Voltaic Cell: A voltaic cell is a type of electrochemical cell, which converts chemical energy to
electrical energy. Chemical energy from movement of electrons, so use redox reactions as they
transfer electrons.
Half Cell: A half-cell contains species present in redox half- equation. If two chemicals in half cell
allowed to mix, uncontrolled electron flow so heat energy released
Setting up Voltaic Cell: Voltaic cell made by connecting two half cells with a wire connected to
voltmeter. Two solutions also connect by salt bridge to allow ions to flow.
– Salt bridge contains electrolyte that doesn’t react with half cells e.g. potassium nitrate
soaked in paper. KCl can’t be salt bridge as chloride ions may react with copper ions to form
CuCl4-.
Metal Half Cells: Metal rod dipped into aqueous metal ion solution. Shown by Zn2+(aq) l Zn(s).
Zn2+ + 2e- <-> Zn.
Ion Half Cells: Solution contains ions of same element in different oxidation states e.g.
Fe3+ (aq) + e- <-> Fe2+(aq). Unreactive platinum electrode used to transport electrons in or
out. Both metal ions must have same concentration.
Electrode Potentials: Direction of electron flow.
– The negative electrode with more reactive metal loses electrons and is oxidised. The
reducing agent.
– The positive electrode with less reactive metal gains electrons and is reduced. The oxidising
agent.
Standard Electrode Potential: E
o
. It is the e.m.f. of a half cell connected to a standard hydrogen
half- cell under standard conditions- 1 mol dm-3, 298K and 100kPa. E
o of standard hydrogen
electrode is 0V.
Standard Hydrogen Half Cell: Half cell contains H2 (g) and solution contains H+ ions at 1moldm3. Inert platinum electrode used to transport electrons. Since a gas, it’s pumped through gas
cylinder. pH is zero in standard hydrogen half cell.
E
o
found in data books. Always shown so that forward reaction is reduction.
Standard electrode potential values… Why is the metal more reactive, explain using electrode
potentials.
– The more negative the Eo
value, the greater the tendency to lose electrons more easily and
undergo oxidation. Metals. So more reactive in losing electrons. Stronger the reducing
agent. System shifts to the left (using equation in table).
– The more positive the Eo
value, the greater the tendency to gain electrons more easily and
underdo reduction. Non-metals. So more reactive in gaining electrons. Stronger the oxidising
agent. System would shift to the right.
– Signs of electrode match up to E cell value.
– If question in terms for oxidising and reducing agent, I need to talk about oxidising and
reducing agents instead.
Why bubbles during reaction? Metal reacts with H+ to form H2 gas.
Calculating Standard Cell Potential or Cell Potential: Cell potential calculated from electrode
potential. Cell potential measured using voltmeter.
E
o
cell = large number – small number. In V units.
Be careful with negatives when one of the Eo missing but given Ecell.
Do not use potassium manganate instead of potassium dichromate in a titration because Ecell
value for MnO4– is more positive than Cl2. So the Cl2 added right at the start will still be present
and will react with the potassium manganate added later on. MnO4– reacts with Cl–
.
Draw electrochemical cell- compete circuit with voltmeter. Salt bridge. Pt electrode in ion
solution- where both ions are same concentration. ‘Ion is 1 moldm-3’.
Charge carriers/ current flow- Electrons through wire. Ions through solution and salt bridge.
Species which reduces… Make sure state correct one e.g. if Ag or Ag+.
If table of standard electrode potentials given, use them to help formulate equation.
If Fe (s), Fe2+ and Fe3+ given, if use Fe (s) produces other two since a continuing reaction. All three
could be involved in reactions
‘Predict the products’- state products and give a reason why using E cell values.