What are electrochemical cells
•A cell has two half-cells.
•The two half cells have to be connected with a salt bridge.
•Simple half cells will consist of a metal (acts an electrode)
• and a solution of a compound containing that metal (eg Cu and CuSO4).
•These two half cells will produce a small voltage if connected into a circuit. (i.e. become a Battery or cell).
diagram 1
Why does a voltage form
In diagram1
- When connected together zinc half-cell has more of a tendency to oxidise to the Zn2+ ion and release electrons
- more a tendency than the copper half-cell.
(Zn→ Zn2+ + 2e-)
..
- More electrons will so build up on zinc electrode than copper electrode.
- A potential difference is created between two electrodes.
- zinc strip is negative terminal and copper strip is the positive terminal.
..
- This potential difference is measured with a high resistance voltmeter
- and is given the symbol E, which is always positive
Why use a high resistance voltmeter?
- The voltmeter needs to be of very high resistance to stop current from flowing in the circuit.
- In this state it is possible to measure the maximum possible potential difference (E).
- The reactions will not be occurring bc the very high resistance voltmeter stops current from flowing.
Why do we have salt bridges
- The salt bridge is used to connect up the circuit. The free moving ions conduct the charge.
- A salt bridge is usually made from a piece of filter paper (or material)
- soaked in a salt solution, usually potassium nitrate.
- salt should be unreactive with electrodes/electrode solutions
> eg KCl isnt suitable for copper systems as chlordie ions form complexes with copper ions
..
- A wire is not used because the metal wire would
- set up its own electrode system with the solutions.
What happens if current is allowed to flow?
- If the voltmeter is removed and replaced with a bulb or if the circuit is short circuited, a current flows.
- reactions will then occur separately at each electrode. voltage will fall to zero as reactants are used
..
- The most positive electrode will always undergo reduction.
- Cu2+ (aq) + 2e → Cu(s) (positive as electrons are used up)
- The most negative electrode will always undergo oxidation.
- Zn(s) → Zn2+(aq) + 2e (negative as electrons are given off)
What are cell diagrams
diagram 2
• The solid vertical line represents the boundary between phases e.g. solid (electrode) and solution (electrolyte)
•The double line represents salt bridge between two half cells
•the voltage produced is indicated
the more positive half cell is
• written on the right if possible (but this is not essential)
If systems dont include metals
- If a system doesnt include a metal that can act as an electrode, then a platinum electrode
- ## must be used and included in cell diagram. It provides conducting surface for electron transfer.
- A platinum electrode is used because it is unreactive and can conduct electricity.
..
e.g. for Fe2+ (aq) → Fe3+ (aq) + e-
- there is no solid conducting surface, a Pt electrode must be used.
- The cell diagram is drawn as:
|| Fe3+ (aq), Fe2+ (aq) |Pt
-
- still with more oxidised form near double line
- a comma separates oxidised from reduced species
..
- If the system contains several species e.g. MnO4 + 8H+ + 5e → Mn2+ +4H2O
- then in cell diagram the balancing numbers H+ ions and H2O can be left out
If a half equation for a cell diagram has several physical states (and doesnt include metals)
- If half equation has several physical states then solid vertical line
- should be used between each diff state boundary
4e + 2H20 (l) +02 (g)→ 4OH (aq)
→||O2 | H20, OH- | Pt
Cl2 (g) + 2e → 2Cl-(aq)
||Cl (g)| Cl-(aq) |Pt
As phase line also separates oxidised and reduced terms, a comma is not necessary here
Measuring the electrode potential of a cell
- its not possible to measure absolute potential of a half electrode on its own.
- its only possible to measure potential difference between two electrodes
..
• To measure it, it has to be connected to another half-cell of known potential,
• and potential difference between the two half-cells measured.
•
• by convention we can assign a relative potential to each electrode by linking it to reference electrode
• (hydrogen electrode), which is given a potential of zero Vol
Whats standard hydrogen electrode
diagram3
- potential of all electrodes are measured by comparing their
- potential to standard hydrogen electrode (SHE)
> assigned potential of 0v
Hydrogen electrode equilibrium is
H2(g) <>2H+ + 2e-
For cell diagram: Pt| H2(g) | H+(aq)
Components of standard hydrogen electrode
> to make electrode a standard reference electrode some conditions apply:
1. Hydrogen gas at pressure of 100kPa
2. Solution containing hydrogen ion at 1.0 mol dm-3 (solution usually 1 mol dm-3 HCI)
3. Temperature at 298K
4. Platinum electrode
..
- Because equilibrium does not include a conducting metal surface, a platinum wire is used
- which is coated in finely divided platinum.
- (The platinum black is used bc it is porous; absorbs the hydrogen gas.)
..
- Standard conditions are needed because position of redox equilibrium will change w conditions.
- For example, in the equilibrium:
M^n+(aq) + n e- <> M(s)
-
- An increase in concentration of M^n+ would move equilibrium to right, so making potential more positive.
Secondary standards
Secondary standards
-The SHE is difficult to use, so often a diff standard is used; easier to use.
- These other standards are themselves calibrated against the SHE
-
- This is known as using a secondary standard - i.e. a standard electrode
- wch has been calibrated against primary standard.
- The common ones are: silver / silver chloride (calomel electrode)
E = +0.22 V
E =+0.27 V
Standard electrode potentials
Understand diagram4
- When an electrode system is connected to the hydrogen electrode system,
- ## and standard conditions apply , potential difference measured is called the standard electrode potential
- The standard conditions are
•all ion solutions at 1 mol dm-3
•temperature 298 K
•gases at 100 kPa pressure
•no current flowing
..
- Standard electrode potentials are found in data books; quoted as:
Li (aq) | Li (s)
E= -3.03V
more oxidised form on left
may also be quoted as half equations
Li (aq) + e- > Li (s)
E= -3.03V
but again the more oxidised form is on the left
Calculating emf of a cell 🟡
Mg(s) | Mg2+(aq) | | Cu2+(aq) | Cu (s) E= +1. 1V
> In order to calculate Ecell, we must use ‘standard electrode potentials for half cells
🟡Ecell= Erhs - Elhs
For the cell diagram above
Ecell = 0.34 - -2.37
=+2.71V
(Youll be given the values of standard electrode potential from each side)
Using Electrode Potentials (first example)
The easiest way to use electrode potentials is as follows:
- most useful application of electrode potentials is to show direction
- of spontaneous change for redox reactions
..
- any two half equations
- The more negative half cell will always oxidise (go backwards)
Mg2+ (aq) + 2e → Mg(s) E= -2.37V
Cu2+ (aq) + 2e → Cu (s) E = +0.34V
-
- The more positive half cell will always reduce (go forwards)
The reaction would be
Mg + Cu2+ → Cu + Mg 2+
- To work out Ecell that corresponds to this spontaneous change use
🟡 Ecell = Ered - Eox
- A spontaneous change will always have a positive Ecell.
Using electrode potentials (second example)
- The most negative electrode will oxidise and go from right to left
- The half equation is therefore
Zn(s) → Zn2+ (aq) +2e-
- - Electrons are given off (lost) and travel to positive electrode.
Zn2+(aq) + 2e → Zn(s) E= - 0.76V
Fe2+(aq) + 2e → Fe(s) E= -0.44V
- - The more positive electrode will reduce and go from left to right
- Fe2+ (aq) +2e-→ Fe(s)
- Electrons arrive at this electrode and are absorbed (gained)
..
To get the full equation of the reaction add the two half reactions together,
> cancelling out the electrons.
Zn + Fe2+ → Fe + Zn2+
Using series of standard electrode potentials
diagram 5
diagram 6 examples
- Most strong reducing agents found here as more +ve
- increasing tendency for species on left to reduce, and act as oxidising agents
- As more -ve increasing tendency for species on right to oxidise, and act as reducing agents
..
- To work out the Ecell from two standard electrode potentials then use
Ecell = Ered - Eox
- most powerful reducing agents will be found at most negative end of series on right
- (i.e. the one with the lower oxidation number)
-
- The most powerful oxidising agents will be found at the most positive end
- of the series on the left (i.e. the one with the higher oxidation number).
Effect of conditions on Cell voltage Ecell
- ## The effects of changing conditions on E cell can be made by applying le Chatelier’s principle.> If current is allowed to flow, cell reaction will occur and Ecell will fall to zero
as reaction proceeds and reactant concs drop.
..
- Ecell is a measure of how far from equilibrium the cell reaction lies.
- The more positive the Ecell the more likely the reaction is to occur.
Effect of concentration on Ecell
- Looking at cell reactions is a straight forward application of le Chatelier.
- So increasing conc of ‘reactants’ would increase Ecell
- and decreasing them would cause Ecell to decrease.
..
Zn2+ (aq) + 2e- > Zn(s) E= -0.76
Fe2+ (aq) + 2e- > Fe(s) E= -0.44
> Zn + Fe2+ → Fe + Zn2+ E= +0.32
- increasing the conc of Fe2+, increasing conc of Zn2+ wd cause Ecell to increase.
Effect of temperature on Ecell
- Most cells are exothermic in the spontaneous direction so applying Le Chatelier
- to a temp rise to these would result in a decrease in Ecell
- because the equilibrium reactions would shift backwards.
.. - if ecells positive it indicates reaction might occur ; but theres still a possibility it may not
- or will occur so slow it effectively wont happen- if has high activation energy reaction wont occur
If the reaction has a high activation energy the reaction will not occur.
Electrochemical cells
- Electrochemical cells can be used as a commercial source of electrical energy
- Cells can be non-rechargeable (irreversible), rechargeable and fuel cells.
> Should be able to work out Ecell for given half reactions.
..
- You should be able to convert between standard electrode potential half cells,
- full cell reactions and
- cell diagrams
- and be able to calculate potentials from given data.
LEARN THESE HARD EXAMPLES
diagram 7 >primary
diagram8> secondary
Fuel cells
- A fuel cell uses energy from reaction of a fuel with oxygen to create a voltage.
- You should learn details of lithium cell and hydrogen fuel cell in alkaline conditions
.. - Hydrogen fuel cell (potassium hydroxide electrolyte)
- Alkaline Conditions
Standard electrode potentials
2e + 2H20 → H2 + 20H- E=-0.83V
4e + 2H20 + 02 → 40H- E=+0.4V - Equations at electrodes
Negative: H2 + 2OH → 2e + 2H20
Positive: 4e + 2H20 + 02 → 40H-
Overall reaction 2H2 + 02 → 2H20 E=1.23V
..
- Fuel cells will maintain a constant voltage over time bc they are continuously fed with fresh O2 and H2 - so maintaining constant conc of reactants. This differs from ordinary cells
- where voltage drops over time as the reactant concs drop.
diagram9
Using standard conditions (hydrogen fuel cell-KOH electrolyte)
- Using standard conditions: The rate is too slow to produce an appreciable
current - Higher temps are so used to increase rate but reaction is exothermic
- so by applying le Chatelie would mean E cell falls
- A higher pressure can help counterac this.
Fuel cells pros and cons
• Advantages of fuel cells over conventional petrol or diesel-powered vehicles
— less pollution and less CO2. (Pure hydrogen emits only water whilst
- hydrogen-rich fuels produce only small amounts of air pollutants and CO2).
— greater efficiency
..
• Limitations of hydrogen fuel cells
— expensive
— storing and transporting hydrogen, in terms of safety, feasibility of a
- pressurised liquid and a limited life cycle of a solid ‘adsorber’ or ‘absorber’
— limited lifetime (requiring regular replacement and disposal) and high production costs
— use of toxic chemicals in their production
..
- Hydrogen is readily available by the electrolysis of water, but this is expensive.
- To be a green fuel the electricity needed would need to be produced from renewable resources
Hydrogen can be stored in fuel cells
(i) as a liquid under pressure,
(ii) adsorbed on the surface of a solid material,
(iii) absorbed within a solid material;
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1.1 Atomic Structure31
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1.3 Bonding27
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equations!6
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1.2 Amount Of Substance41
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3.1 Intro To Organic Chemistry13
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3.2 Alkanes22
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3.3 Haloalkanes17
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Oxidation, Reduction And Redox Equations14
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Periodicity, Group 2 And 738
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Alkenes22
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Energetics19
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Alcohols19
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Kinetics9
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Organic Analysis12
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Chemical Equilibria and Le Chatelier's principle29
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Kp5
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Period 3 Oxides (A Level)21
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Organic Naming And Isomerism (A Level)15
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Rate Equations (Kinetics A Level)19
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Carbonyls: Aldehydes And Ketones // Carboxylic Acids And Derivatives (A Level)33
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Thermodynamics31
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Physical Exam Questions56
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Electrode Potentials26
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A2 - Physical148
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