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Chemistry: CHEM 5 > Redox Equilibria > Flashcards

Redox Equilibria Flashcards

1
Q

What is a reducing agent?

A

A species that donates electrons.

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2
Q

What is an oxidising agent?

A

A species that accepts electrons.

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3
Q

What happens in reduction?

A

Electrons are gained and oxidation state decreases.

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4
Q

What happens in oxidation?

A

Electrons are lost and oxidation state increases.

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5
Q

What is an electrochemical cell?

A

Electrochemical cells can be made from two different metals dipped in salt solutions of their own ions and connected by a wire (the external circuit). There are always a reduction and oxidation reaction taking place.

  • The negative electrode is always on the left, it contains the species that loses electrons more easily, it is the better reducing agent, and is oxidised itself. It releases electrons into the external circuit.
  • The positive electrode is always on the right, it contains the species that gains electrons more easily, it is the better oxidising agent, and is reduced itself. It accepts electrons from the external circuit.
  • ELECTRONS FLOW FROM THE MOST REACTIVE TO THE LEAST.
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6
Q

How is a zinc-copper electrochemical cell set up?

A

-A zinc electrode is placed in a solution of Zn2+(aq) (1M conc ZnSO4). This is a half cell and can be represented as Zn2+(aq) | Zn(s).
-A copper electrode is placed in a solution of Cu2+(aq) (1M conc CuSO4). This half cell can be represented as Cu2+(aq) | Cu(s).
-Both half cells are known as the redox couple.
-The whole electrochemical cell can be represented as:
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s).

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7
Q

What is the salt bridge?

A

The electrochemical cell will not work until the circuit is completed, by the salt bridge.
-It is usually a strip of filter paper soaked in a solution of a soluble ionic compound, usually KNO3(aq). This allows ions to flow through and balance out the charges.

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8
Q

What is cell potential?

A

Each electrode has an electrode potential. The potential difference between the two half cells is known as the cell potential, or the electromotive force (e.m.f) or Ecell.

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9
Q

How is electrode potential calculated?

A

Electrode potential is the e.m.f of the half-cell.

-This is calculated experimentally by measuring the half cell relative to the standard hydrogen electrode (SHE).

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10
Q

What is the standard electrode potential, Eᶿ?

A

What is the standard electrode potential, Eᶿ? The standard electrode potential, Eᶿ, of a half ell is the e.m.f of the half cell, relative to the standard hydrogen electrode measured when the concentrations of all the ions are 1.0 moldm-3, the temperature is 298K and any gases are at 100kPa pressure.

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11
Q

What is the standard hydrogen electrode?

A

The standard hydrogen electrode consists of pure hydrogen gas (at 100kPA and 298K) bubbled across a platinum electrode, placed in a 1M solution of acid (1.00moldm-3).

  • The half cell can be represented as: Pt(s) | H2(g) | H+(g).
  • The potential difference of the standard hydrogen electrode is defined as 0.00v.
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12
Q

What changes would increase Eᶿ of the standard hydrogen electrode above 0.00V?

A

2H+(aq) + 2e- H2(g)

  • Increase [H+(aq) abouve 1.00moldm-3. This would shift the equilibrium to the right to oppose the disturbance, thereby increasing the value of Eᶿ.
  • Decrease the pressure of H2(g) below 100kPa. This would shift the equilibrium to the right to oppose the disturbance, thereby increasing the value of Eᶿ.
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13
Q

What is an alternative to the standard hydrogen electrode?

A

The Calomel electrode.
-It is made of a platinum electrode, placed in liquid mercury in contact with a mixture of mercury (I) chloride and potassium chloride.
-Pt(s) | Hg(l) | Hg2CL2(sat) ∙ Kcl.
Eᶿ = +0.27V (against SHE).
-It can potentially be dangerous as the vapours of mercury are poisonous ad can affect the brain.

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14
Q

What is an electrochemical series?

A

A list of electrode potentials written in order from most negative to most positive.

  • The half equations are always written as reduction reactions.
  • When two half-equations are put together in an electrochemical cell, the more negative is oxidised, the more positive is reduced.
  • Top: more reducing, are oxidised, lose e-.
  • Bottom: more oxidising, are reduced, gain e-.
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15
Q

How do you calculate cell potential/e.m.f/Eᶿ?

A 
Eᶿcell = EᶿRHS - EᶿLHS
Eᶿcell = Eᶿred - Eᶿox
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16
Q

How can you decide whether the reactions is feasible using Eᶿcell?

A

If Eᶿcell is positive overall, the reaction is feasible.

17
Q

What are non-rechargeable cells?

A

Non-rechargeable cells use irreversible reactions; a common type is a dry cell alkaline battery. They are useful for gadgets that don’t use a lot of power or are only used for short periods of time e.g. TV remote, torch.

  • Zinc/carbon cells
  • Zinc chloride cells
  • Long life alkaline batteries
18
Q

What is a zinc/carbon cells?

A

A non-rechargeable cell.

  • Zn(s) | Zn2+(aq) || MnO2(s) Z Mn2O3(s)
  • e.m.f. = 1.5V
  • The positive electrode is made of carbon, which acts like the inert platinum electrode in the hydrogen electrode.
  • The commercial form of the cell consists of a zinc canister, filled with a paste of ammonium chloride and water. This acts as the electrolyte.
  • In the centre is a carbon rod. It is surrounded by a mixture of manganese (IV) oxide and powdered carbon.
19
Q

What reactions occur in a zinc/carbon cell?

A

At the zinc: Zn(s) –> Zn2+(aq) + 2e-
At the carbon rod: 2NH4+(aq) + 2e- –> 2NH3(g) + H2(g)
-Overall: 2NH4+(aq) + Zn(s) –> 2NH3(g) + H2(g) + Zn2+(aq)

20
Q

What happens in a zinc/carbon cell?

A
  • The hydrogen gas is oxidised to water by the manganese (IV) oxide (preventing a build-up of pressure), while the ammonia dissolves in the water of the paste.
  • As the cell discharges, the zinc is used up and the walls of the zinc canister become thin and prone to leakage. The ammonium chloride electrolyte is acidic and can be corrosive.
21
Q

What is a zinc chloride cell?

A

A zinc chloride cell is a variant to zinc/carbon cells (non-rechargeable cell).

  • Zinc chloride is used as the electrolyte.
  • Such batteries are better at supplying high currents and are marked as ‘extra life’ batteries (for radios, torches, shavers).
22
Q

What are long life alkaline batteries?

A

Rechargeable cells can be recharged by reversing the cell reactions. This is done by applying an external voltage to drive the electrons in the opposite direction. This is possible because none of the substances in a rechargeable battery escape or are used up.

  • Lead-acid batteries
  • Nickel/cadmium
  • Lithium ion
23
Q

What are rechargeable cells?

A

Rechargeable cells can be recharged by reversing the cell reactions. This is done by applying an external voltage to drive the electrons in the opposite direction. This is possible because none of the substances in a rechargeable battery escape or are used up.

  • Lead-acid batteries
  • Nickel/cadmium
  • Lithium ion
24
Q

What are lead-acid batteries?

A

They are rechargeable batteries used to operate the starter motor of cars. They consist of six 2V cells connected in a series to give 12V. Each cell consists of two plates dipped into a solution of sulfuric acid (electrolyte).

  • Pb(s) | PbSO4(s) || PbO2(s) | PbSO4(s)
  • e.m.f. = 2V
  • The positive plate is made out of lead coated with lead (IV) oxide, PbO2. This is a lead (IV) dioxide cathode.
  • The negative plate is made out of lead.
  • Both electrodes end up coated in lead (II) sulfate.
25
Q

What reactions occur in a lead-acid battery?

A

On discharging, the battery drives the electrons from the lead plate to the lead (IV) dioxide plate.

  • At the lead plate: Pb(s) + SO42-(aq)  PbSO4(s) + 2e-
  • At the lead (IV) dioxide plate: PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- –> PbSO4(s) + 2H2O(l)
  • Overall: PbO2(s) + 4H+(aq) +2SO42-(aq) + Pb(s) –> 2PbSO4(s) + 2H2O(l)
  • These reactions are reversed as the battery is charged up and electrons flow in the reverse direction, driven by the car’s generator.
26
Q

What is a nickel/cadmium battery?

A

-A rechargeable battery that can be recharged up to 500 times and contains an alkaline electrolyte.
-e.m.f. = 1.2V
-Cd(OH)2(s) + 2e- Cd(s) + 2OH-(aq)
-NiO(OH)(s) + H2O(l) +e- Ni(OH)2(s) + OH-(aq)
Overall: 2NiO9OH)(s) + Cd(s) + 2H2O(l) 2Ni(OH)2(s) + Cd(OH)2(s)
The reaction goes from left to right on discharge (from Cd to NI) and right to left on charging.

27
Q

What is a lithium ion battery?

A

A rechargeable battery that is light because they are based on lithium rather than heavier metals.

  • They have polymer electrolytes rather than liquids or pastes, meanng they cannot leak.
  • A single cell gives a voltage of 3.5-4.0V, compared with 1.5V for a zinc/carbon cell.
  • The cell reactions are complex, involving lithium cobalt oxide.
28
Q

What are the arguments for non-rechargeable vs rechargeable batteries?

A
  • Cost: NR are cheaper however they have to be replaced everytime so R are cheaper in the long term.
  • Lifetime: NR normally works for longer but R can be reused
  • Power: NR can’t supply as much power, no use in mobile phones/laptops.
  • Uses of resources and waste: More NR are produced, can only be used once, so also create more waste. (Both can be recycled).
  • Toxicity: NR less likely to contain lead and cadmium, so less hazardous in landfill.
29
Q

What is a fuel cell?

A

In a fuel cell, the chemicals are stored separated outside the cell and fed in when electricity is required, e.g. hydrogen-oxygen fuel.

  • Hydrogen and oxygen gases are fed into two separate platinum-containing electrodes. These electrodes are usually made by coating a porous ceramic material with a thin layer of platinum, (cheaper than using a solid platinum rods, provides a larger surface area).
  • The electrodes are separated by an ion-exchange membrane that allows protons (H+) to pass through it, but stops electrons going through it.
30
Q

What happens in a fuel cell?

A
  • At the left hand (-ve) electrode, hydrogen enters and gives up electrons to form H+ ions; H2(g) –> 2H+(aq) + 2e-.
  • The electrons flow through an electrical circuit where they can power a motor. The H+ ions flow through the electrolyte to a second electrode. Here they react with oxygen and electrons to form water - the only emission from the unit: 4H+(g) + O2(g) –> 2H2O(l).
  • Overall: 2H2(g) + O2(g) –> 2H2O(l) (emf = 1.2V)
31
Q

What are the advantages of fuel cells?

A

+They don’t need recharging, as long as hydrogen and oxygen are supplied, the cell will continue to produce electricity.
+The only product is water, so no toxic chemicals need to be disposed of and no CO2 emissions from the cell itself.

32
Q

What are the disadvantages of fuel cells?

A
  • Energy is needed to produce a supply of hydrogen and oxygen. This can be produced by the electrolysis of water, reusing the waste product, but this requires energy (usually from burning fossil fuels), so the whole process isn’t carbon neutral.
  • Hydrogen is highly flammable so it needs to be handled carefully when stored and transported.
33
Q

Are fuel cells carbon neutral?

A

Fuel cells can’t be considered carbon neutral because hydrogen is produced from crude oil, a non-renewable resource. Alternatively it could be made by electrolysis of water, but the electricity for this is often generated from crude oil, emitting CO2.
-Hydrogen-powered vehicles are require an infrastructure of hydrogen filling stations to be built.

Chemistry: CHEM 5 (6 decks)

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