DM - Electrochemical cells

Question 1

What do you get when electrons move?

Answer 1

Electricity

Question 2

What do electrochemical cells make?

Answer 2

Electricity

Question 3

What can electrochemical cells be made from?

Answer 3

Two different metals dipped in salt solutions of their own ions and connected by a wire (the external circuit).

Question 4

What do two different metals dipped in salt solutions of their own ions and connected by a wire (the external circuit) make up?

Answer 4

Electrochemical cells

Question 5

There are always how many reactions within an electrochemical cell?

Answer 5

2

Question 6

What two reactions are there always within an electrochemical cell?

Answer 6

Oxidation

Reduction

Question 7

What type of process is going on in an electrochemical cell?

Answer 7

A redox process.

Question 8

Where in an electrochemical cell does oxidation always happen?

Answer 8

At the anode (the positive electrode).

Question 9

Where in an electrochemical cell does reduction always happen?

Answer 9

At the cathode (the negative electrode).

Question 10

Which electrode is positive?

Answer 10

Anode.

Question 11

Which electrode is negative?

Answer 11

Cathode.

Question 12

Is the anode positive or negative?

Answer 12

Positive.

Question 13

Is the cathode positive or negative?

Answer 13

Negative.

Question 14

What happens in a zinc/copper electrochemical cell?

Answer 14

1) Zinc loses electrons more easily than copper. So in the zinc electrode half-cell, zinc (from the zinc electrode) is OXIDISED to form Zn^2+ (aq) ions.
Zn (s) –> Zn^2+ (aq) + 2e-
This releases electrons into the external circuit.

2) In the other half-cell, the same number of electrons are taken from the external circuit, REDUCING the Cu^2+ ions to copper atoms.
Cu^2+ (aq) + 2e- –> Cu (s)

Question 15

Where do electrons flow in an electrochemical cell?

Answer 15

Through the wire from the most reactive metal to the least.

Question 16

How are the solutions in the two half-cells connected?

Answer 16

By a salt bridge made from filter paper soaked in a salt solution, e.g. KNO3 (aq).

Question 17

What is the salt bridge for?

Answer 17

The salt ions flow through the salt bridge to complete the cell, and balance out the charges in the beakers.

Question 18

What is included in the external circuit?

Answer 18

A voltmeter.

Question 19

What does the voltmeter do in the external circuit?

Answer 19

Shows the voltage between the two half-cells.

Also measures the direction of the flow of electrons.

Question 20

What is the cell potential or EMF, Ecell?

Answer 20

The voltage between the two half-cells.

Question 21

What is the voltage between the two half-cells called?

Answer 21

The cell potential or EMF, Ecell.

Question 22

What can half-cells involve?

Answer 22

Solutions of two aqueous ions of the same element, such as Fe^2+ (aq)/Fe^3+ (aq).

Question 23

Where does the conversion from Fe^2+ to Fe^3+ (or vice versa) happen in an electrochemical cell?

Answer 23

On the surface of the electrode.

Question 24

Equations showing conversion of Fe^2+ to Fe^3+ (and vice versa)

Answer 24

Fe^2+ (aq) –> Fe^3+ (aq) + e-

Fe^3+ (aq) + e—> Fe^2+ (aq)

Question 25

What is important about the electrode if the half-cell involves solutions of two aqueous ions of the same element?

Answer 25

Because neither the reactants nor the products are solids, you need something else for the electrode. It needs to conduct electricity and be very inert, so that it won’t react with anything in the half-cell.

Platinum is an excellent choice, but is very expensive, so graphite is often used instead.

Question 26

What do the electrodes need to be able to do?

Answer 26

Conduct electricity and be very inert, so that it won’t react with anything in the half-cell.

Question 27

What are the electrodes usually made out of if neither the reactants nor the products are solids?

Answer 27

Platinum is an excellent choice, but is very expensive, so graphite is often used instead.

Question 28

Are electrochemical cells always made out of metals?

Explain

Answer 28

Can also be made from non-metals. For systems involving a gas (e.g. chlorine), the gas can be bubbles over a platinum electrode sitting in a solution of its aqueous ions (e.g. Cl-).

Question 29

When drawing electrochemical cells, which half-cell is drawn on which side?

Answer 29

The half-cell where oxidation happens (the anode) should always be drawn on the left, and the half-cell where reduction happens (the cathode) is drawn on the right.

Question 30

Explain the method to set up an electrochemical cell involving two metals

Answer 30

1) Get a strip of each of the metals you’re investigating. These are your electrodes. Clean the surfaces of the metals using a piece of emery paper (or sandpaper).
2) Clean any grease or oil from the electrodes using some propanone. From now, be careful not to touch the surfaces of the metals with your hands as you could transfer grease back onto the strips.
3) Place each electrode into a beaker filled with a solution containing ions of that metal. E.g. if you had a zinc electrode, you could place it in a beaker of ZnSO4 (aq). Sometimes you’ll have to add acid to the half-cell too.
4) Create a salt bridge to link the two solutions together by simply soaking a piece of filter paper in salt solution, e.g. KCl (aq) or KNO3 (aq) and draping it between the two beakers. The ends of the filter paper should be immersed in the solutions.
5) Connect the electrodes to a voltmeter, using crocodile clips and wires. You should get a reading on your voltmeter if this is done right.

Question 31

How will your method be different if your electrochemical cell is made up of half-cells where neither the oxidised or reduced species are solid (e.g. their both aqueous ions)?

Answer 31

For example, you’ll need to use an inert electrode (e.g. platinum).

Question 32

Are the reactions at each electrode reversible or irreversible?

Answer 32

Reversible

Question 33

What does the direction in which each reaction goes in depend on?

Answer 33

How easily each metal lose electrons (i.e. how easily it’s oxidised).

Question 34

What does each half-cell have?

Answer 34

An electrode potential.

Question 35

What is an electrode potential?

Answer 35

The potential difference between the electrode and the solution.

Question 36

Which way does the value for the half-reaction with the more positive standard electrode potential go?

Answer 36

Forwards (more positive).

Question 37

Which way does the value for the half-reaction with the more negative standard electrode potential go?

Answer 37

Backwards (more negative).

Question 38

Symbol for standard electrode potential

Answer 38

E°

Question 39

What are the standard conditions?

Answer 39

298K and 100kPa

Question 40

In a cell, does the half-reaction in the half-cell with the more negative electrode potential go in the direction of oxidation or reduction?

Answer 40

Oxidation (backwards).

Question 41

In a cell, does the half-reaction in the half-cell with the more positive electrode potential go in the direction of oxidation or reduction?

Answer 41

Reduction (forwards).