3.1.11 Electrode potentials and electrochemical cells

Define electrochemical series

List of electrode potentials in numerical order

What is the IUPAC convention for writing half-equations?

Write them as reduction reactions

How can these be identified from an electrochemical series

1. Most powerful oxidising agent

2. Weakest oxidising agent

3. Most powerful reducing agent

4. Weakest reducing agent

1. Species on LHS with the most positive electrode potential

2. Species on LHS with the most negative electrode potential

3. Species on RHS with the most negative electrode potential

4. Species on RHS with the most positive electrode potential

How do you calculate the e.m.f of a cell?

E_cell= E^⦵R - E^⦵L= red - ox

Positive right red cat

For spontaneous electrochemical cells, the more positive half-equation goes right and this is reduction and occurs at the cathode.

The more positive half-equation also goes on the RHS in the spontaneous conventional representation

What are the charges of the electrodes?

Cathode = +ve electrode

Anode = -ve electrode

What are the standard conditions?

• 298K

• All gaseous species at 100kPa

• All aqueous species at 1 mol dm^-3

How do you draw a standard hydrogen electrode?

• Hydrogen gas at 1atm

• 1M HCl solution

• Platinum electrode

• 298K

• high-resistance voltmeter

Why is the hydrogen electrode the reference electrode?

It is 0.00V by definition

How do you draw a cell?

• solid metal or platinum electrode (Pt is inert so doesn’t create a potential difference, and is a conductor so allows the transfer of electrons)

• salt bridge to provide an electrical connection from ions that are free to move and carry charge- usually KNO₃ (unreactive as if a reaction occurs, there will be a change in the concentration, affecting the E_cell)

• high-resistance voltmeter: doesn’t allow current to flow so allows emf to be measured

• wire/ammeter: allows current to flow; E_cell eventually falls to 0.00 as the reactants are used up and the cell goes to equilibrium

How do you write the conventional representation of a cell?

• || - salt bridge

• right of salt bridge: positive terminal of voltmeter (half cell with more +ve electrode potential for a spontaneous reaction)

• left of salt bridge: negative terminal of voltmeter (half cell with more -ve electrode potential for a spontaneous reaction) OR standard hydrogen electrode (always on LHS)

• higher oxidation state goes next to salt bridge

• | for phase change, comma for no phase change

What is a galvanic cell?

• has a positive E_cell

• cell is discharging: more positive E is reduced (cathode) and more negative E is oxidised (anode)

• overall equation- flip more negative equation

• the emf is determined when no current flows

• the emf changes when electrodes are connected and voltmeter is replaced with a component that allows current to flow eg wire, bulb, ammeter; current flows so the concentration of the ions change and are no longer standard

What is an electrolytic cell?

• -ve E_cell/ emf

• cell is recharging: more positive E is oxidised and more negative E is reduced

• overall equation occurs in the reverse direction- flip more positive equation

What is the effect on electrode potential if a change in concentration causes equilibrium of the right hand half cell to shift to the RHS?

1. More reduction occurs (less oxidation)

2. More electrons absorbed

3. Electrode becomes more positive

4. Electrode potential becomes more positive

What is the effect on electrode potential if a change in concentration causes equilibrium of the right hand half cell to shift to the LHS?

1. More oxidation occurs (less reduction)

2. More electrons released

3. Electrode becomes more negative

4. Electrode potential becomes more negative

Describe a non-rechargable electrochemical cell and its advantages and disadvantages

• an irreversible cell that can’t be charged by electric current

• reaction not reversible

• E_cell changes over time and eventually falls to 0.00V as the reactants are used up and cell is at equilibrium

Advantages

• cheaper initial cost

• longer storage times

Disadvantages

• cell can’t be reused (higher batter production which depletes supplies of materials used to make the battery so more mining, increases CO₂ production as energy to extract the metal is generated from the burning of fossil fuels; disposal contributes to landfill problems, recycled separately to prevent pollution of the environment by toxic or dangerous substances and to make it easier to recycle valuable components)

• may leak after a long period of use- dry cell specifically (Zn anode is oxidised/ used up)

Describe a rechargeable electrochemical cell and its advantages and disadvantages

• a reversible cell that can be recharged by an electric current

• reaction is reversible

• E_cell changes over time and eventually falls to 0.00V as the reactants are used up

• the cell can be recharged by changing the products back into reactants

Advantages

• the cell can be reused (cheaper in the long run, supplies not depleted so less mining/ less energy required/ less CO₂ released, fewer batteries go into landfill)

Disadvantages

• battery may lose ability to maintain capacity over time

• batteries must be recycled separately to prevent pollution of the environment by toxic or dangerous substances and to make it easier to recycle valuable components

What are the simplified electrode reactions in a lithium cell?

Positive electrode: Li⁺ + CoO₂ + e^- → Li⁺[CoO₂]^-

Negative electrode: Li → Li⁺ + e^-

What do electrochemical cells use instead of a salt bridge?

A porous separator which allows ions to flow

Describe a fuel cell and its advantages and disadvantages

• a cell where electricity is generated from the continuous oxidation of an external fuel source

• overall reaction= combustion

• hydrogen and oxygen supplied continuously (E_cell remains constant as long as H₂ and O₂ are constantly supplied)

Advantages

• can be refuelled quickly compared to the time required to charge an electric battery

• more efficient than combustion of hydrogen in an internal combustion engine and only produces H₂O as the waste product/ does not produce CO₂ (but CO₂ is produced to make hydrogen: CH₄ + 2H₂O → 4H₂ + CO₂)

Disadvantages

• hydrogen is more expensive than petrol

• refuelling is difficult as fewer fuel stations sell hydrogen gas

• hydrogen is explosive so heavy and strong material needed for its storage

• hydrogen may need to be made using an energy source that is not ‘carbon neutral’ (however can use carbon neutral method eg electrolysis of water, generate the electricity using a method that doesn’t produce CO₂ eg solar, wind, geothermal, tidal)

What are the electrode reactions in an alkaline hydrogen-oxygen fuel cell?

Positive: O_2(g) + 2H₂O_(l) + 4e^– → 4OH^– _(aq)

Negative: 4H₂O_(l) + 4e^– → 4OH^– _(aq) + 2H_2(g)

Overall: O_2(g) + 2H_2(g) —> 2H₂O_(l)

What are the electrode reactions in an acidic hydrogen-oxygen fuel cell? Why is the voltage the same for alkaline and acidic?

Positive: O_2(g) + 4H⁺_(aq) + 4e^- → 2H₂O_(l)

Negative: 4H⁺_(aq) + 4e^- → 2H_2(g)

Overall: O_2(g) + 2H_2(g) —> 2H₂O_(l)

The overall reaction is the same