3.3 Redox

Redox reaction (MnO4-)
Achievement Standard 3.3 - Redox (3 Credits)
"Describe oxidation-reduction processes"

This Achievement Standard has two main parts to it.

  • Oxidation-Reduction reactions
  • Electrochemistry
Oxidation-Reduction reactions

  • Observations of redox reactions
  • Oxidation Numbers (ON)
  • Oxidants and Reductants


-The basics of redox-

Oxidation (LEO)                                     LEO goes GER                        Reduction (GER)                    
-loss of electrons                                                                                     -gain of electrons
-increase in ON                                                                                       -decrease in ON
-reductants get oxidised                                                                           -oxidants get reduced
-reduce other species                                                                               -oxidise other species

-Oxidation Numbers-



  • Elements have ON of 0..... E.g Cl2 = 0
  • Oxygen is -2 except in H2O2 where its is -1...... E.g K2O O = -2
  • Hydrogen is +1 except in metal hydrides where it is -1...... E.g H2O H = +1
  • Monoatomic Ions ON is the charge...... E.g Fe2+, Fe= +2
  • Polyatomic Ions all ON add up to charge...... E.g SO42-, S= +6
  • Molecules/Compounds all ON add up to 0
  • Halides are -1..... E.g HCl, Cl = -1


~Worked Example~
C2O42-, a polyatomic ion, ON must add up to -2
so, the ON of C is +6-8=-2, C=+3

-Half Equations-


-Acid Conditions- (don't forget the states, you will have to learn them!)

Cr2O7 2-(aq)                             → 2Cr3+(aq)  -Balance main atom.
Cr2O7 2-(aq)                             → 2Cr3+(aq) + 7H2O(l) -Balance O by adding H2O to opposite side.
Cr2O7 2-(aq) + 14H+(aq)          → 2Cr3+(aq) + 7H2O(l) -Balance H by adding H+ to opposite side.
Cr2O7 2-(aq) + 14H+(aq) + 6e  → 2Cr3+(aq) + 7H2O(l) - Balance the charge.

-Alkaline and Neutral Conditions-

Just like acidic conditions but instead of adding H+ ions, add OH- to both sides. This creates water (OH- + H+ → H2O).

~Worked Example~

Neutral Conditions:
MnO4-(aq)                             → MnO2(s)                            Balance the main atom (balanced already)
MnO4-(aq)                             → MnO2(s) + 2H2O(l)           Balance the O
MnO4-(aq) + 4H+(aq) +4OH-(aq)→ MnO2(s) + 2H2O(l) +4OH-(aq) Balance the H (refer to alkaline)
MnO4-(aq) + 2H2O(l) +3e → MnO2(s) +4OH-(aq)              Balance the charge etc.

-Oxidants and Reductants-
For a list of oxidants and reductants at Level 3 along with colour changes and half equations, I will upload a picture from a book soon.

That's basically what you need to know for the redox reactions part of this unit.

-Electrochemistry-



  • Electrochemical Cells
  • Cell Diagrams
  • E° Values and E° of Cells
  • Commercial Cells
This is the more challenging part to redox this year as it is completely opposite to the electrolysis section last year.

In level 3, we use redox reactions to create electrical current.

-Electrochemical Cells-

<Will post a picture soon>

Salt Bridge contains an electrolytic (ionic) solution which completes the circuit. It's main function (which is usually asked in the exam) is to balance the build up charge as the redox reactions take place in the half cells. It is usually a solution of the unreactive KNO3. K+ ions move to the negative anode to  balance the charge, where NO3- ions do the same at the cathode.

Standard Conditions
Temperature: 25°C (298K)
Pressure: 101.3 kPa (1 atm)
Concentration: 1 molL-1

-Cell Diagrams-


This is a representation of a cell (instead of drawing it)

E.g <Will post a picture soon>

-E° Values-



  • Standard Reduction Potentials/Electrode Potentials
  • These are values obtained by connecting a half cell to the standard hydrogen half cell
  • The more +ve the E° value, the stronger the oxidant (more easily reduced)
  • The more -ve the E° value, the stronger the reductant (more easily oxidised)
~Worked Example~


Redox Couple                       E°/V
Cl2/Cl-                                 +1.36
SO4 2-/SO2                        +0.20
Mg2+/Mg                            -2.37

The strongest oxidant is Cl2 (Oxidants are on the left of redox couples)
The strongest reductant is Mg (reductants are on the right of redox couples)
The weakest reductant is Cl-
The ion which is the strongest oxidant is SO4 2-

-E° of Cells-


This is the voltage for the entire cell and can be calculated in two ways:


  • E°cell = E°RHE - E°LHE
OR
  • E°cell = E°REDUCTION - E°OXIDATION

If you are given information about a cell (e.g. a picture or an explanation), then it is strongly advise you use the  first equation for calculating the E°cell.



-Commercial Cells-


A commercial cell or battery, is just a cell that is used in everyday life. They follow the exact same process as normal cells, only they appear to be more complex, putting students off.

-The Lead Acid Battery-


  • 6 cells connected in series
  • rechargeable
  • Cell is made up of Lead (Pb) and PbO2
When this battery is discharging, an oxidation reaction occurs at the anode

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