Oxidation, reduction and redox reactions ✅

Oxidation

Loss of electrons, increase in oxidation state, gain of oxygen

Oxidising agents

- A chemical that causes another chemical to be oxidised by accepting electrons.

- Are reduced themselves: the reactant which GAINS electrons

Describe an example of an oxidising agent

- e.g. sulfuric acid, not sulphur.

- Only sulphur is reduced but we still say sulphuric acid is the oxidising agent.

Give an example of a half equation to show how copper is oxidised.

Cu → Cu²⁺ + 2e−

Reduction

Gain of electrons, decrease in oxidation state, oxygen removed

Reducing agent

- A chemical that causes another chemical to be reduced by donating electrons.

- Are oxidised themselves: the reactant which LOSES electrons.

Give an example of a half equation to show how oxygen is reduced.

O₂ + 4e− ’→2O

Redox reaction

Chemicals are both oxidised and reduced in the reaction.

. What are half equations?

They show the movement of electrons within reactions.

What are the rules for half equations?

- Atoms on each side of the arrow must be balanced. Overall charge on each side of the arrow must also be balanced.

- Electrons on the right of the arrow is oxidation

- Electrons of the left is reduction

Write the redox reaction for the following half equations: Cu → Cu²⁺ + 2e−

O₂ + 4e− → 2O

Ionic: 2Cu + O₂ → Cu²⁺ + 2O (Cancel out electrons)

Oxidation state/number

- Used to see what has been oxidised / reduced in a redox reaction.

-Tells us about the distribution of electrons between elements of different electronegativities in a molecule.

What does it mean if the oxidation state goes up by 1 in an ionic compound?

1 electron has been lost to another element.

Define oxidation states in terms of covalent compounds.

Where the electrons spend more of their time (as if were ionic)

Oxidation state rules (best to see table in notes)

1. Uncombined elements: Always 0 e.g. Cl₂

2. Monatomic ions: Ox. number same as charge on ion e.g. Ca²⁺ = +2

3. Group 1: Always +1 e.g. KCL

4. Group 2: Always +2 e.g. CaO

5. Aluminium: Always +3 e.g. Al₂O₃

6. Hydrogen: +1 (except in hydrides where it is -1) e.g. HF (hydride example

- NaH)

7. Chlorine: -1 (except if in a compound with F and O - it would have a positive

value) e.g. KCl (Cl has a value of +3 in CIF₃)

8. Fluorine: Always -1 e.g. KF

9. Oxygen: -2 (except in peroxides it's -1 and in OF₂ it's +2) e.g. Li₂O (O has the value -1 in H₂O₂

Give the rules for oxidation states in order of priority. (Preferred table, uplearn one)

1. Uncombined elements = 0

2. Monatomic ions = charge on ion

3. Compounds = sum of oxidation states is equal to overall charge

4. Group 1 atoms = +1

5. Group 2 atoms = +2

6. Fluorine = -1

7. Hydrogen = +1

8. Oxygen = -2

9. Chlorine = -1

What are the 'usually' rules?

Fluorine: -1 Hydrogen: +1 Oxygen: -2 Chlorine: -1

Fairies Hate Orange Clover 'Usually' because they are overruled by previous rules.

What are the rules for oxidation states?

1. Every element in it's uncombined state has an oxidation of 0

2. Some elements always have the same oxidation state in all their compounds.

3. The sum of all the oxidation states in a compound must equal 0.

4. The sum of the oxidation states in a complex ion must equal the *overall

charge* of the ion.

5. The more electronegative element is given the negative* oxidation

state.

What does a positive/negative number show in oxidation states?

- Positive number = element has lost electrons and been oxidised.

- Negative number = element has gained electrons and been reduced.

Work out the oxidation states of SiF₄

SiF₄

Si = +4

F = -1

Work out the oxidation states of H₂O₂

H₂O₂

H = +1

O = -1

Work out the oxidation states of S₂O₃²⁻

S₂O₃²⁻

S₂ = +2

O₃ = -2

Work out the oxidation states of C₂O₄²⁻

C₂O₄²⁻

C₂ = +6

O₄ = -2

Work out the oxidation states of Cl₂

Cl₂ = 0

Work out the oxidation states of IO₃⁻

IO₃⁻

I = +5

O = -2

What is the most electronegative element?

Fluorine

What is oxidation and reduction in terms of oxidation number?

- Reduction is a decrease in oxidation number

- Oxidation is an increase in oxidation number

What can we work out using oxidation states in redox reactions?

Which elements are oxidised and reduced.

What are spectator ions?

Ions that do not participate in a reaction

What is oxidised and what is reduced? Which is the oxidising and reducing agent? 2Na + Cl₂ --> 2NaCl

2Na + Cl₂ → 2NaCl

Na: 0 to +1 so oxidised, reducing agent

Cl: to -1 so reduced, oxidising agent

What are the rules for balancing half equations?

1. Write down the species before and after a reaction

2. Balance any atoms apart from oxygen and hydrogen (deal with these later)

3. Balance any oxygens with H₂O

4. Balance any hydrogens with H⁺ ions

5. Balance charges with electrons (e⁻)

What are the conditions of these reactions?

Acidic, so has H⁺ ions

Aqueous, so has H₂O

Write a half equation showing the conversion of MnO₄⁻ to Mn²⁺

1. MnO₄⁻ ’ Mn²⁺ (before and after)

2. MnO₄⁻ ’ Mn²⁺ + 4H₂O (balance oxygens with water)

3. MnO₄⁻ + 8H⁺ ’ Mn²⁺ + 4H₂O (balance hydrogens with H⁺)

4. MnO₄⁻ + 8H⁺ + 5e⁻ ’ Mn²⁺ + 4H₂O (balance charges with electrons)

. What are ionic equations?

2 half equations combined to make a full ionic equation.

Make sure electrons balance.

Then cancel out electrons.

Write an ionic equation for these 2 half equations: Fe²⁺ --> Fe³⁺ + e⁻

MnO₄⁻ + 8H⁺ + 5e⁻ --> Mn²⁺ + 4H₂O

Fe²⁺ → Fe³⁺ + e⁻

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

5Fe²⁺ → 5Fe³⁺ + 5e⁻

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

Ionic:

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

How do we show oxidation states in ionic compounds?

Bracketed Roman numerals to show oxidation states e.g. copper (I) oxide

What can be said about the oxidation states of ionic compounds?

It is equal to the charge on the ion.

e.g. in lithium chloride, Li⁺ has a charge of 1+ and an oxidation state of +1, Cl⁻

has a charge of 1- and an oxidation state of -1

. How do we name compounds with multiple ions?

According to it's oxidation state.

E.g. ClO₂⁻ is called chlorate (III) ion as chlorine has an oxidation state of +3

Give the full name of this ion. SO₄²⁻

SO₄²⁻

Sulfate (VI) ion

Name the following ions: ClO⁻ ClO₂⁻ClO₃⁻ ClO₄⁻

ClO⁻ = Chlorate (I)

ClO₂⁻ = Chlorate (III)

ClO₃⁻ = Chlorate (V)

ClO₄⁻= Chlorate (VII)

Named after oxidation state of Cl

Are acid base reactions redox

No. Oxidation states don't change.

In a covalent compound, which element is the reducing agent?

The element which is reduced is the most electronegative element and is the oxidising agent

Give the full name of this sulfate ion: SO₄²⁻

Oxidation state O₄ = -8, S = +6

So, Sulfate (VI)

What is a disproportionation reaction?

A reaction in which an element is both oxidised and reduced in the same reaction is called a disproportionation reaction.

Balancing Equation

How do we balance half equations?

1) Identify the oxidising/reducing agent

2) Balance the atoms

3) Balance the charges

Deduce the half

equation for chlorine in the equation

below.

Cl₂ + Fe²⁺ → Cl⁻ + Fe³⁺

Cl₂ + 2e⁻ → 2Cl⁻

Given the ionic equation below, write the balanced half equations for the reaction.

Half-equations should always be given with the lowest coefficients. Then state the oxidising and reducing agents. 2Al (s) + 3Br₂ (l) → 2AlBr₃ (s)

Reduction: Br₂ + 2e⁻ → 2Br⁻

Oxidation: Al → Al³⁺ + 3e⁻

Reducing agent: Al

Oxidising agent: Br₂

Consider the following reaction:

2CrCl₃ + Zn → 2CrCl₂ +

ZnCl₂

State which species

is the oxidising

agent and the reducing agent. Derive the half equations for each redox

agent.

2CrCl₃ + Zn → 2CrCl₂ + ZnCl₂

Reducing agent: Zn

Half equation of reducing agent: Zn →Zn²⁺ + 2e⁻

Oxidising agent: CrCl₃

Half equation of oxidising agent: Cr³⁺ + e⁻ → Cr²⁺

What are the steps when balancing equations using oxidation states?

- Work out the initial oxidation state and the final oxidation state of both

elements

- Work out the change in oxidation state. E.g. N goes from +5 to +4 so the

change is -1. I goes from 0 to +5 so the change is +5

- Work out the number of atoms that react, so you need 5 N and 1 O