Electrochemistry & Fuel Cells Chemistry A OCR Unit 5

HideShow resource information
Preview of Electrochemistry & Fuel Cells Chemistry A OCR Unit 5

First 128 words of the document:

Electrochemistry & Fuel Cells
Electrochemical Cells
Shorthand Notation
The Standard Hydrogen Electrode
Standard Electrode Potential (E)
Electromotive Force (emf)
Measuring Standard Electrode Potentials
E for Metal/Metal Ion Half-Cell
E for Metal Ion/Metal Ion Half-Cell
E for Non-Metal/Non-Metal Ion Half-Cell
The Electrochemical Series
Oxidising and Reducing Agents
Calculating Cell Potentials
The Reaction in the Electrochemical Cell
The Feasibility of Redox Reactions
Limitations to Predictions
The Effect of Concentration on Electrode Potentials
Storage Cells & Fuel Cells
Storage Cells
Fuel Cells
Hydrogen-Oxygen Fuel Cell
Methanol Fuel Cells
Advantages & Disadvantages of Hydrogen Fuel Cells
The Hydrogen Economy

Other pages in this set

Page 2

Preview of page 2

Here's a taster:

An electrochemical cell uses the electron transfer from a redox reaction to produce a voltage
(electrical energy).
The cell consists of two halfcells that correspond to the two halfequations of the redox
reaction, and each halfcell consists of two species of the same element in different
oxidation states.
The copper halfcell consists of copper
with an oxidation state of zero (Cu) and
copper with an oxidation state of +2
(Cu2+).…read more

Page 3

Preview of page 3

Here's a taster:

The Copper/Zinc Cell
By convention, the oxidised halfcell is
drawn on the left and the reduced
halfcell on the right.…read more

Page 4

Preview of page 4

Here's a taster:

The Standard Hydrogen Electrode
When measuring the electrode potential of a single halfcell, the standard hydrogen electrode
is used as a reference. It is given an electrode potential of 0.00V.
H+(aq) from an acid solution (e.g. HCl) 1 moldm=3
H2 (g) at 100kPa and 298K
Platinum as an electrode as there is no solid present.
Half Equation: 2H+(aq) + 2e H2 (g)
Emf: 0.…read more

Page 5

Preview of page 5

Here's a taster:

Measuring Standard Electrode Potentials
The standard electrode potential is found by measuring the voltage of a cell with the
standard hydrogen electrode. The polarity relative to the hydrogen halfcell tells us the sign
of the electrode potential.
E for Metal/Metal Ion HalfCell
E.g. Cu2+/Cu
Measured Voltage: 0.34 V
Polarity of Copper
Electrode: Positive
Hence E = +0.34V
E for Metal Ion/Metal Ion HalfCell
These are halfcells where both species are aqueous ions.…read more

Page 6

Preview of page 6

Here's a taster:

E.g. Fe3+/Fe2+
Measured voltage: +0.77 V
Polarity of Iron III/Iron II Cell: Positive
Hence E = +0.77 V
E for NonMetal/NonMetal Ion HalfCell
Platinum electrodes are used for the oxidised and reduced species.…read more

Page 7

Preview of page 7

Here's a taster:

Measured Voltage: +1.36 V
Polarity of Chloride/chloride HalfCell: Positive
Hence E = +1.36 V
The Electrochemical Series
This is a list of standard electrode potentials in numerical order & all halfequations are
written as reductions.…read more

Page 8

Preview of page 8

Here's a taster:

The more positive the E the greater the tendency that the species will be reduced.
The more negative the E the easier it is to oxidise the species.
Oxidising and Reducing Agents
An oxidising agent is reduced itself as it accepts electrons, so strong oxidising agents
have highly positive potentials.
From the list F2 is a good oxidising agent as is Au+.
A reducing agent is oxidised itself as it donates electrons, so strong reducing agents
have highly negative potentials.…read more

Page 9

Preview of page 9

Here's a taster:

If done correctly, the Ecell value will always be positive otherwise the reaction will not
The Reaction in the Electrochemical Cell
In cells, chemical energy is converted into electrical energy, and once the chemicals are used
up, the reaction stops. You can view the E values to predict what reaction occurs.
For example:
Zn2+(aq) + 2e Zn(s) E = 0.76V
Ag+(aq) + e Ag(s) E = +0.80V
The equation with the most positive E will be reduced and hence gains electrons.…read more

Page 10

Preview of page 10

Here's a taster:

Fe2+ Fe3+ + e-
Then check the electrons balance, which they do and combine the equations:
½Cl2 + Fe2+ Fe3+ Cl-
Now compare the result with the original question; "Can chlorine oxidise Iron (II)
to Iron (III)?" The equation corresponds as chlorine is oxidising iron (II) to iron (III). So the
reaction is feasible.
To double check, the cell potential can be calculated and it should come out as a positive
result if the reaction is feasible:
1.36 ­ 0.77 = + 0.…read more



thanks for sharing :)

Similar Chemistry resources:

See all Chemistry resources »See all resources »