Unit 3.1.11 - Electrode potentials and electrochemical cells
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Summary Electrode Potentials and Electrochemical Cells Revision Notes
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Unit 3.1.11 - Electrode potentials and electrochemical cells
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AQA
Book
AQA Chemistry A Level Year 2 Student Book
Here are my revision notes for the Electrode Potentials and Electrochemical Cells section of the course.
Feel free to check out my other computer typed notes for parts of the 2nd year course on here.
I also have hand written notes of matching quality for pretty much the rest of the course (1st ...
Unit 3.1.11 - Electrode potentials and electrochemical cells
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krishanduggal
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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
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