Answers Chemistry, 12e (Brown, LeMay, Bursten, Murphy) chapter 5 & 6
Oxidation and Reduction Gain of Electrons is Reduction …web.uvic.ca/~asirk/Electro_2.pdf ·...
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Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species to another. Substance which receives electrons is the oxidising agent or oxidant and is itself reduced. Substance which gives up its electrons is the reducing agent or reductant and is itself oxidised. Oxidation and Reduction GER LEO Gain of Electrons is Reductio Loss of Electrons is Oxidation
Transcript of Oxidation and Reduction Gain of Electrons is Reduction …web.uvic.ca/~asirk/Electro_2.pdf ·...
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Electrochemical Reactions
In electrochemical reactions, electrons are transferred from one species to another.
Substance which receives electrons is the oxidising agent or oxidant and is itself reduced.
Substance which gives up its electrons is the reducing agent or reductant and is itself oxidised.
Oxidation and Reduction
GER
LEO
Gain of Electrons is Reduction
Loss of Electrons is Oxidation
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Balancing Redox Equations (Chem 120)
Reduction Potentials
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Reduction Potentials
Voltaic Cells In spontaneous oxidation-reduction (redox)
reactions, electrons are transferred and energy is released.
Figure 16.2
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Voltaic Cells
We can use that energy to do work if we make the electrons flow through an external device.
We call such a setup a voltaic cell. Figure 16.3
Voltaic Cells A typical cell looks
like this.
The oxidation occurs at the anode.
The reduction occurs at the cathode.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Voltaic Cells
Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop.
Figure 16.4
Voltaic Cells
Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced: Cations move toward
the cathode.
Anions move toward the anode.
Figure 16.4
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Voltaic Cells
In the cell, electrons leave the anode and flow through the wire to the cathode.
As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.
Voltaic Cells
As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode.
The electrons are taken by the cation, and the neutral metal is deposited on the cathode.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Electromotive Force (emf)
Water only spontaneously flows one way in a waterfall.
Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy.
Figure 16.6
Electromotive Force (emf)
The potential difference between the anode and cathode in a cell is called the electromotive force (emf). It is measured in volts (V):
It is also called the cell potential, and is designated Ecell.
1 V = 1 JC
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Standard Cell Potentials
The cell potential at standard conditions can be found through this equation:
Because cell potential is based on the potential energy per unit of charge, it is an intensive property.
Ecell = Ered (cathode) − Ered (anode)
Standard Hydrogen Electrode Their values are referenced to a standard
hydrogen electrode (SHE).
By definition, the reduction potential for hydrogen is 0 V.
2 H+ (aq, 1M) + 2 e− H2 (g, 1 atm)
Figure 16.7
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Standard Reduction (Half-Cell) Potentials
Reduction potentials for many electrodes have been measured and tabulated.
Table 16.1
Cell Potentials
For the oxidation in this cell:
For the reduction:Ered = +0.34 V
Ered = −0.76 V
Figure 16.8
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Cell Potentials
= +0.34 V − (−0.76 V)= +1.10 V
Ecell = Ered (cathode) − Ered (anode)
Strengths of Oxidisingand Reducing Agents
The strongest oxidisers have the most positive reduction potentials.
The strongest reducers have the most negative reduction potentials.
Figure 16.11
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Strengths of Oxidisingand Reducing Agents
The greater the difference between the two, the greater the voltage of the cell.
Figure 16.9
Cell emf and G G for a redox reaction can be found by using
the equation:G = −nFE
A positive value of E and a negative value of Gboth indicate that a reaction is spontaneous.
Consequently, under standard conditions:G = −nFE
n is the number of moles of electrons transferred.F is called Faraday’s constant: 1 F = 96,485 C/mol = 96,485 J/V mol
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Cell emf Under Non-standard Conditions - The Nernst Equation
Remember that:G = G + RT ln Q
This means:−nFE = −nFE + RT ln Q
Cell emf Under Non-standard Conditions - The Nernst Equation
Dividing both sides by −nF, we get the Nernst equation:
E = E − RTnF ln Q
or, using base-10 logarithms,
E = E − 2.303 RTnF log Q
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Cell emf Under Non-standard Conditions - The Nernst Equation
At room temperature (298 K):
Thus the equation simplifies to:
E = E − 0.0592n log Q
2.303 RTF = 0.0592 V
Concentration Cells Notice that the Nernst equation implies that a cell
could be created that has the same substance at both electrodes.
For such a cell, would be 0, but Q would not. Therefore, as long as the concentrations are
different, E will not be 0.
Ecell
Figure 16.12
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Batteries and Fuel Cells -The Hydrogen Fuel Cell
Figure 16.16
Corrosion of Iron
Figure 16.17
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Preventing the Corrosion of Iron
Figure 16.18
Electrolysis
Voltaic cells are based on a spontaneous redox reaction, but it is possible to use electricity to cause non-spontaneous redox reactions to occur.
For example, the electrolysis of molten sodium chloride to produce sodium metal and chlorine gas.Figure 16.20
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Education Australia
Electrolysis Electroplating is another instance where
elecrolysis is used to deposit a thin layer of one metal on another to improve beauty or resistance to corrosion, for example the nickel-plating of steel.
One point worth noting is that in both voltaic and electrolytic cells, the anode is where oxidation occurs (and the cathode is where reduction occurs).
Figure 16.21
Quantitative Aspects of Electrolysis Half-reactions show how many electrons are
needed to achieve an electrolytic process and so it follows that the amount of a substance oxidised (or reduced) is directly proportional to the number of electrons passed into the cell.
Figure 16.22
