STANDARD ELECTRODE POTENTIALS. THE STANDARD HYDROGEN ELECTRODE In order to measure the potential of...
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Transcript of STANDARD ELECTRODE POTENTIALS. THE STANDARD HYDROGEN ELECTRODE In order to measure the potential of...
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STANDARD ELECTRODE POTENTIALS
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THE STANDARD HYDROGEN ELECTRODE
In order to measure the potential of an electrode, it is compared to a reference electrode – the standard hydrogen electrode.
The standard half reaction is the reduction of hydrogen:
2H+ + 2e- → H2
The electrode potential of this half cell = 0,00V
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Operating conditions:
STP
Temperature = 298K
Pressure = 1 atm
or 101,3 2kPa
[H+] = 1 mol.dm-3
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Measuring the standard electrode potential of zinc by connecting a zinc electrode to the hydrogen electrode.
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The hydrogen electrode is the cathode.
2H+(aq) + 2e- → H2(g)
The zinc electrode is the anode.
Zn(s) → Zn2+(aq) + 2e-
The reading on the voltmeter is 0,76V and because electrons flow from the zinc to the hydrogen electrode
Eo for the zinc half cell = - 0,76V
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Abbreviated redox table
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Full Redox potential table
Any substance on the right will spontaneously react with something above it on the right of the table – and vice versa.
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When using the table:
• strong reducing agents are at the top of the table.
• when combining half reactions, the one higher up the table will be the reducing agent.
• when combining half reactions, the half reaction located higher up the table is written in reverse (from left to right) and the one lower down from right to left.
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Using the table of standard electrode potentials:
• determine the emf of a cell.
• predicting whether a redox reaction will occur spontaneously;
• balancing redox reactions.
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Determining the emf of a cell:
Eo cell = Eo reduction – Eo
oxidation
= Eo oxidising agent – Eo
reducing agent
= Eo cathode – Eo
anode
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For the cell:
Zn / Zn2+ // Cu2+ / Cu
The zinc electrode is the reducing agent and the anode as oxidation occurs at that electrode.
Eo cell = Eo Cu – Eo
Zn
= 0,34 – (-0,76)
= +1,10V
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Predicting whether a redox reaction will occur spontaneously.
• Write down the equation as you expect it to occur.
• from the equation decide which half cell is the reducing/oxidising agent.
• Determine Eocell based on this information.
• If Eocell is a positive number, the reaction, as
written, is non-spontaneous.
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Will the following reaction occur spontaneousky:
Pb2+(aq) + 2Br-
(aq) → Br2(l) + Pb(s)
Pb2+ + 2e- → Pb Eo = - 0.13V reduction
2Br- → Br2 + 2e- Eo = + 1,06V oxidation
Eocell = Eo
Pb – EoBr = (- 0,13) – (+ 1,06)
= - 1,19V
Eo is negative and so the reaction is non-spontaneous
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Balancing redox equations.
Hydrogen sulphide reduces potassium dichromate. Write a balanced equation for this reaction.
H2S → 2H+ + S + 2e-
Cr2O72- + 14H+ + 6e- → Cr3+ + 7H2O
Multiply the first equation by 3 to balance the electrons, cancel out any ions/molecules that occur on opposite sides of the equation and add up the remaining ions/molecules to give the balanced ionic equation.
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3H2S → 6H+ + 3S + 6e-
Cr2O72- + 14H+ + 6e- → Cr3+ + 7H2O
3H2S + Cr2O72- + 8H+ → Cr3+ + 7H2O + 3S