reduction and oxidation reactions and oxidation reactions • predict what might happen when a piece...
TRANSCRIPT
Chapter 12
Electrochemical cells and
electrode potentials
Reduction and Oxidation Reactions
• Predict what might happen when a piece of copper wire in a solution of 2% AgNO3.
• If you try this experiment, you will initially see that the copper is a shiny copper color and the solution is clear. In less than one hour the solution is light blue and the wire is covered with shiny silver needles. What happened?
• Copper metal became copper ions in solution and silver ions became silver metal.
• Cu(s) + Ag+(aq) Cu2+
(aq) + Ag(s) (unbalanced)
• The Cu(s) loses electrons to become Cu2+(aq) ions and the Ag+
(aq) ions gain electrons to become Ag(s).
• Reactions that involve the exchange of electrons are called reduction and oxidation (redox) reactions.
• When a chemical species loses electrons we say that it is oxidized, and when a chemical species gains electrons we say that it is reduced.
• The Cu(s) loses electrons to be oxidized to Cu2+(aq).
The Ag+(aq) gain electrons to be reduced to Ag(s).
• We can break the reaction into the following two half reactions. – Cu(s) Cu2+
(aq) + 2e-
– Ag+(aq) + e- Ag(s)
• What would you predict if you placed a piece of Ag metal in a solution of Cu2+?
• Since we observed that the reaction of Ag+ and Cu is spontaneous, we would not expect the reverse reaction to be spontaneous.
• So no reaction occurs between Ag metal and Cu2+. Or we call it nonspontaneous reaction
Terminology
• Reduction: gaining of electrons (decrease in the oxidation
state
• Oxidation: loss of electrons (increase in the oxidation
state)
• Reducing agent (reductant): species that donates
electrons to reduce another reagent.
(The reducing agent get oxidized.)
• Oxidizing agent (oxidant): species that accepts electrons
to oxidize another species.
(The oxidizing agent gets reduced.)
Oxidation-reduction reaction (redox reaction): a reaction in which electrons are transferred from one reactant to another. • For example, the reduction of cerium(IV) by iron(II):
Ce4+ + Fe2+ Ce3+ + Fe3+
• The reduction half-reaction is given by...
Ce4+ + e- Ce3+
• The oxidation half-reaction is given by...
Fe2+ e- + Fe3+
• The half-reactions are the overall reaction broken down into
oxidation and reduction steps. • Half-reactions cannot occur independently, but are used
conceptually to simplify understanding and balancing the equations
Rules for Balancing Oxidation-Reduction Reactions
• Write out half-reaction "skeletons."
• Balance the half-reactions by adding H+, OH- or H2O as needed, maintaining electrical neutrality.
• Combine the two half-reactions such that the number of electrons transferred in each cancels out when combined.
• For example, consider the following reaction of the peroxydisulfate ion with manganese ion:
• S2O82- + Mn2+ SO4
2- + MnO4-
1. The reduction step is:
S2O8
2- + 2 e- 2 SO42-
(Each sulfur atoms goes from +7 to +6 oxidation state.)
2. The oxidation step is:
Mn2+ + 4 H2O 5 e- + MnO4- + 8 H+
(Manganese(II) loses 5 electrons, going from +2 to +7.)
3. In combining the two equations, the oxidation step must be multiplied by "2," and the reduction step must be multiplied by "5" to cancel out the electrons:
2x[Mn2+ + 4 H2O 5 e- + MnO4- + 8 H+]
5x[S2O82- + 2 e- 2 SO4
2-]
4. Adding these two equations together:
S2O82- + 2 Mn2+ + 8 H2O 10 SO4
2- + 2 MnO4- + 16 H+
5. Note that the half-reactions are charge-balanced before adding them together.
Consider the next example:
Cr2O72- + I- Cr3+ + I3
-
• The reduction step is given by...
Cr2O72- + 6 e- + 14 H+ 2 Cr3+ + 7 H2O
(Cr(IV) in the dichromate on is reduced to Cr(III).)
• The oxidation step is give by...
3 I- 2 e- + I3-
• Multiplying the oxidation half-reaction by 3x and adding the two half-reactions together:
Cr2O72- + 9 I- + 14 H+ 2 Cr3+ + 3 I3
- + 7 H2O
Oxidation-Reduction Reactions in Electrochemical Cells
• It is possible to separate the half-reactions of an oxidation-reduction reaction in an electrochemical cell.
• Consider the following reaction:
2 Ag+ + Cu(s) 2 Ag(s) + Cu2+
• The reduction half-reaction is given by...
Ag+ + e- Ag(s)
• The oxidation half-reaction is given by...
Cu(s) 2 e- + Cu2+
Redox reactions can be "separated" in a galvanic cell (also
called a voltaic cell or battery):
Charactristics of the of the galvanic cell (battery)
• The left container (half cell) contains of 0.020 M CuSO4.
• The right cell contains 0.0200 M AgNO3.
• A Cu electrode is immersed in the CuSO4 solution.
• An Ag electrode in immersed in the AgNO3 solution.
• The two solutions "communicate" via a salt bridge which consists of a saturated KCl solution in a tube
• with glass frits in both ends.
• At the anode, oxidation takes place:
Cu(s) 2 e- + Cu2+
• At the cathode, reduction takes place:
2 Ag+ + 2 e- 2 Ag(s)
• Chloride ions move into the CuSO4 solution to maintain electrical neutrality.
• Potassium ions move into the AgNO3 solution to maintain electrical neutrality.
• The volt meter reads 0.412 V at the instant the connection is made between the two electrodes. (This represents the difference is voltage ( Ecell) between the two electrodes.
a. The copper electrode has an initial voltage of 0.2867 V.
b. The silver electrode has an initial voltage of 0.6984 V.
c. Voltage Difference = Voltagecathode - Voltageanode
Voltage Difference = 0.6984 - 0.2867 = + 0.412 V
• This initial voltage drops as the reaction proceeds toward equilibrium as soon as connection is made. At equilibrium, the voltage read zero volts.
• The potential difference (voltage) between the anode and cathode is a measure of the tendency of the reaction to proceed from nonequilibrium to equilibrium
Operation of an electrolytic cell
• An external voltage source with a voltage larger than that of the battery is connected to the galvanic cell - positive pole to the silver electrode; negative pole to the copper electrode.
• The external voltage source reversed the direction of electron flow (and reverses the direction of the reactions at each electrode): – The silver electrode switches from the cathode to
the anode (the site of oxidation).
– The copper electrode switches from the anode to the cathode (the site of reduction.
– The voltage meter will read "negative" if still connected in the original fashion.
– The battery will "recharge.
Schematic Representation of Cells
• The copper(II) sulfate/silver nitrate systems described above would be symbolized:
• Cu|CuSO4 (0.02 M)||AgNO3 (0.02 M)|Ag – Each vertical line ("|") represents a phase
boundary or interface where a potential develops.
– Each double vertical line ("||") represents two phase boundaries (e.g., the salt bridge).
• The direction of electron flow is from left to
right: Cu--->Ag.
The original, initial voltage for the cell
• The original, initial voltage for the cell is given by...
• Ecell = Eright - Eleft = Ecathode - Eanode = EAg - ECu
• The potentials (absolute voltages) at the two electrodes cannot be determined experimentally.
• Only the differences between electrodes (via a volt meter) can be measured.
• Potentials at electrodes are assigned relative values, based upon comparison to a standard.
• By convention, the standard hydrogen electrode (SHE) is used as the "agreed upon" reference half-cell against which all others are compared.
• The voltage or cell potential is related to the free energy of the reaction driving the cell:
G = -nFEcell = -2.303RTlog(Keq)
Standard Hydrogen Electrode (SHE)
• The SHE (standard hydrogen electrode) is the reference point for
determining relative electrode potentials.
• The SHE is symbolized:
Pt, H2 ( hydrogen = 1.00 atm)|H+ (aH+ = 1.00)||Ag+ (aAg+ = 1.00)|Ag)
The platinum electrode is a platinized platinum electrode (platinum
coated with finely divided platinum called platinum black).
The aqueous acid solution has an activity of 1.00 (i.e., approximately
1.00 M hydrogen ion) and is saturated with hydrogen gas, bubbled in
at 1.00 atmosphere.
The SHE is connected via the salt bridge and connecting wire to the
other half-cell of the battery.
The half-cell reaction is given by: 2 H1+ + 2 e- H2(g)
• The SHE acts as either the anode or cathode, depending upon whether electrons are given up to the other half-cell or taken in.
• By convention, the SHE is assigned an absolute potential (voltage) = 0.00 volts at ALL temperatures.
• All other half-cells (and half-reactions) are measured relative to the SHE.
• By definition, the electrode potential is the potential of a cell with the standard hydrogen electrode acting as the anode and the other half-cell acting as the cathode.
– For example, consider the cell formed with the SHE as one half-cell and an Ag/Ag+ half-cell on the other side consisting of a silver electrode and an aAg+ = 1.00 silver nitrate solution.
• At the anode (oxidation): SHE
2 H+ + 2 e- H2(g)
• At the cathode (reduction): Ag/Ag+
Ag+ + e- Ag(s)
• As the instant the connection between half-cells is made, the experimentally observed difference in potential between the two electrodes measured by the volt meter (i.e., the battery voltage) is +0.799 V.
• Electrons flow from the SHE to the Ag electrode.
• By convention, the sign of the voltage is positive if electrons leaves the SHE; negative if electrons are taken in.
•
• Consider what happens if the Ag/Ag1+ half-cell is replaced with a Cd/Cd2+ half-cell (aCd+2+ = 1.00) attached to an SHE.
• 1) At the anode (oxidation): Cd/Cd2+
• Cd Cd2+ + 2 e1-
• 2) At the cathode (reduction): SHE
• 2 H1+ + 2 e1- H2(g)
• 3) At the instant the connection between half-cells is made, the experimentally observed difference in potential between the two electrodes measured by the volt meter (i.e., the battery voltage) is a negative voltage (-0.403 V).
• 4) Electrons flow to the SHE from the Cd electrode.
• 5) By convention, the sign of the voltage is negative if electrons are taken in by the SHE. This means that the half-cell reaction opposite the SHE is more reducing than the SHE.
• 6) A voltage greater than 0.403 V would have to be applied to reverse the flow of electrons.
Accepted IUPAC Conventions and Electrode Potentials
• Half-reactions are always written as reductions (i.e., electrodes potentials are by definitions reduction potentials).
• The sign of the electrode potential is determined relative to the SHE. – Positive (+) means electrons flow out of the SHE to the other
electrode (i.e., the SHE acts an the anode).
– Negative (-) means electrons flow to the SHE from the other electrode (i.e., the SHE becomes the cathode).
• The sign of the electrode potential signifies whether the net reaction of the battery is spontaneous "to the right" or "to the left." For example: – The Ag/Ag+ half-reaction has an electrode potential of +0.799 V,
meaning electrons flow to the Ag electrode.
– The Cd/Cd2+ half-reaction has an electrode potential of -0.403 V, meaning electrons flow to the SHE (and the reactions occur opposite to the directions written above).
• The effect of species concentration of electrode potentials is
described by the Nernst equation. For the reversible half-
reaction:
aA + bB + ... ne1- cC + dD + ...
Nernst equation
where...
E = the electrode potential
Eo= the standard electrode potential (i.e., the potential observed
when species in the half-cell are at a = 1.00 or pressure = 1.00 atm)
R = the gas constant (8.314 JK-1mol-1)
T = Kelvin temperature
n = number of moles of electrons in balanced half reaction
F = Faraday's constant (96,485 coulombs/mole; the charge on a
mole of electrons).
• Note that the standard electrode potential (Eo) is
measured under standard conditions (a = 1, pressure
= 1 atm).
• The Nernst equation corrects for nonstandard
concentrations.
• 2.303RT/nF simplifies to 0.0592/n.
• E = Eo if all species are at a = 1 and pressure = 1 atm.
Calculating battery voltages
• The battery voltage is given by:
Ecell = Eright - Eleft = Ecathode - Eanode
• Each half-cell electrode potential is
calculated, and the difference between
electrode determined.
Applications of electrode potentials
• Redox titrations utilized to measure analytes
via oxidation-reduction titrations will be
discussed. This will include:
• Equilibrium constants for redox reactions.
• Titrations curves - shapes, endpoints,
calculations, etc.
• Indicators for oxidation-reduction reactions.
Measuring Equilibrium Constants with Redox
Reactions • Consider the reaction:
2 Ag+ + Cu(s) 2 Ag(s) + Cu2+
• with associated half-reactions:
2 Ag+ + 2 e- 2 Ag(s) (reduction step)
Cu(s) 2 e- + Cu2+ (oxidation step)
• The equilibrium constant for this reaction is given by: •
•The reaction above is the same as the galvanic cell...
Cu|Cu2+ (?? M)||Ag+ (?? M)|Ag
• When this galvanic cell completely discharges to reach equilibrium, Ecell = Eright - Eleft = Ecathode - Eanode = EAg - ECu = 0
which means: EAg = ECu
• Substituting for each from the Nernst equation:
Note that Keq is so large,
the equilibrium lies heavily
toward the right
Oxidation-Reduction Titrations
Oxidation-Reduction Titrations
• When a oxidation-reduciton (redox) reaction is used to measure
an analyte, the titrations usually follow the electrode potential
as a function of titrant (or analyte) concentration.
• Since the electrode potential is a log function of the
concentrations, it behaves as a "p" function.
• We will discuss the redox titrations in terms of a
specific reaction and look at the calculations for
each of the different regions of the titration curve.
This would include:
• Initially (before any titrant was added)
• Before the equivalence point
• At the equivalence point
• After the equivalence point
• Consider titration of iron(II) with cerium(IV)
Fe2+ + Ce4+ Fe3+ + Ce3+
• Oxidation half-reaction:
Fe2+ Fe3+ + e-
• Reduction half-reaction:
e- + Ce4+ Ce3+ – The reactions must be fully reversible (i.e., the system must be at
equilibrium at all times through the titration).
– To be fully reversible means that:
Esystem = ECe = EFe = Eindicator and Ecell = 0 • The equilibrium concentration ratios of the oxidized and reduced forms
of the two species are such that their attraction for electrons are identical.
• The Nernst equation applies.
– Data from the titration can be used to calculate the titration curve using the Nernst equation for either the cerium(IV) or the iron(II) half-reactions.
• In practice: – Before equivalence, the Fe(III) and Fe(II) concentrations are used to
calculate the Esystem.
– After equivalence, Ce(III) and Ce(IV) concentrations are used to calculate the Esystem.
– At equivalence, simplifying assumptions are made based upon the stoichiometric relationships to calculate the Esystem.
• Example: 50.00 mL of 0.05000 M Fe2+ is titrated with 0.1000 M
Ce4+ in a medium containing 1.0 M H2SO4
• Derivation/calculation of the titration curve
1. Initial region: no Ce(IV) has been added
• Since the amount of Fe3+ cannot be calculated (and is
essentially zero), and the amount of Ce3+ is zero, the potential
cannot be calculated.
• Some amount of reaction must occur before numbers can be
plugged into the Nernst equation.
2. Before equivalence
assume 5.00 mL of Ce4+ solution have been added
3. At equivalence
• All the iron(II) is converted to iron(III); all the cerium(IV) is converted to cerium(III)
– Note that at the equivalence point: [Fe3+] = [Ce3+]
[Fe2+] = [Ce4+] ==> a very small amout!!
– The system potential can be calculated by combining the Nernst equations for both species as follows:
4. Beyond the equivalence point
• assume 25.10 mL of cerium(IV) have been added – The iron(II) is completely titrated. b. The cerium(III) and
cerium(IV) concentrations must be used to to calculate the potential of the system.
– Note that the number of moles of cerium(III) is equal to the number of moles of iron(II) originally present in the sample.
– Cerium(IV) present is due to the excess cerium(IV) added past equivalence.
– The system potential is calculated as follows:
Substituting into the Nernst equation:
Summary Of Redox Titration Curve
1. Titration begins beyond Eo for reactant
2. 50% titrated E = Eo for reactant
3. EP
200% titrated E = Eo2
Oxidation-Reduction Indicators • Two types of indicators are commonly used:
1. Specific indicators.
2. Oxidation-reduction indicators ("true" redox indicators).
• Specific indicators are species that react with one of the titration species to produce a color change.
• "True" redox indicators respond to the system potential to produce a color change. – Redox indicators are more versatile than specific indicators, since
they depend on the system potential, not the specific reaction.
– The half-reaction for the redox indicator is:
In(oxid) + n e- In(red)
• Like pH indicators, the color of redox indicators must change by about 10-fold to be seen. This means to see a full color transition, ratio of the oxidized and reduced indicator species must change from 1:10 to 10:1. – If these ratios are plugged into the Nernst equations for various
"n" values, the change in potential at the endpoint should correspond to EIn
o +/- 0.0592/n to get an appropriate color change. • For n = 1, the range around the endpoint is EIn
o +/- 0.0592/1 = +/- 0.0592 V (meaning the equivalence point must span this "window" for the indicator to work).
• For n = 2, the range around the endpoint is +/- 0.0296 V.
• For n = 3, the range around the endpoint is +/- 0.0197 V
– In practice, the equivalence point of a titration is first calculated, then a redox indicator whose EIn
o most closely matches the equivalence point is chosen.
– Consideration must also be given to the titration "break" at endpoint. If the standard electrode potentials of the analyte and titrant are too close to each other (<~0.40 V), it is nearly impossible to titrate the species. (The reaction is not complete enough to give a clear endpoint transition, and an indicator can not be "fit" to the titration.)
• Examples of Specific Indicators
• Starch-I3- complex
• In the presence of I3-, starch forms a deep blue/black color due
to the formation of a complex between the starch and the I3-.
I2(s) + I-(aq) I3-(aq)
I3-(aq) + Starch I3
-(aq)
.Starch (blue color)
• The presence/absence of I3-(aq) in solution can be used as an
indicator. For example, consider the following titration of copper(II):
2 Cu2+ + 5 I- 2 CuI(s) + I3-
– If excess iodide is added, the reaction is driven to completion, and
the I31- is back titrated with S2O3
2- :
I3
- + 2 S2O32- 3 I- + S4O6
2- (tetrathionate ion)
– I3- appears in solution as a yellowish color, and is
visible to approximately 5 x 10-6 M concentration.
– During back-titration, as the yellow disappears, starch may be added to produce a blue color that disappears when the last I3
1- is titrated (i.e., an endpoint).
– The blue color of the I3-(aq)
.starch indicator is visible to approximately
2 x 10-7 M concentration.
Thiocyanate.Iron(III) Complex
• Thiocyanate reacts with iron(III) to produce a deep,
red color:
• Fe3+ + SCN- FeSCN2+ (red)
• The red color can be used to detect the presence of
Fe3+ in titrations.
Example of redox indicators
1,10-phenanthrolines
Applications
Oxidizing Agents
• KMnO4 - titrant for Fe, Sn, and Oxalate
• Ce4+ - Very stable in strong sulfuric
acid, Titration of Fe and organics
• K2Cr2O7 - little used now due to hazards
and waste disposal of Cr
• I2 - add KI to increase solubility - Vit C,
wide range of applications
• Karl Fischer Detn of H2O
Reducing Agents
• Not too many of these. Fe2+ not very stable
Fe(NH3)2(SO4)2 Iodide
Thiosulfate - S2O32- ===> S4O6
2- + 2e-
• Many common REDOX procedures use a
reducing agent to reduce the sample and
then use an oxidizing agent titrant for
analysis rather than a reducing agent titrant.
eg. reduce Fe3+ to Fe2+ then titrate with Ce.