1. Half Cells
• An electrode can be either positive or negative
• 2 electrodes/half cells are joined together to make a 6. Electrochemical Cells
complete circuit. • Zinc/Copper Cells
• Measuring the electrical potential difference of the cell • Developed in 1830s
indicates how readily electrons are released by one of the • Not practical for portable devices
electrodes • Electrons are transferred from more reactive metal to the
• The circuit is completed by a salt bridge (generally other
saturated potassium nitrate is used).
allows the flow of ions • Zinc/Carbon Cells
doesn’t impact the balance of the circuit. • Carbon is the positive electrode (cathode)
• this is all done under standard conditions • Basis for most disposable battery cells
Temperature: 298K • Electrolyte is a paste rather than a liquid
Pressure: 100KPa • Commercial form has a zinc cannister with aluminium chloride paste
Concentration: 1.00mol/dm3 • Hydrogen gas is oxidised to water by manganese oxide – stops pressure
• The Zinc electrode is the anode and electrons are lost build up
from here, the Zn electrode will decrease in size but the Note: • As cell discharges, Zn is used up, walls become thin and prone to leakage
concentration of Zn2+ will increase. Metals react by losing electrons • NH4Cl is corrosive
• The electrons flow through the wire to the copper and Non-metals react by gaining • Used as doorbells – small current needed
react with the Cu2+ ions and make Cu. Causing the electrons
concentration of Cu2+ to decrease and the size of the Cu • Alkali Batteries
electrode to increase. • Based on same system
• Electrolyte is potassium hydroxide instead of aluminium chloride
2. Hydrogen Electrode • Powdered zinc is used – larger s.a allows higher currents
Physical Chemistry: Electrochemical Cells • Cell is in a steel container to prevent leakages
• Used to compare the tendency of metals to release
electrons. • Used for stereos
• Standard Hydrogen Electrode has electrode potential of
0.00v 3. Electrochemical Series 5. Representing Cells
• Hydrogen gas is bubbled into a solution of H+ ions (most • Written as a table of half equations with electrode
likely from HCl). potentials on the side.
• Hydrogen doesn’t conduct electricity, so a platinum • Each equation can be thought about in terms of
electrode is used instead. equilibria • The singular vertical line indicates a boundary phase
unreactive/inert • The most negative value means a better reducing • The double vertical line indicates a salt bridge
conducts electricity agent, so the oxidation reaction is favoured. Equilibria • Need to include the states of each compound
• The platinum is coated finely on the surface to increase the lies to the left as the backwards reaction is favoured • If using Fe2+ and Fe3+, they will both be aqueous, depending on
surface area and allow the reaction to occur rapidly. • The most positive value means a better oxidizing agent, if they are oxidation or reduction will determine the order, but
• This is done under standard conditions. so the reduction reaction is favoured. Equilibria lies to they will use a platinum electrode and be separated by a
the right as the forwards reaction is favoured comma because they are in the same state
• The number of electrons involved in the reaction has 4. Calculating E.M.F
no impact on the value • E.M.F=Reduction – Oxidation
• The Oxidation occurs at the anode (negative electrode) • In the conventional diagram above, the right side is always
• The Reduction occurs at the cathode (positive reduction
electrode) • The left side is oxidation
• The value given in the table is the same for either the • The value always needs to be positive otherwise the reaction isn’t
reduction or oxidation reaction, it just depends on feasible
what other element is being used
• An electrode can be either positive or negative
• 2 electrodes/half cells are joined together to make a 6. Electrochemical Cells
complete circuit. • Zinc/Copper Cells
• Measuring the electrical potential difference of the cell • Developed in 1830s
indicates how readily electrons are released by one of the • Not practical for portable devices
electrodes • Electrons are transferred from more reactive metal to the
• The circuit is completed by a salt bridge (generally other
saturated potassium nitrate is used).
allows the flow of ions • Zinc/Carbon Cells
doesn’t impact the balance of the circuit. • Carbon is the positive electrode (cathode)
• this is all done under standard conditions • Basis for most disposable battery cells
Temperature: 298K • Electrolyte is a paste rather than a liquid
Pressure: 100KPa • Commercial form has a zinc cannister with aluminium chloride paste
Concentration: 1.00mol/dm3 • Hydrogen gas is oxidised to water by manganese oxide – stops pressure
• The Zinc electrode is the anode and electrons are lost build up
from here, the Zn electrode will decrease in size but the Note: • As cell discharges, Zn is used up, walls become thin and prone to leakage
concentration of Zn2+ will increase. Metals react by losing electrons • NH4Cl is corrosive
• The electrons flow through the wire to the copper and Non-metals react by gaining • Used as doorbells – small current needed
react with the Cu2+ ions and make Cu. Causing the electrons
concentration of Cu2+ to decrease and the size of the Cu • Alkali Batteries
electrode to increase. • Based on same system
• Electrolyte is potassium hydroxide instead of aluminium chloride
2. Hydrogen Electrode • Powdered zinc is used – larger s.a allows higher currents
Physical Chemistry: Electrochemical Cells • Cell is in a steel container to prevent leakages
• Used to compare the tendency of metals to release
electrons. • Used for stereos
• Standard Hydrogen Electrode has electrode potential of
0.00v 3. Electrochemical Series 5. Representing Cells
• Hydrogen gas is bubbled into a solution of H+ ions (most • Written as a table of half equations with electrode
likely from HCl). potentials on the side.
• Hydrogen doesn’t conduct electricity, so a platinum • Each equation can be thought about in terms of
electrode is used instead. equilibria • The singular vertical line indicates a boundary phase
unreactive/inert • The most negative value means a better reducing • The double vertical line indicates a salt bridge
conducts electricity agent, so the oxidation reaction is favoured. Equilibria • Need to include the states of each compound
• The platinum is coated finely on the surface to increase the lies to the left as the backwards reaction is favoured • If using Fe2+ and Fe3+, they will both be aqueous, depending on
surface area and allow the reaction to occur rapidly. • The most positive value means a better oxidizing agent, if they are oxidation or reduction will determine the order, but
• This is done under standard conditions. so the reduction reaction is favoured. Equilibria lies to they will use a platinum electrode and be separated by a
the right as the forwards reaction is favoured comma because they are in the same state
• The number of electrons involved in the reaction has 4. Calculating E.M.F
no impact on the value • E.M.F=Reduction – Oxidation
• The Oxidation occurs at the anode (negative electrode) • In the conventional diagram above, the right side is always
• The Reduction occurs at the cathode (positive reduction
electrode) • The left side is oxidation
• The value given in the table is the same for either the • The value always needs to be positive otherwise the reaction isn’t
reduction or oxidation reaction, it just depends on feasible
what other element is being used
