Whilst acid/base reactions were proton (H⁺) reactions, redox reactions are electron reactions. 

Oxidation – loss of electrons

Reduction – gain of electrons

Oxidation States/Numbers

Oxidation state – charge of an atom if all of its bonds were hypothetical

• The number of electrons an atom has gained or lost by forming bonds

• Also involves covalent bonds

• Notation: symbol first, and then number

  • Example: +2, -3

• Oxidation number – given in roman numerals

This is important as the oxidation state of an atom has a significant impact on its chemistry. 

★ Oxidation = hypothetical charge which applies to both ionic and covalent bonds

Charge = actual charge which only applies to ionic bonds ★

Calculating Oxidation States

• Oxidation state of an element = 0

• Oxidation states of a compound sum to zero, and of an ion sum to the ion’s charge

• The more electronegative atom in an ion assumes a negative oxidation state

• The less electronegative atom in an ion assumes a positive oxidation state

• Some rules of thumb: //need to know this//

Oxidation States & Names

Oxidation states are used in the names of compounds

• ‘-ate’ → an element is in a positive oxidation state (usually because it bonded with oxygen)

• ‘-ide’ → an element is in a negative oxidation state

Oxidation state of transition metals is given in Roman numerals.

Examples:

1. FeCl → iron (II) chloride

• Iron (II) means that Fe = +2 (ox. state)

• Chloride means that Cl = neg. ox. state

2. FeCl₃ → iron (III) chloride

• Iron (III) means that Fe = +3 (ox. state)

3. KClO₃ → potassium chlorate

• Chlorate means that Cl = pos. ox. state

Redox Reactions

Whenever an oxidation occurs, a reduction also occurs, hence, redox. 

Examples:

Zn (s) + Cu²⁺ (aq) → Cu (s) + Zn²⁺ (aq)

  0     +2     0       +2

• Zn = oxidised; it loses 2 electrons

  • Zn is the reducing agent because it reduces the copper

• Cu = reduced; it gains 2 electrons

  • Cu²⁺ is the oxidising agent because it oxidises the zinc

Disproportionation * not part of syllabus *

When some atoms of an element are oxidised and others are reduced. 

Example:

2H₂O₂ →2H₂O + O₂

 +1 -1       +1 -2   0

• The oxygen ending in the H₂O gained an e^- and is reduced

• The oxygen ending in the O₂ lost an e^- and is oxidised

Activity Series of Metals

The activity series of metals shows which metals can reduce other metals. The metals at the top of the series can reduce the metals below it, in other words, they oxidise the easiest. //can be found in the data booklet//

Redox & Reactivity

• Redox behaviour is closely linked to reactivity

• The most reactive metals = best reducing agents

• The most reactive nonmetals = best oxidising agents

• The least reactive elements = bad at both, oxidation and reduction

Reactivity & Displacement Reactions

• Reactive metals = better reducing agents than unreactive metals

• As such, a reactive metal can displace less reactive metals from their compounds

Example:

Zn (s) + CuSO₄ (aq) → ZnSO₄ (aq) + Cu (s)

Zn (s) + Cu²⁺ (aq) → Zn²⁺ (aq) + Cu (s)

• As a good reducing agent, the zinc reduces the Cu²⁺, causing it to gain 2 electrons

Half equations/reactions

• Show changes to individual species in redox reactions

Example:

Fe₂O₃ + 2Al → 2Fe + Al₂O₃

Fe³⁺ + 3e^- → Fe ←reduction

Al → Al³⁺ + 3e^- ← oxidation

//can be found in the data booklet//

Balancing Redox Reactions

Step #1: Write the skeletons of the oxidation and reduction half-reactions

Step #2: Balance all elements other than H and O

Step #3: Balance oxygen atoms by adding H₂O molecules where needed

Step #4: Balance hydrogen atoms by adding H⁺ ions where needed

Step #5: Balance the charge by e^-

Step #6: If the # of e^- lost in the oxidation half-reaction is not equal to the # of e^- gained in the reduction half-reaction, multiple one or both of the half-reactions to make the # of e^- gained equal to the # of e^- lost

Step #7: Add the 2 half-reactions: electrons always cancel out; if the same formulas are found on opposite sides, cancel them; if they are on the same side, add them

Step #8: Check that the atoms and charges are balanced

Example: Cr₂O₇^2- (aq) + HNO₂ (aq) → Cr³⁺ (aq) + NO₃^-

1. Cr₂O₇^2- → Cr³⁺

HNO₂ → NO₃^-

2. Cr₂O₇^2- → 2Cr³⁺

HNO₂ → NO₃^-

3. Cr₂O₇^2- → 2Cr³⁺ + 7H₂O

HNO₂ + H₂O → NO₃^-

4. Cr₂O₇^2- + 14H⁺→ 2Cr³⁺ + 7H₂O

HNO₂ + H₂O → NO₃^- + 3H⁺

5. Cr₂O₇^2- + 14H⁺ + 6e^- → 2Cr³⁺ + 7H₂O

HNO₂ + H₂O → NO₃^- + 3H⁺ + 2e^-

6. Cr₂O₇^2- + 14H⁺ + 6e^- → 2Cr³⁺ + 7H₂O

3(HNO₂ + H₂O → NO₃^- + 3H⁺ + 2e^-)

7. Cr₂O₇^2- + 14H⁺ + 3HNO₂ + 3H₂O → 2Cr³⁺ + 7H₂O + 3NO₃^- + 9H⁺

8. Cr₂O₇^2- + 5H⁺ + 3HNO₂ + 3H₂O → 2Cr³⁺ + 4H₂O + 3NO₃^- 

Redox Titration

Redox titrations are done the same way as an acid-base titration. 

CaVan=CbVbn n = molar ratio

Example:

5Fe²⁺ + MnO₄^- + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O

V: 25.0      20.0

C: _____      0.050

∴  C25.05=0.05020.01

      C = 0.200 mol dm^-3

Connection between O₂ and Organic Matter

Organic substances need O₂ to break down. When the organic substance is in water, O₂ comes from the supply of dissolved oxygen in the water. There is only a limited amount of dissolved O₂ as water is polar and O₂ is nonpolar. If the organic material uses O₂, then the aquatic life that is dependent on it [dissolved O₂] suffocate. 

Winkler Method

• Used to measure the amount of dissolved O₂ in water

• High concentration = little to no pollution

• Biochemical oxygen demand – amount of O₂ required to oxidise organic matter in water samples at a definite temperature over a period of 5 days

  • Measured in ppm (parts per million)

  • Measures degree of organic pollution in water sample

Voltaic Cells

• Forces each half of the reaction to take place in a separate container with the electrons moving through a circuit to get from one side to the next

  • exothermic reaction; energy is lost as electrical energy

• Usually involves metals, but does not have to 

Example: reaction of Mg and Cu²⁺ ions:

Mg (s) + Cu²⁺ → Mg²⁺ (aq) + Cu (s) 

• The magnesium reduces the copper ions as it is more reactive

• Energy is released as heat

//need to know how to draw//

• CROA → cathode = reduction; anode = oxidation

  • Cations go towards the cathode → reduction ∴ maintain charge 

  • Anions go towards the anodes → oxidation ∴ balance out charges

★not 100% efficient → Mg loses ability to become Mg²⁺★

• Voltaic cell notation – anode II cathode; solids on the exterior, ions on the interior

★always use K⁺ and NO₃^- for the salt bridge★

• Electron flow – anode to cathode

Key Parts of a Voltaic Cell

Anode:

• electrode/half-cell where oxidation happens

• Contains the more reactive metal

• Negative electrode →produces electrons

Cathode:

• electrode/half-cell where reduction happens

• Contains the least reactive metal

• Positive electrode → accepts electrons

Salt bridge:

• Contains a neutral salt; example: KNO₃

• Made of a tube of jelly or a filter paper soaked in the salt solution

• Ions diffuse in and out to balance the charge and complete the circuit

Voltmeter:

• Measures the difference in potential between half-cells

• Could be replaced by other circuitry to do useful work

★ over time, anode mass decreases (it dissolves) ★

★ over time, cathode mass increases (it grows) ★

Electrolytic Cells

• Opposite of a voltaic cell

• Uses electricity to provide the energy for an endothermic reaction

• Cathode and anode determined by reactivity of the metal

• Current (anode to cathode) is carried by moving ions:

  • Cations (+) move to the cathode (negative electrode)

  • Anode (-) move to the anode (positive electrode)

• Ionic compounds must be either molten or dissolved in a solution, so that the ions can be able to move

• Anode forms a gas and the cathode forms a precipitate

Example: electrolysis of lithium chloride

Li⁺ (l) + Cl^- (l) → Li (s) + ½ Cl₂ (g)

//need to know how to draw//