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Chemistry A-Level > Electrode Potentials (Chapter 23.4, 23.5, 23.6) > Flashcards

Electrode Potentials (Chapter 23.4, 23.5, 23.6) Flashcards

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1
Q

What is a voltaic cell?

A

A type of electrochemical cell that converts chemical energy into electrical energy

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2
Q

What is required for a voltaic cell?

A

Chemical reactions that transfer electrons from one species to another (redox reactions) - bc electrical energy results from the movement of electrons

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3
Q

What does a half-cell contain?

A

The chemical species present in a redox half-equation

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4
Q

How is a voltaic cell made?

A

By connecting together two different half-cells, allowing electrons to flow

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5
Q

What is necessary in a voltaic cell and why?

A

The chemicals in the two half-cells must be kept apart - if they were allowed to mix, electrons would flow in a uncontrolled way and heat energy, instead of electrical, would be released

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6
Q

What do metal/metal ion half-cells consist of?

A

A metal rod dipped into a solution of its aqueous metal ion

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7
Q

What is the place where the metal is in contact with its ions called?

A

The phase boundary

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8
Q

What happens at the phase boundary in a metal/metal ion half-cell?

A

An equilibrium will be set up in which the forwards reaction show reduction and the reverse reaction shows oxidation e.g. Zn2+ + 2e- <=> Zn

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9
Q

What happens in an isolated half-cell?

A

There is no net transfer of electrons either into or out of the metal

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10
Q

What does the direction of electron flow depend on when two half-cells are connected?

A

The relative tendency of each electrode to release electrons

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11
Q

Describe ion/ion half-cells

A
  • They contain ions of the same element in different oxidation states e.g. a mixture of Fe2+ and Fe3+
  • e.g. redox equilibrium: Fe3+ + e- <=> Fe2+
  • There is no metal to transport electrons ∴ an inert metal electrode made out of platinum (Pt) is used
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12
Q

What happens in a cell with 2 metal/metal ion half-cells connected?

A

The more reactive metal releases electrons more readily and is oxidised

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13
Q

Describe the electrodes in an operatic voltaic cell

A
  • The electrode with the more reaction metal loses electrons and is ∴ oxidised - this is the negative electrode
  • The electrode with the less reactive metal gains electrons and is ∴ reduced - this is the positive electrode
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14
Q

Define the standard electrode potential

A

The e.m.f (electron motive force) of a half-cell connected to a standard hydrogen half-cell under standard conditions of 298K , solution concentrations of 1mol/dm3 and pressure of 100 KPa

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15
Q

What does the standard electrode potential measure?

A

The tendency of a species to be reduced and gain electrons

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16
Q

What is the standard half-cell?

A

A half-cell containing hydrogen gas and a solution of H+ ions with a platinum electrode to allow electrons into and out of the half-cell

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17
Q

What are the standard conditions for a half-cell?

A
  • Concentration of solutions = 1M
  • Temperature = 298K (25 degrees)
  • Pressure = 100KPa (1 atm)
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18
Q

What is the electrode potential of a standard hydrogen half-cell?

A

0V

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19
Q

What does the sign of an electrode potential show?

A

The sign of the half-cell connected to the standard hydrogen half-cell and the relative tendency to gain electrons compared with the hydrogen half-cell

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20
Q

Describe how to measure a standard electrode potential?

A

1) the half-cell electrode is connected to a standard hydrogen electrode by a wire, allowing a controlled flow of electrons
2) the two solutions are connected with a salt bridge which allows ions to flow and typically contains a concentrated solution of an electrolyte that does not react with either solution e.g. a strip of filter paper soaked in aqueous potassium nitrate (KNO3)

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21
Q

What do you need to label on a voltaic cell diagram?

A

1) solutions and concentrations
2) electrodes (the metals and charges)
3) voltmeter
4) salt bridge

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22
Q

The more negative the electrode potential value…

A

1) the greater the tendency to lose electrons and undergo oxidation
2) the less the tendency to gain electrons and undergo reduction

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23
Q

The more positive the electrode potential value…

A

1) the greater the tendency to gain electrons and undergo reduction
2) the less the tendency to lose electrons and undergo oxidation

24
Q

What elements tend to have negative electrode potential values?

A

Metals

25
Q

What elements tend to have positive electrode potential values?

A

Non-metals

26
Q

What can you conclude in general about electrode potentials?

A

1) the more negative the electrode potential value the greater the reactivity of the metal in losing electrons
2) the more positive the electrode potential value, the greater the reactivity of a non-metal in gaining electrons

27
Q

What is the e.m.f. of a cell called?

A

The cell potential, Ecell

28
Q

What can the standard Ecell be recorded from the voltmeter?

A

Using two standard half cells, connected by a salt bridge with the electrodes connected to a voltmeter and wire

29
Q

How do you work out the standard Ecell?

A

electrode potential of positive electrode - electrode potential of negative electrode

30
Q

How can you use electrode potentials to predict redox equations?

A
  • An oxidising agent takes electrons way from the species being oxidised ∴ oxidising agents are reduced and are on the left of the equation
  • A reducing agent adds electrons to the species being reduced ∴ reducing agents are oxidised and on the right the equation
31
Q

Where is the strongest oxidising agent in a table of equations?

A

At the bottom on the left

32
Q

Where is the strongest reducing agent in a table of equations?

A

At the top on the right

33
Q

How can you predict the feasibility of redox reactions from electrode potential values?

A
  • A reaction should take place between an oxidising agent and a reducing agent, provided that the redox system of the oxidising agent has a more positive electrode potential than the redox system of the reducing agent
  • e.g. an oxidising agent will not react if there are no redox systems with a less positive electrode value
34
Q

Describe the direction of reaction of redox systems

A
  • The redox systems with the more positive electrode potential value will react from left to right
  • The redox system with the more negative electrode potential value will react from right to left
35
Q

What are the two limitations of predictions using electrode potential values?

A

1) reaction rate
2) standard conditions
3) concentration

36
Q

Why is reaction rate a limitation when making predictions using electrode potential values?

A
  • Even though electrode potentials may indicate thermodynamic feasibility of a reaction, they give no indication of the rate of reaction
  • ∴ a reaction deemed feasible may not actually be so, if the reaction has a very large activation energy and ∴ a very slow reaction rate (like with deltaG)
37
Q

Why are standard conditions a limitation when making predictions using electrode potential values?

A
  • Actual conditions may be different from standard conditions used to record standard electrode potential values - this will affect the value of the electrode potential
  • Standard electrode potentials apply to aqueous equilibria but many reactions that take place are not aqueous
38
Q

Why is concentration a limitation when making predictions using electrode potential values?

A
  • Electrode potentials are measured using concentrations of 1M but many reactions take place using concentration or dilute solutions - if the concentration of the solution is not 1M, then the value of the electrode potential will be different from the standard electrode potential
  • e.g. in a Zn half cell, if [Zn2+] is greater than 1M, the equilibrium will shift to the right, removing electrons from the system and making the electrode potential less negative
  • if [Zn2+] is less than 1M, the equilibrium will shift tot eh left, increasing the number of electrons in the system and making the electrode potential more negative
  • Any change to the electrode potential will affect the value of the overall cell potential
39
Q

What is the key-requirement for modern day cells and batteries?

A

For the two electrodes to have different electrode potentials

40
Q

What are the 3 types of cells?

A

1) primary
2) secondary
3) fuel

41
Q

Describe primary cells

A
  • Non-rechargeable and designed to be used once only
  • Electrical energy is produced by oxidation and reduction at the electrodes, but the reactions cannot be reversed
  • Eventually the chemicals will be used up, voltage will gall, the battery will go flat and the cell is discarded or recycled
  • Typically alkaline, based on Zn/MnO2 and a KOH alkaline electrolyte
  • Uses: low-current, long-storage devices e.g. wall clocks and smoke detectors
42
Q

Why are secondary cells rechargeable?

A

Bc the cell reaction producing the electrical energy can be revered during recharging - the chemicals in the cell are ∴ regenerated and the cell can be used again

43
Q

What are three examples of secondary cells?

A

1) lead-acid batteries used in car batteries
2) nickel-cadmium cells and nickel-metal hydride cells - cylindrical batteries used in radios and torches
3) lithium-ion and lithium-ion polymer cells used in modern appliances e.g. laptops and cameras

44
Q

Describe lithium ion and lithium-ion polymer cells

A
  • Negative electrode: Li => Li + e-
  • Positive electrode: Li+ + CoO2 + e- => LiCoO2
  • Limitations of lithium ion cells: can become unstable at high temperatures and have ignited laptops and phones ∴ care must be taken in their recycling as lithium is a very reaction metal
  • Risks of lithium based cells: toxicity and fire
  • Benefits of lithium based cells: flexible so can come in various shapes and sizes
45
Q

Describe fuel cells

A
  • They use the energy from the reaction of a fuel and oxygen to create a voltage
  • The fuel and oxygen flow into the fuel cell, the products flow out and the electrolyte remains in the cell
  • They can operate continuously, provided that the fuel and oxygen are supplied into the cell
  • They do not have to be recharged
46
Q

Why are hydrogen/hydrogen-rich (e.g. methanol) fuel cells very common?

A

They produce no CO2 during combustion, only water (using an acid or alkali electrolyte)

47
Q

What are the equations for an alkali hydrogen fuel cell?

A
  • Oxidation: H2 + 2OH- => 2H2O + 2e-
  • Reduction: 1/2O2 + H2O + 2e- => 2OH-
  • Overall: H2 + 1/2O2 => H2O
48
Q

What are the equations for an acid hydrogen fuel cell?

A
  • Oxidation: H2 => 2H+ + 2e-
  • Reduction: 1/2O2 + 2H+ + 2e- => H2O
  • Overall: H2 + 1/2O2 => H2O
49
Q

What is the overall equation and Ecell for an acid and alkali hydrogen fuel cell?

A

Overall equation: H2 + 1/2O2 => H2O

Ecell = 1.23V

50
Q

What direction do electrons flow in?

A

Negative electrode to positive electrode

51
Q

Describe gas/ion half-cells

A
  • Electrode: Pt
  • Gas at 1atm
  • e.g. Cl2 + 2e- => 2Cl-
52
Q

What does the salt bridge do?

A

It allows ions to flow

53
Q

Why is a platinum electrode used?

A

It is unreactive but conducts electricity so that electrons can be transferred

54
Q

When is a reaction feasible?

A

When the Ecell is positive

55
Q

Why are weak acids not used as a source of H+ ions for a hydrogen half-cell?

A

Because the [H+] is unpredictable and keeps changing

56
Q

What are 2 advantages and two disadvantages of using hydrogen fuel cells?

A
  • Advantages: 1) they operate continuously
    2) they release no greenhouse or toxic gases e.g. CO2
  • Disadvantages: 1) hydrogen gas is very flammable and ∴ dangerous to store
    2) the fuel cell reaction has a slow reaction rate

Chemistry A-Level (27 decks)

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