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A Level Physical Chemistry > 1.11 Electrode Potentials and Cells > Flashcards

1.11 Electrode Potentials and Cells Flashcards

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1
Q

What happens when a rod of a metal is dipped into a solution of its own ions?

A

An equilibrium is set up between the solid metal and the aqueous metal ions

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2
Q

Write a half-equation for zinc (s) to zinc (II)

A

Zn(s) ⇌ Zn2^+ (aq) + 2e-

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3
Q

Write a half-equation for copper (II) to copper (III)

A

Cu2+(aq) ⇌ Cu3+ (aq) + e-

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4
Q

What is the simplest salt bridge made of?

A

Filter paper soaked in saturated solution of KNO3 (potassium nitrate)

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5
Q

Why are salt bridges necessary?

A

Complete the circuit, but avoid further metal/ion potentials as does not perform electrochemistry. Allow ion movement to balance the change. Do not react with electrodes

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6
Q

What symbol is used to represent a salt bridge in standard notation?

A

||

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7
Q

What type of species goes on the outside (furthest from the salt bridge) in standard cell notation

A

The most reduced species

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8
Q

What does | indicate?

A

Phase boundary (solid/liquid/gas)

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9
Q

How would an Aluminium/copper cell be represented?

A

Al(s) | Al3+ (aq) || Cu2+ (aq) | Cu(s)

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10
Q

What happens at the left hand electrode?

A

Left hand electrode is where oxidation occurs. Left hand electrode is the half cell with the most negative E^0 value

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11
Q

What happens at the right-hand electrode?

A

Right hand electrode is where reduction occurs. Right hand electrode is the half cell with the most positive E^0 value

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12
Q

Which side of the cell has the most negative E0 value? What happens to the metal with the most negative E^0 value?

A

Oxidation- left hand electrode

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13
Q

Draw the standard hydrogen electrode

A
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14
Q

What conditions is the standard hydrogen electrode used in?

A

Temp = 298K
Pressure = 100kPa
[H+] = 1.00 mol dm^-3

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15
Q

Why might you use other standard electrodes occasionally?

A

They are cheaper/easier/quicker to use and can provide just as good a reference

Platinum is expensive

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16
Q

If an E^0 value is more negative, what does it mean in terms of oxidising/reducing power?

A

Better reducing agent (easier to oxidise)

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17
Q

If an E^0 value is more positive, What does it mean in terms of oxidising/reducing power?

A

Better oxidising agent (easier to reduce)

18
Q

What factor will change E^0 values?

A

Concentration of ions
Temperature

19
Q

What happens if you reduce the concentration of the ions in the left hand half cell?

A

Equilibrium moves to the left to oppose the change of removing ions; this releases more electrons, the E^0 of the left hand cell becomes more negative, so the e.m.f. Of the cell increases

20
Q

How do you calculate the emf of a cell from E^0 values

A

E^0cell = E^0right - E^0left

21
Q

When would you use a platinum electrode?

A

When both the oxidised and reduced forms of the metal are in aqueous solution

22
Q

Why is platinum chosen

A

Inert so does not take part in the electrochemistry

Good conductor to complete circuit

23
Q

How would you predict if a reaction would occur?

A

Take the 2 half equations

Find the species that is which reduced(this is effectively the right hand electrode)

Calculate its E^0 value minus the E^0 value of the species that is being oxidised (effectively the left hand cell)

If E^0 overall > 0, reaction will occur

24
Q

What was the first commercial cell made from (Daniell cell)

A

Zinc/copper (II)

25
Q

What are zinc/carbon cells more commonly known as?

A

Disposable batteries

26
Q

What are the two reactions that take place in zinc / carbon cells?

A

Zn oxidised to Zn2+

NH4+ reduced to NH3 at carbon electrode

27
Q

What are the reactions that occur in a lead/acid battery (car batteries)?

A

Pb + SO4 2- —> PbSO4 (S) + 2e-

PbO2 + 4H+ + SO4 2- + 2e- —> PbSO4 + 2H2O

28
Q

How are cells recharged (if they are rechargeable)

A

Reactions are reversible and are reversed by running a higher voltage through the cell than the cells E^0

29
Q

Nickel/cadmium cells are rechargeable AA batteries etc. what reactions occur at the electrodes?

A

Cd(OH)2 (s) + 2e- —> Cd(s) + 2OH-

NiO(OH) (s) + H2O + e- —> Ni(OH)2 (s) + OH-

30
Q

Where are lithium-ion cells used?

A

Mobile phones

Laptops

31
Q

What reactions occur on discharge in lithium-ion cells?

A

Li+ + CoO2 + e- —> Li+[CoO2]-

Li—> Li+ + e-

32
Q

What is a fuel cell?

A

A cell that is used to generate electric current; doesn’t require electrical recharging

33
Q

what are the reactions that take place at the two electrons in an alkaline hydrogen fuel cell?

A

2H2 + 4OH- —> 4H2O + 4e-

O2 + 2H2O + 4e- —> 4OH-

34
Q

Draw a diagram of a hydrogen fuel cell

A
35
Q

Why is it better to use a fuel cell than to burn H2 in air, even though the same overall reaction occurs

A

In combustion, sulfer containing compounds (SO2, SO3) and nitrogen containing compounds (NO2, NOx) are produced due to the high temperatures and the S and N in the air. These are bad for the environment.
This does not occur in fuel cell; the only product is water.
More efficient

36
Q

Disadvantages of fuel cells?

A

Hydrogen is a flammable gas with a low b.p. —> hard and dangerous to store and transport —> expensive to buy

Fuel cells have a limited lifetime to use toxic chemicals in their manufacture

37
Q

How do you find the weakest reducing agent from a table of electrode potential date?

A

Most positive E^ θ value. Then it is the PRODUCT of the reduction equation I.e. imagine equation going from right to left

38
Q

What is the reason that some cells cannot be recharged?

A

Reaction of the cell is not reversible - a product is produced that either dissipates or cannot be converted back into the reactants

39
Q

Why might the e.m.f of a cell change after a period of time?

A

Concentration of the ions change - the reagents are used up

40
Q

How can the e.m.f. Of a cell be kept constant

A

Reagents are supplied constantly, so the concentrations of the ions are constant; E^0 remains constant

A Level Physical Chemistry (12 decks)

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