidation state becoming more positive/less negative.
Reduction = gain of electrons or oxidation state becoming less positive/more negative.
NB do not mix up charges and oxidation states – they are similar but not the same: +3
and -2 are oxidation states but 3+ and 2- are charges.
, Oxidation state rules
Elements have an oxidation state of 0. This includes non-bonded elements (e.g. Xe)
and elements bonded to itself (O2 and C60).
The sum of all oxidation states of atoms in a compound equals the charge on the
species – for a neutral species, the sum equals zero.
The oxidation state of a common ion is equal to the charge of the common ion e.g. -2
for SO42-:
The oxidation state is negative for the most electronegative atom in a compound.
Some elements almost always have the same oxidation numbers in their compounds
(CHOF 12 – CHO have exceptions):
Element Oxidation Exceptions (F MH PF – F off MH u Pig F)
state
Chlorine -1 Except in compounds with O or F (e.g. NaClO).
Hydroge +1 Except in metal hydrides (made up of metal ion and hydride ion(s)
n e.g. NaH or AlH3) where it is -1. NB this is because hydrogen is
present as the hydride ion, H–.
Oxygen -2 Except in peroxides (compound contain the peroxide ion e.g.
H2O2) where -1. NB this is because oxygen is present as the
peroxide ion, O2–2.
Except in F2O where O is +2 since fluorine is more
electronegative (O is the second most electronegative element)
and is -1.
Fluorine -1
Group 1 +1
metals
Group 2 +2
metals
Disproportionation reaction
Disproportionation = a type of redox reaction where atoms of the same element
become oxidised and reduced to form two different products.
Equations
Ionic equations
Specification
(a) describe redox in terms of electron transfer, use oxidation states (numbers) to identify
redox reactions and decide which species have been oxidised and which reduced;
(b) write ion-electron half equations for redox reactions for which stoichiometric information
is supplied, and use titration and other data to carry out appropriate calculations;
(c) show awareness that electrode processes represent oxidations and reductions;
(d) recall and use the redox systems specified below, including the appropriate colour change
and ion/electron half-equations: Cu2+(aq)|Cu(s); Zn2+(aq)|Zn(s); H+(aq)|H2(g) Pt; Fe3+(aq),
Fe2+(aq)|Pt; MnO4– (aq), Mn2+(aq)|Pt; X2(g)|2X (aq) (X = Cl , Br , I );
(e) use redox systems in addition to those in (d), for which all relevant information is
supplied;
(f) describe simple electrochemical cells involving;
(i) metal/metal ion electrodes, and
(ii) electrodes based on different oxidation states of the same element.
(g) explain and use the term standard electrode potential especially (i) the use of the standard
hydrogen electrode in determining standard electrode potential; (ii) to calculate standard
potentials of cells formed by combining different electrodes and; (iii) to predict the feasibility
of specified reactions.
Redox reactions
Redox (reduction-oxidation) reaction = chemical reaction in which electrons are
transferred from one species (the reducing agent), which is oxidised by losing an
electron, to another species (the oxidising agent), which is reduced by receiving the
electron.
Reducing agent = reduces another species by giving it electrons and so is oxidised by
its loss of electrons. A reducing agent contains at least one element which is oxidised.
Oxidising agent = oxidises another species by removing electrons from it and so is
reduced by its gain of electrons. An oxidising agent contains at least one element
which is reduced (e.g. although Cr is reduced, potassium dichromate is the oxidising
agent).
Oxidation states
Oxidation states measure how much an atom has been oxidised compared with the
element i.e. it shows how many electrons must be added (if negative) or lost (if
positive) from the element to get to the atom in its current state.
The point of oxidation numbers is to allow us to track electrons in redox reactions as
it is normally hard to track changes in electrons just with a balanced equation.
Oxidation states allow us to measure oxidation and reduction in covalent compounds
as well as ionic compounds.
Oxidation = loss of electrons or oxidation state becoming more positive/less negative.
Reduction = gain of electrons or oxidation state becoming less positive/more negative.
NB do not mix up charges and oxidation states – they are similar but not the same: +3
and -2 are oxidation states but 3+ and 2- are charges.
, Oxidation state rules
Elements have an oxidation state of 0. This includes non-bonded elements (e.g. Xe)
and elements bonded to itself (O2 and C60).
The sum of all oxidation states of atoms in a compound equals the charge on the
species – for a neutral species, the sum equals zero.
The oxidation state of a common ion is equal to the charge of the common ion e.g. -2
for SO42-:
The oxidation state is negative for the most electronegative atom in a compound.
Some elements almost always have the same oxidation numbers in their compounds
(CHOF 12 – CHO have exceptions):
Element Oxidation Exceptions (F MH PF – F off MH u Pig F)
state
Chlorine -1 Except in compounds with O or F (e.g. NaClO).
Hydroge +1 Except in metal hydrides (made up of metal ion and hydride ion(s)
n e.g. NaH or AlH3) where it is -1. NB this is because hydrogen is
present as the hydride ion, H–.
Oxygen -2 Except in peroxides (compound contain the peroxide ion e.g.
H2O2) where -1. NB this is because oxygen is present as the
peroxide ion, O2–2.
Except in F2O where O is +2 since fluorine is more
electronegative (O is the second most electronegative element)
and is -1.
Fluorine -1
Group 1 +1
metals
Group 2 +2
metals
Disproportionation reaction
Disproportionation = a type of redox reaction where atoms of the same element
become oxidised and reduced to form two different products.
Equations
Ionic equations
