Week 5 - Balancing redox reactions, voltaic and electrolytic cells, primary and secondary batteries

Describe the process of balancing redox reactions in acidic solution

• Write the skeletal form of each half equation

• Balance all atoms except O and H

• Balance O atoms with H₂O molecules

• Balance H atoms with H⁺ atoms

• Add the electrons gained/lost

• Balance both equations so the electrons are equal

• Done

Describe the process of balancing redox reactions in basic solution

• Balance as if the reactions were taking place in acidic solution

• Cancel all H+ using OH^- ions by adding OH^- to both sides of the equation

• The H⁺ and OH^- ions form H₂O molecules so cancel these out with the others

Describe the process of writing molecular equations from skeletal ionic equations

• Balance the ionic equations

• Add the spectator ion to each of the ions present in the ionic equation

Describe the features of a voltaic cell of reactant electrodes and draw one using X⁺ and Y⁺ cations

State the purpose of a salt bridge

• To neutralize the solutions of both the anode and cathode

Draw a diagram of the voltaic cell of 2I^- (to I₂) and MnO₄^- (to Mn²⁺)

Define cell potential (E_cell)

• The difference in electrical potential between half-cells

• Also called voltage or e.m.f

State the what it means E_cell > 0, E_cell < 0 and E_cell = 0

• When E_cell > 0, the reaction is spontaneous

• When E_cell < 0 the reaction is non-spontaneous

• When E_cell = 0 the reaction is in equilibrium.

Define Voltage

• Work done per unit charge in driving that charge around a complete circuit

• 1V = 1 J/C

Define standard cell potential (E°_cell)

• The potential of a cell under standard conditions:

  • 298K

  • 1mol/dm³ concentration of all electrolytes

  • all reactants in their standard states

  • 1 atm for gases

  • pure solids are used for electrodes

Define standard electrode potential (E°_half-cell)

• Standard half cell potential

• The potential of an electrode measured against hydrogen under standard conditions

  • 298K

  • 1mol/dm³ concentration of all electrolytes

  • all reactants in their standard states

  • 1 atm for gases

  • pure solids are used for electrodes

Describe how to find E°_cell of the half-cells from thetheir E°_half-cell

• E°_{cell =} E°_cathode - E°_anode

Describe the trend of half cell potentials

• More positive = more readily receives electrons (Stronger oxidizing agent)

• More negative = more readily loses electrons (Stronger reducing agent)

Define a primary battery

• A non-rechargable battery

• Once the reaction reaches equilibrium it cannot be ‘reversed’ even when supplied with external energy

Name 3 examples of primary batteries

• Dry cells

• Alkaline battery

• Mercury and silver button battery

Define a secondary battery

• It is rechargeable

• The electrochemical reaction is reversible

Define verdigris

• Deterioration of copper by its reaction with the environment

Define rust

• Deterioration of iron by its reaction with the environment

Define tarnish

• Deterioration of Silver by its reaction with the environment