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A-level Chemistry > Electrochemical Cells > Flashcards

Electrochemical Cells Flashcards

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1
Q

What is a half cell?

A

One of two electrodes in a galvanic cell or battery.

A reactant (often metal) is placed in an electrolyte solution.

One half cell is oxidised and the other is reduced in a full working cell.

All half cells are equilibriums

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2
Q

What is a metal solution half cell?

What happens in terms of oxidation and reduction?

A

A half cell where a metal (in one oxidation state) is placed in a solution of its own ions (in another oxidation state)

Eg Cu(s) electrode in Cu^2+ (aq) solution.

Metal is oxidised and ions are reduced forming an equilibrium.

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3
Q

What is the forward reaction of any half cell equation?

A

The forwards reaction is always reduction

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4
Q

What is the is concentrate for a standard half cell solution?

A

1 moldm^-3

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5
Q

How do we construct half cells when an oxidation state contains a gas?

A

We bubble the gas through a liquid and give the reaction a surface to take place on.

Often use a platinum electrode

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6
Q

Why are platinum electrodes used when the half cell does not contain a metal?

A

Can conduct electricity and is used to transfer electrons from one half cell to another.

Inert and hence will not react or affect the redox reaction

Electrode is coated in platinum black, porous substance to increase the surface area for the reaction to take place on.

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7
Q

How do we construct a half cell when both the oxidation states present are in aqueous solution?

A

Construct the cell using an equimolar solution and a platinum electrode to provide or remove electrons from the half cell.

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8
Q

Explain which reaction is favoured when equilibrium lies to the left and right in a half cell equation.

A

Equilibrium lies to the left:
The half cell is better at releasing electrons (being oxidised)

Equilibrium lies to the right:
The half cell is better at accepting electrons (being reduced) since the forwards reaction is always reduction.

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9
Q

What is the standard electrode potential?

A

Used to measure the tendency of a half cell to accept or release electrons

The emf generated by a half cell when it is connected to the standard hydrogen electrode (left) under standard conditions and the electrode in question (right)

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10
Q

What are the standard conditions for a half cell?

A

298K
100KPa
1 moldm^-3 solution

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11
Q

What is the hydrogen electrode and how is it constructed?

A

Used to compare the how different metals release or accept electrons

Hydrogen gas is bubbles into a solution of H+ ions. Since hydrogen does not conduct electricity, electrical contact is made using a platinum electrode.

Hydrogen half cell has an electrode potential of 0v

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12
Q

How is an electrical circuit created from half cells?

A

The 2 half cells are connected via a salt bridge (allows free ions to move completing the circuit) an have a volt meter attached to record the electrical current between the half cells.

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13
Q

Why is a voltmeter (high resistance) used?

A

To tell us the difference in standard potentials of the half cells measured in volts

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14
Q

What is the purpose of the salt bridge?

A

Contains free ions to complete the circuit and allow the charge to flow

It redistributes the ion charges

Usually a piece of filter paper soaked in potassium of ammonium nitrate.

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15
Q

In a full cell which half cell is the positive and which is the negative terminal?

A

The half cell with the more negative E potential is the negative terminal as it will more readily donate electrons.

The more positive e potential half cell more readily accept electrons and is the positive terminal

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16
Q

What is an electrochemical series?

A

A lift of standard electrode potentials

17
Q

Which half cells are better oxidising agents?

A

The more positive an E potential, the more the forward reaction is favoured (reduction) and hence will have the strongest oxidising agents.

18
Q

In an Electrochemical series which side of the equation are the reducing agents, and which are the oxidising agents.
Also which species are the most powerful oxidising and reducing agents?

A

LHS - Oxidising agents (positive = stronger)
RHS - Reducing agents (negative = stronger)

The more negative an electrode potential, the more favoured the backwards reaction is and hence oxidation takes place.

This means more negative are stronger reducing agents

19
Q

What is the e cell?

A

The difference between the positive electrode potential and negative electrode potential.

Ecell = E positive terminal - E negative terminal

20
Q

What is the short hand way of writing a half cell for Zn and Zn^2+?

A

Zn^2+(aq) | Zn(s)

E potential = -0.76v

represents a phase boundary

21
Q

How do you write a half cell when there is one phase?

Fe^2+ and Fe^3+

A

Pt (s) | Fe^3+(aq), Fe^2+(aq)

22
Q

How do you write the hydrogen half cell?

A

Pt(s) | 1/2 H2(g) | H+(aq)

E potential = 0.00v

23
Q

How do you write a combine half cell between Zn and Cu?

A

Zn(s) | Zn^2+(aq) : : Cu^2+(aq) | Cu(s)

E potential = 1.10v

: : represents a salt bridge

24
Q

What are the rules for writing half cells and full cells?

A

: : salt bridge

Reduced species on the outside

Positive electrode on the right

phase boundary

25
Q

What is the anti-clockwise rule?

A

When working out which half cell favours which reaction,

Write the more negative on top of the more positive half cell equation

Then the reaction occurs in the anti clockwise direction

Hence the top half cell will favour the backward reaction and Bottom half cell favours the forward reaction.

26
Q

How can we tell the feasibility of a reaction?

A

Cell potentials tells us the extent at which electrons flow between the half cells.

The greater the cell potential, the more likely the reaction is feasible.

If the E cell is less than 0.4v then the reaction is unlikely to take place

27
Q

What limits how accurate a cell calculation is?

A

If the standard conditions for the half cells are not met then equilibrium will move and the potentials of the half cells will be different

28
Q

What are the different types of cell? And give examples of each

A

Non rechargeable (primary cells):
Zinc/copper cells used in telegraphs
Zinc/carbon cells used in disposable batteries

Rechargeable (secondary cells):
Lead acid cells used in cars
Lithium ion used in laptops and phones

Fuel cells:
Alkaline hydrogen-oxygen fuel cell

29
Q

What are the electrodes and their equations in a lithium ion cell?

A

Positive electrode:
Li+ CoO2 + e- ——> Li+[CoO2]-

Negative electrode:
Li ——> Li+ + e-

30
Q

How do fuel cells work?

A

Used to create an electrical current and do not need to be recharged as long as they are supplied with their reactants.

31
Q

What are the reactions that take place in an alkaline hydrogen -oxygen fuel cell?

A

Hydrogen and oxygen flow into the cell and react and water is produced which comes out.

2H2(g) +4OH-(aq) ——> 4H20(l) + 4e-
E potential = -0.83v

O2(g) +2H2O(l) + 4e- ——> 4OH-(aq)
E potential = +o.40v

Overall:
2H2(g) + O2(g) ——> 2H2O(l)
E cell = 1.23v

32
Q

What are the benefits and risks of using hydrogen fuel cells?

A

Benefits:
Only waste product is water
Do not need to be recharged
Will continue to produce electricity

Risks:
Hydrogen highly flammable and so could have issues with transporting and storing
Tricky to store gases
Storing gas takes up a lot of space
Produced from burning of fossil fuels (non renewable source)

33
Q

Why may some reactions not occur?

A

Reactions may be too slow

Activation energy may be too high

34
Q

In an Electrochemical series, how do we know which species will oxidise others?

A

The species with the more positive Electrode potential can oxidise those with a more negative electrode potential.

35
Q

Explain how a battery is recharged?

A
  1. current is applied
  2. electrons flow in opposite directions
  3. reverse reactions take place
  4. write equation showing reverse reactions
36
Q

Explain how the hydrogen oxygen fuel cell works.

A

Hydrogen is fed into the negative electrode and oxygen is fed into the positive electrode.
both electrodes are made of platinum.

At the negative electrode:
2OH- + H2 —-> 2H2O + 2e-
electrons flow through the external circuit and the OH- ions are attracted to the oxygen electrode.

At the positive electrode:
O2 + 2H2O + 4e- —-> 4OH-

Overall:
2H2 + O2 —-> 2H2O

A-level Chemistry (33 decks)

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