Transcript Electrochemistry Redox reactions
Electrochemistry
Redox reactions
Ox
1 Re
d
2 Re
d
1
Ox
2
oxidant reduc
tan
t examples
:
M a
ne
M
(
a
n
)
M a
M
(
a
n
)
ne
The half-reactions for Sn 2+ (
aq
) + 2Fe 3+ (
aq
) Sn 4+ (
aq
) + 2Fe 2+ (
aq
) Oxidation: Sn Reduction: 2Fe 2+ 3+ ( (
aq aq
) are Sn 4+ (
aq
) +2e ) + 2e 2Fe 2+ (
aq
) In oxidation: electrons are products.
In reduction: electrons are reactants.
Voltaic or Galvanic Cells
• The energy released in a spontaneous redox reaction can be used to perform electrical work.
• Voltaic or galvanic cells are devices in which electron transfer occurs via an external circuit.
• Voltaic cells use spontaneous reactions.
• If a strip of Zn is placed in a solution of CuSO 4 , Cu is deposited on the Zn and the Zn dissolves by forming Zn 2+ .
• Zn is spontaneously oxidized to Zn 2+ by Cu 2+ .
• The Cu 2+ is spontaneously reduced to Cu 0 by Zn.
• Galvanic cells consist of – Anode: Zn(
s
– Cathode: Cu ) 2+ (
aq
Zn 2+ (
aq
) + 2e 2 ) + 2e Cu(
s
) – Salt bridge (used to complete the electrical circuit): cations move from anode to cathode, anions move from cathode to anode.
• The two solid metals are the electrodes (cathode and anode).
• As oxidation occurs, Zn is converted to Zn 2+ and 2e . The electrons flow towards the anode where they are used in the reduction reaction.
A shorthand convention exists for describing batteries. For the Cu/Zn battery, it would be described as follows: Zn(s)|Zn 2+ (aq)|| Cu 2+ (aq)|Cu(s) The ANODE... The CATHODE...
-supplies electrons to the external circuit (wire) -accepts electrons from the external circuit (wire) -is the negative pole of the battery -is the positive pole of the battery -is the site of OXIDATION -is the site of REDUCTION -is written on the left hand side if the convention -is written on the right hand side if the is followed convention is followed -is the half-cell with the lowest electrode potential -is the half-cell with the highest electrode potential
Cell EMF
• Electromotive force (emf) is the force required to push electrons through the external circuit.
• Cell potential:
E
cell is the emf of a cell (in Volts).
• For 1
M
solutions at 25 C (standard conditions), the standard emf (standard cell potential) is called
E
cell .
Standard Reduction Potentials
• Convenient tabulation of electrochemical data.
• Standard reduction potentials,
E
red are measured relative to the standard hydrogen electrode (SHE).
1 V
1 J 1 C
Standard reduction potentials
To predict the reactivity of oxidants or reductants we need to measure the potential of each half-reaction.
impossible!!....for every oxidation we have a reduction Define a
standard half-cell
of potential = 0V against which all other half-
cell reduction potentials
are measured. Each component in these standard cells having
unit concentration.
Pt(s) | H 2 (g) | H + (a) || Ag + (a) | Ag(s) |_______________| H +
NHE
(aq) + e H 2 (g)E 0 =0V By convention: Standard (or Normal) Hydrogen Electrode is used
The cell potential E = E+ - E- (written as reductions)
Standard Reduction Potentials • Consider Zn(
s
) Zn 2+ (
aq
) + 2e . We measure
E
cell relative to the SHE (cathode):
E
cell =
E
+ (cathode) 0.76 V = 0 V -
E
(anode)
E
red (anode).
• Therefore,
E
red (anode) = -0.76 V.
• Standard reduction potentials must be written as reduction reactions: Zn 2+ (
aq
) + 2e Zn(
s
),
E
red = -0.76 V.
• Since
E
red = -0.76 V we conclude that the reduction of Zn presence of the SHE is not spontaneous.
2+ in the
Electrochemical series
half-rxns oxidant reductant E 0 (V) 1
F
D
G = -nFE
96 , 500 C/mol 96,500 J/V.mol
Therefore if E cell is positive, the reaction is spontaneous
stronger oxidant
F 2 (g) + 2e- <=> 2F Ce 4+ + e- <=> Ce 3+ 2.890
1.720
Ag + Fe 3+ + e- <=> Ag(s) 0.799
+ e- <=> Fe 2+ 0.771
O 2 + 2H + Cu 2+ + 2e- <=> H 2 O 2 0.695
+ 2e- <=> Cu(s) 0.339
2H + Cd 2+ + 2e- <=> H 2 (g) 0.000
+ 2e- <=> Cd(s) -0.402
Zn 2+ K + + 2e- <=> Zn(s) -0.762
+ e- <=> K(s) -2.936
Li + +e- <=> Li(s) -3.040
stronger reducer
Applications of Galvanic Cells
Potentiometry and Ion Selective Electrodes
the measure of the cell potential to yield chemical information (conc., activity, charge ....) 0 .
0 1 M C 2 + C 2 + 0 .
1 M C 2 + ( 0 .
0 1 + ) M + a + + C 2 + ( 0 .
1 ) M a 2 + 0 .
0 2 M C l 0 .
2 M C l + 0 .
0 2 M C l + 0 .
2 M C l Calcium selective molecular recognition ligand A difference in the activity of an ion on either side of a selective membrane results in a thermodynamic potential difference being created across that membrane D
G
RT
ln
A
1
A
2
nFE
E
RT nF
ln
A
1
A
2 0. 05916
n
log
A
1
A
2 (
à
25C )
The glass pH electrode
Ag(s) | AgCl(s) | Cl (aq) || H (aq, ext) H (aq, int), Cl (aq) | AgCl(s) | Ag(s) ref ext analyte H + int ref int
Corrosion
Fe 2+ +2e Fe E 0 =-0.44V
2H + + 2e H 2 E 0 =0V 2H 2 O + O 2 =4e 4OH E 0 =1.23V
Iron is oxidized in water or humid conditions to give rust.
Inhibit this by coating with another material (Zn for example that forms a protective oxide on the iron), or by providing a sacrificial anode (b).
Batteries-providing electricity from chemistry
The Lead Acid Storage Battery was developed in the late 1800's and has remained the most common and durable of the battery technologies (in vehicles).
Lead-acid batteries
When the battery is used as a voltage supply, electrons flow from the Pb metal to the Pb(IV)oxide. The reactions aren't quite the reverse of the formation reactions, because now the sulfate ions in the solution begin to play a role. The two reactions are: PbO 2 + 4H + + 2e + SO 4 -2 PbSO 4 + 2H 2 O Pb + SO 4 -2 PbSO 4 + 2e The overall reaction if we combine the hydrogen ions and the sulfate: PbO 2 + Pb + 2H 2 SO4 2 PbSO 4 + 2 H 2 O Lead sulfate is fairly insoluble so that as soon as Pb(II) ions are formed by either reaction, the ions immediately precipitate as lead sulfate. The beauty is that this lead sulfate stays attached to the grids so that it is there for recharging of the battery.
Other batteries
Primary battery non rechargeable Longer shelf-life Rechargeable Offer higher efficiencies compared to burning fuels
Leclanche
Images of batteries
Alkaline Fuel Cell
Electrolysis
Use Faraday’s Laws to evaluate the number of moles of a substance oxidised or reduced by passage of charge (current over a given period of time = I.t) through an electrode Faraday:
Q (charge) = nF
N=number of moles of electrons F=constant of 96500 Coulomb/mole Example (try it): What current is needed to deposit 0.500g of chromium from a solution containing Cr 3+ over a one hour period (MW for Cr=52)? (Ans=0.77A)
Applications of Electrolytic Cells
Aluminium refining:
The major ore of aluminium is bauxite, Al 2 O 3 .
Anhydrous Al 2 O 3 melts at over 2000°C. This is too high to permit its use as a molten medium for electrolytic formation of free aluminium. The electrolytic process commercially used to produce aluminium is known as the Hall process, named after its inventor, Charles M. Hall. Al 2 O 3 is dissolved in molten cryolite, Na 3 AlF 6 , which has a melting point of 1012 o C and is an effective conductor of electric current. Graphite rods are employed as anodes and are consumed in the electrolysis process. The cell electrolytic reaction is: 2Al 2 O 3 + 3C 4Al(l) + 3CO 2 (g)
Electrolysis of brine
Chlorine and sodium hydroxide are both manufactured by electrolysis of brine (aqueous sodium chloride) using inert electrodes. Chlorine is evolved at the anode, Cl 1/2Cl 2 + e Hydrogen is evolved at the cathode: H + + e 1/2H 2 The removal of chloride ions and hydrogen ions leaves sodium ions and hydroxide ions in solution.
Chlorine is used to disinfect municipal water supplies and water in swimming pools. It is used to manufacture household bleaches and disinfectants. It is used to manufacture plastics (e.g. PVC), pesticides, anaesthetics, CFCs etc. Sodium hydroxide is used in the manufacture of synthetic fibres, soaps and detergents.
Electroplating
In all aspects of our lives we are surrounded by products with electroplated surfaces. Whether we are looking at a silver-plated watch through gold-plated glasses, watching television, using the washing machine, getting into a car or boarding a plane: electroplating plays an important part in all of these situations. The objective is to prevent corrosion and wear,produce hardness and conductivity, and give products an attractive appearance.
The principle: thin metallic layers with specific properties are deposited on base materials including steel, brass, aluminium, plastic and die-cast parts.
Silver electroplating was the first large scale use of electrolysis for coating base metal objects with a higher value decorative finish.