Redox Reactions — IMAT Chemistry Formulas and Traps

Oxidation Numbers: Rules and Calculation

Working out an oxidation number from a formula

Sum of oxidation numbers

\sum (\text{number of atoms} \times \text{oxidation number}) = \text{overall charge}

overall charge0 for a neutral compound; the ion's charge for an ion

The range of oxidation numbers, and fractional averages

Limits for a main-group element

\text{lowest} = \text{outer electrons} - 8 \qquad \text{highest} = \text{outer electrons}

outer electronsvalence electrons: 5 for N and P, 6 for O and S, 7 for the halogens

The rules for assigning oxidation numbers

Atom or situation	Oxidation number	Exception or example
An uncombined element	0	$\mathrm{Na}$ , $\mathrm{O_2}$ , $\mathrm{Cl_2}$ , $\mathrm{P_4}$ , $\mathrm{S_8}$ are all 0
A simple (monatomic) ion	Equal to its charge	$\mathrm{Mg^{2+}}$ is +2, $\mathrm{Cl^-}$ is −1
Fluorine in a compound	−1 always	The most electronegative element, no exceptions
Group 1, Group 2, aluminium	+1, +2, +3	Almost no exceptions in IMAT
Hydrogen	+1	−1 in metal hydrides: $\mathrm{NaH}$ , $\mathrm{LiH}$ , $\mathrm{CaH_2}$
Oxygen	−2	−1 in peroxides ( $\mathrm{H_2O_2}$ , $\mathrm{Na_2O_2}$ ); +2 in $\mathrm{OF_2}$ Peroxides are the exception IMAT likes most: in hydrogen peroxide each O is −1.
Cl, Br, I	−1 in halides	Positive when bonded to O: +1 in $\mathrm{HClO}$ , +7 in $\mathrm{HClO_4}$

Common traps

Oxygen is not always −2, and hydrogen is not always +1

In peroxides each oxygen is −1, and in

\mathrm{OF_2}

oxygen is +2. In metal hydrides such as NaH, hydrogen is −1. Applying the usual values blindly to these compounds gives a wrong oxidation number for every other atom.

Divide by the number of atoms of the unknown element

In

\mathrm{Cr_2O_7^{2-}}

the oxygens contribute −14 and the ion is −2, so the two chromium atoms together are +12 and EACH is +6. Forgetting to divide gives +12, an impossible value that IMAT sometimes offers.

Oxidation numbers can be fractions, but not irrational

An average such as

+8/3

in

\mathrm{Fe_3O_4}

or +2.5 in

\mathrm{S_4O_6^{2-}}

is a genuine result. Every oxidation number is a whole-number total shared among whole numbers of atoms, so it is always a ratio of integers: never irrational.

Recognising Redox: Oxidation, Reduction and the Agents

Learn this subtopic in the notes

Oxidation and reduction as electron loss and gain

Two half-equations, one transfer

\underbrace{\mathrm{Zn \rightarrow Zn^{2+} + 2e^-}}_{\text{oxidation}} \qquad \underbrace{\mathrm{Cu^{2+} + 2e^- \rightarrow Cu}}_{\text{reduction}}

$\mathrm{e^-}$ an electron; on the right in oxidation, on the left in reduction

Common oxidising and reducing agents

Agent	Acts as	Becomes (in acid)	Electrons per particle
Permanganate $\mathrm{MnO_4^-}$ (purple)	Oxidising agent	$\mathrm{Mn^{2+}}$ (almost colourless)	5 gained
Dichromate $\mathrm{Cr_2O_7^{2-}}$ (orange)	Oxidising agent	$\mathrm{Cr^{3+}}$ (green)	6 gained (3 per Cr)
Chlorine $\mathrm{Cl_2}$	Oxidising agent	$\mathrm{Cl^-}$	2 gained
Concentrated $\mathrm{HNO_3}$	Oxidising agent	$\mathrm{NO_2}$ (dilute acid gives NO)	1 gained (3 for NO)
Hydrogen peroxide $\mathrm{H_2O_2}$	Either	$\mathrm{H_2O}$ as oxidant; $\mathrm{O_2}$ as reductant	2
Reactive metals (Mg, Zn, Fe)	Reducing agent	Metal ions	Equal to the ion's charge, lost
Carbon, CO, $\mathrm{H_2}$	Reducing agent	$\mathrm{CO_2}$ , $\mathrm{H_2O}$	Used to extract metals from oxides

Common traps

Oxidation raises the oxidation number

Losing electrons, which are negative, makes an atom more positive. So oxidation is an INCREASE in oxidation number and reduction is a DECREASE. Mixing up the direction reverses every later answer about agents.

No free element does not mean no redox

A reaction can be redox with compounds on both sides: in

\mathrm{2FeCl_3 + SnCl_2 \rightarrow 2FeCl_2 + SnCl_4}

, iron falls from +3 to +2 and tin rises from +2 to +4. Always check the oxidation numbers rather than looking for an element.

A decomposition is not automatically redox

Breaking a carbonate into an oxide and carbon dioxide changes no oxidation numbers. Breaking a chlorate into a chloride and oxygen does. Decide by the numbers, not by the reaction type.

The oxidising agent is the one that is reduced

The name says what it does to the OTHER substance. An oxidising agent takes electrons, so its own oxidation number goes down. Picking the species that is oxidised as the oxidising agent is the classic slip.

Name the reactant, not the product

When zinc reacts and becomes

\mathrm{Zn^{2+}}

, the reducing agent is zinc metal, not the zinc ion it turns into. Options often list both forms of the same element.

Hydrogen peroxide can be either agent

With oxygen at −1, between 0 and −2,

\mathrm{H_2O_2}

oxidises iodide ions (becoming water) but reduces permanganate (becoming oxygen gas). Options that call it only an oxidising agent are incomplete.

Balancing Redox Equations

Learn this subtopic in the notes

Balancing by matching electrons lost and gained

The balancing condition

(\text{atoms oxidised}) \times (\text{rise per atom}) = (\text{atoms reduced}) \times (\text{fall per atom})

Half-equations in acidic solution

Two common reduction half-equations in acid

\mathrm{MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O} \qquad \mathrm{Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O}

Disproportionation and comproportionation

Splitting one species between two products

(\text{atoms going up}) \times (\text{rise}) = (\text{atoms going down}) \times (\text{fall})

Common traps

Atom balance alone can hide a wrong electron count

Some wrong options balance the metal atoms but not the electrons or the spectator ions. Check three things: the electron ratio, every element, and (for ionic equations) the total charge. An equation is balanced only when all three agree.

Electrons go on the side that is more positive

After adding water and

\mathrm{H^+}

, compare the total charges. Electrons are added to the side with the larger positive charge, enough to make the two sides equal. In a reduction they end up on the left; in an oxidation, on the right.

In disproportionation one species is both agents

In

\mathrm{Cl_2 + H_2O \rightarrow HCl + HClO}

, chlorine is both the oxidising agent and the reducing agent: one Cl atom falls to −1 while the other rises to +1. An option saying the reaction is not redox because only one element changes is wrong.

The Reactivity Series and Displacement Reactions

Learn this subtopic in the notes

Predicting displacement reactions of metals and halogens

Metal displacement, as an ionic equation

\mathrm{M_{(more\ reactive)} + N^{n+}_{(aq)} \rightarrow M^{n+}_{(aq)} + N_{(less\ reactive)}}

$\mathrm{M}$ the metal higher in the series (reducing agent)
$\mathrm{N^{n+}}$ the ion of the metal lower in the series (oxidising agent)

The reactivity series of metals

Metal	With cold water	With dilute HCl	How it is extracted
Potassium, sodium	Violent: hydroxide and $\mathrm{H_2}$	Dangerously violent	Electrolysis of the molten chloride
Calcium	Steady: hydroxide and $\mathrm{H_2}$	Vigorous	Electrolysis of the molten chloride
Magnesium	Very slow (fast with steam)	Vigorous, $\mathrm{H_2}$	Electrolysis
Aluminium	None (oxide layer)	Slow at first, then steady	Electrolysis of molten $\mathrm{Al_2O_3}$
Zinc, iron	None (react with steam when hot)	Moderate, $\mathrm{H_2}$	Heating the oxide with carbon or CO
Lead	None	Very slow	Heating the oxide with carbon
Copper, silver, gold	None	None	Reduction of ores; silver and gold also found as the metal

Higher in the table means a stronger reducing agent; the ions of metals lower down are stronger oxidising agents.

Common traps

For halogens, the more reactive one is the oxidising agent

A reactive metal is a strong REDUCING agent because it loses electrons easily. A reactive halogen is a strong OXIDISING agent because it gains electrons easily. In

\mathrm{Cl_2 + 2NaI \rightarrow 2NaCl + I_2}

, chlorine is reduced and the iodide ion is oxidised.

Electrochemical Cells and Electrolysis

Learn this subtopic in the notes

Standard reduction potentials and the cell voltage

Standard cell potential

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

$E^\circ_{\text{cathode}}$ reduction potential of the half-cell that is reduced (the higher value)
$E^\circ_{\text{anode}}$ reduction potential of the half-cell that is oxidised

Galvanic and electrolytic cells compared

Feature	Galvanic (voltaic) cell	Electrolytic cell
Energy change	Chemical to electrical; the reaction is spontaneous	Electrical to chemical; the reaction is forced
Anode	Oxidation; negative electrode	Oxidation; positive electrode
Cathode	Reduction; positive electrode	Reduction; negative electrode
Electrons in the wire	From anode to cathode	Pushed by the supply into the cathode and drawn out of the anode
Examples	Daniell cell (zinc and copper), batteries, fuel cells	Electrolysis of brine, aluminium extraction, electroplating

Electrolysis: what forms at each electrode

Electrolyte (inert electrodes unless stated)	At the cathode (negative)	At the anode (positive)
Molten sodium chloride	Sodium metal	Chlorine
Concentrated aqueous sodium chloride (brine)	Hydrogen; sodium hydroxide is left in solution	Chlorine
Dilute sulfuric acid	Hydrogen	Oxygen, half the volume of the hydrogen
Aqueous copper(II) sulfate	Copper	Oxygen; the solution becomes acidic
Aqueous copper(II) sulfate, copper electrodes	Copper is deposited	The copper anode dissolves as $\mathrm{Cu^{2+}}$ (used to purify copper)
Molten aluminium oxide in cryolite	Aluminium	Oxygen, which burns the carbon anode to $\mathrm{CO_2}$

Common traps

The anode's sign changes, its job does not

Oxidation always happens at the anode, and reduction at the cathode. But the anode is negative in a galvanic cell and positive in an electrolytic cell. Learn the electrodes by their reactions, not by their charges.

Never multiply E° by the equation's coefficients

E° is a property of the half-reaction, like a temperature, not an amount. Doubling a half-equation to balance electrons leaves its E° unchanged. Options that double or triple a potential are built from this mistake.

A reactive metal is not deposited from its aqueous solution

Electrolysing aqueous sodium, potassium or magnesium salts gives hydrogen at the cathode, because water is reduced more easily than those ions. The metal itself is obtained only by electrolysing the molten compound.

Redox Reactions formulas

Oxidation Numbers: Rules and Calculation

Working out an oxidation number from a formula

The range of oxidation numbers, and fractional averages

The rules for assigning oxidation numbers

Common traps

Recognising Redox: Oxidation, Reduction and the Agents

Oxidation and reduction as electron loss and gain

Common oxidising and reducing agents

Common traps

Balancing Redox Equations

Balancing by matching electrons lost and gained

Half-equations in acidic solution

Disproportionation and comproportionation

Common traps

The Reactivity Series and Displacement Reactions

Predicting displacement reactions of metals and halogens

The reactivity series of metals

Common traps

Electrochemical Cells and Electrolysis

Standard reduction potentials and the cell voltage

Galvanic and electrolytic cells compared

Electrolysis: what forms at each electrode

Common traps

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