Electrochemistry

1
Electrochemistry

• Applications of Redox

2
Review

• Oxidation reduction reactions involve a transfer
  of electrons.
• OIL- RIG
• Oxidation Involves Loss
• Reduction Involves Gain
• LEO-GER
• Lose Electrons Oxidation
• Gain Electrons Reduction

3

• Solid lead(II) sulfide reacts with oxygen in the
  air at high temperatures to form lead(II) oxide
  and sulfur dioxide. Which substance is a
  reductant (reducing agent) and which is an
  oxidant (oxidizing agent)?  
• PbS, reductant O2, oxidant 
• PbS, reductant SO2, oxidant 
• Pb2, reductant S2- oxidant 
• PbS, reductant no oxidant 
• PbS, oxidant SO2, reductant

4
Applications

• Moving electrons is electric current.
• 8HMnO4- 5Fe2 5e- Mn2 5Fe3
  4H2O
• Helps to break the reactions into half reactions.
• 8HMnO4-5e- Mn2 4H2O
• 5(Fe2 Fe3 e- )
• In the same mixture it happens without doing
  useful work, but if separate

5

• Connected this way the reaction starts
• Stops immediately because charge builds up.

H MnO4-
Fe2
6
Galvanic Cell
Salt Bridge allows current to flow
H MnO4-
Fe2
7

• Electricity travels in a complete circuit

H MnO4-
Fe2
8

• Instead of a salt bridge

Porous Disk
H MnO4-
Fe2
9
e-
e-
e-
e-
Anode
Cathode
e-
e-
Reducing Agent
Oxidizing Agent
10
Cell Potential

• Oxidizing agent pulls the electron.
• Reducing agent pushes the electron.
• The push or pull (driving force) is called the
  cell potential Ecell
• Also called the electromotive force (emf)
• Unit is the volt(V)
• 1 joule of work/coulomb of charge
• Measured with a voltmeter

11
0.76
H2 in
Cathode
Anode
H Cl-
Zn2 SO4-2
1 M HCl
1 M ZnSO4
12
Standard Hydrogen Electrode

• This is the reference all other oxidations are
  compared to
• Eº 0
• º indicates standard states of 25ºC, 1 atm, 1
  M solutions.

H2 in
H Cl-
1 M HCl
13
Cell Potential

• Zn(s) Cu2 (aq) Zn2(aq) Cu(s)
• The total cell potential is the sum of the
  potential at each electrode.
• Eºcell EºZn Zn2 EºCu2 Cu
• We can look up reduction potentials in a table.
• One of the reactions must be reversed, so change
  it sign.

14
Cell Potential

• Determine the cell potential for a galvanic cell
  based on the redox reaction.
• Cu(s) Fe3(aq) Cu2(aq) Fe2(aq)
• Fe3(aq) e- Fe2(aq) Eº 0.77 V
• Cu2(aq)2e- Cu(s) Eº 0.34 V
• Cu(s) Cu2(aq)2e- Eº -0.34 V
• 2Fe3(aq) 2e- 2Fe2(aq) Eº 0.77 V

15
Reduction potential

• More negative Eº
• more easily electron is added
• More easily reduced
• Better oxidizing agent
• More positive Eº
• more easily electron is lost
• More easily oxidized
• Better reducing agent

16
Line Notation

• solid½Aqueous½½Aqueous½solid
• Anode on the left½½Cathode on the right
• Single line different phases.
• Double line porous disk or salt bridge.
• If all the substances on one side are aqueous, a
  platinum electrode is indicated.

17

• For the last reaction
• Cu(s)½Cu2(aq)½½Fe2(aq),Fe3(aq)½Pt(s)

Fe2
Cu2
18

• In a galvanic cell, the electrode that acts as a
  source of electrons to the solution is called the
  __________ the chemical change that occurs at
  this electrode is called________.  
• a.  cathode, oxidation  
• b.  anode, reduction  
• c.  anode, oxidation  
• d.  cathode, reduction

19

• Under standard conditions, which of the following
  is the net reaction that occurs in the cell?
• CdCd2 Cu2Cu  
• a.  Cu2 Cd ? Cu Cd2  
• b.  Cu Cd ? Cu2 Cd2  
• c.  Cu2 Cd2 ? Cu Cd  
• d.  Cu Cd 2 ? Cd Cu2 

20
Galvanic Cell

• The reaction always runs spontaneously in the
  direction that produced a positive cell
  potential.
• Four things for a complete description.
• Cell Potential
• Direction of flow
• Designation of anode and cathode
• Nature of all the components- electrodes and ions

21
Practice

• Completely describe the galvanic cell based on
  the following half-reactions under standard
  conditions.
• MnO4- 8 H 5e- Mn2 4H2O Eº1.51 V
• Fe3 3e- Fe(s) Eº0.036V

22
Potential, Work and DG

• emf potential (V) work (J) / Charge(C)
• E work done by system / charge
• E -w/q
• Charge is measured in coulombs.
• -w q E
• Faraday 96,485 C/mol e-
• q nF moles of e- x charge/mole e-
• w -qE -nFE DG

23
Potential, Work and DG

• DGº -nFEº
• if Eº gt 0, then DGº lt 0 spontaneous
• if Eºlt 0, then DGº gt 0 nonspontaneous
• In fact, reverse is spontaneous.
• Calculate DGº for the following reaction
• Cu2(aq) Fe(s) Cu(s) Fe2(aq)
• Fe2(aq) e- Fe(s) Eº 0.44 V
• Cu2(aq)2e- Cu(s) Eº 0.34 V

24
Cell Potential and Concentration

• Qualitatively - Can predict direction of change
  in E from LeChâtelier.
• 2Al(s) 3Mn2(aq) 2Al3(aq) 3Mn(s)
• Predict if Ecell will be greater or less than
  Eºcell if Al3 1.5 M and Mn2 1.0 M
• if Al3 1.0 M and Mn2 1.5M
• if Al3 1.5 M and Mn2 1.5 M

25
The Nernst Equation

• DG DGº RTln(Q)
• -nFE -nFEº RTln(Q)
• E Eº - RTln(Q) nF
• 2Al(s) 3Mn2(aq) 2Al3(aq) 3Mn(s) Eº
  0.48 V
• Always have to figure out n by balancing.
• If concentration can gives voltage, then from
  voltage we can tell concentration.

26
The Nernst Equation

• As reactions proceed concentrations of products
  increase and reactants decrease.
• Reach equilibrium where Q K and Ecell 0
• 0 Eº - RTln(K) nF
• Eº RTln(K) nF
• nF Eº ln(K) RT

27
Batteries are Galvanic Cells

• Car batteries are lead storage batteries.
• Pb PbO2 H2SO4 PbSO4(s) H2O

28
Batteries are Galvanic Cells

• Dry Cell Zn NH4 MnO2 Zn2 NH3
  H2O Mn2O3

29
Batteries are Galvanic Cells

• Alkaline Zn MnO2 ZnO Mn2O3 (in base)

30
Batteries are Galvanic Cells

• NiCad
• NiO2 Cd 2H2O Cd(OH)2 Ni(OH)2

31
Corrosion

• Rusting - spontaneous oxidation.
• Most structural metals have reduction potentials
  that are less positive than O2 .
• Fe Fe2 2e- Eº 0.44 V
• O2 2H2O 4e- 4OH- Eº 0.40 V
• Fe2 O2 H2O Fe2O3 H
• Reactions happens in two places.

32
Salt speeds up process by increasing conductivity
Water
Fe2 O2 2H2O Fe2O3 8 H
33
Preventing Corrosion

• Coating to keep out air and water.
• Galvanizing - Putting on a zinc coat
• Has a lower reduction potential, so it is more
  easily oxidized.
• Alloying with metals that form oxide coats.
• Cathodic Protection - Attaching large pieces of
  an active metal like magnesium that get oxidized
  instead.

34
Electrolysis

• Running a galvanic cell backwards.
• Put a voltage bigger than the potential and
  reverse the direction of the redox reaction.
• Used for electroplating.

35
1.10
e-
e-
Zn
Cu
1.0 M Cu2
1.0 M Zn2
Cathode
Anode
36
A battery gt1.10V
e-
e-
Zn
Cu
1.0 M Cu2
1.0 M Zn2
Cathode
Anode
37
Calculating plating

• Have to count charge.
• Measure current I (in amperes)
• 1 amp 1 coulomb of charge per second
• q I x t
• q/nF moles of metal
• Mass of plated metal
• How long must 5.00 amp current be applied to
  produce 15.5 g of Ag from Ag

38
Calculating plating

• Current x time charge
• Charge /Faraday mole of e-
• Mol of e- to mole of element or compound
• Mole to grams of compound
• Or the reverse if you want time to plate

39

• Calculate the mass of copper which can be
  deposited by the passage of 12.0 A for 25.0 min
  through a solution of copper(II) sulfate.

40

• How long would it take to plate 5.00 g Fe from an
  aqueous solution of Fe(NO3)3 at a current of 2.00
  A?

41
Other uses

• Electrolysis of water.
• Separating mixtures of ions.
• More positive reduction potential means the
  reaction proceeds forward.
• We want the reverse.
• Most negative reduction potential is easiest to
  plate out of solution.

42
Redox

• Know the table
• 2. Recognized by change in oxidation state.
• 3. Added acid
• 4. Use the reduction potential table on the front
  cover.
• 5. Redox can replace. (single replacement)

43

• 6. Combination Oxidizing agent of one element
  will react with the reducing agent of the same
  element to produce the free element.
• I- IO3- H I2 H2O
• 7. Decomposition.
• a) peroxides to oxides
• b) Chlorates to chlorides
• c) Electrolysis into elements.
• d) carbonates to oxides

44
Examples

1. A piece of solid bismuth is heated strongly in
   oxygen.
2. A strip or copper metal is added to a
   concentrated solution of sulfuric acid.
3. Dilute hydrochloric acid is added to a solution
   of potassium carbonate.

45

1. Hydrogen peroxide solution is added to a solution
   of iron (II) sulfate.
2. Propanol is burned completely in air.
3. A piece of lithium metal is dropped into a
   container of nitrogen gas.
4. Chlorine gas is bubbled into a solution of
   potassium iodide.

46
Examples

1. A stream of chlorine gas is passed through a
   solution of cold, dilute sodium hydroxide.
2. A solution of tin ( II ) chloride is added to an
   acidified solution of potassium permanganate
3. A solution of potassium iodide is added to an
   acidified solution of potassium dichromate.

47

1. Magnesium metal is burned in nitrogen gas.
2. Lead foil is immersed in silver nitrate solution.
3. Magnesium turnings are added to a solution of
   iron (III) chloride.
4. Pellets of lead are dropped into hot sulfuric
   acid
5. Powdered Iron is added to a solution of iron(III)
   sulfate.

48
A way to remember

• An Ox anode is where oxidation occurs
• Red Cat Reduction occurs at cathode
• Galvanic cell- spontaneous- anode is negative
• Electrolytic cell- voltage applied to make anode
  positive

49
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50

• A student places a copper electrode in a 1 M
  solution of CuSO4 and in another beaker places a
  silver electrode in a 1 M solution of AgNO3. A
  salt bridge composed of Na2SO4 connects the two
  beakers. The voltage measured across the
  electrodes is found to be 0.42 volt.
• (a) Draw a diagram of this cell.
• (b) Describe what is happening at the cathode
  (Include any equations that may be useful.)

51

• A student places a copper electrode in a 1 M
  solution of CuSO4 and in another beaker places a
  silver electrode in a 1 M solution of AgNO3. A
  salt bridge composed of Na2SO4 connects the two
  beakers. The voltage measured across the
  electrodes is found to be 0.42 volt.
• (c) Describe what is happening at the anode.
  (Include any equations that may be useful.)

52

• A student places a copper electrode in a 1 M
  solution of CuSO4 and in another beaker places a
  silver electrode in a 1 M solution of AgNO3. A
  salt bridge composed of Na2SO4 connects the two
  beakers. The voltage measured across the
  electrodes is found to be 0.42 volt.
• (d) Write the balanced overall cell equation.
• (e) Write the standard cell notation.

53

• A student places a copper electrode in a 1 M
  solution of CuSO4 and in another beaker places a
  silver electrode in a 1 M solution of AgNO3. A
  salt bridge composed of Na2SO4 connects the two
  beakers. The voltage measured across the
  electrodes is found to be 0.42 volt.
• (f) The student adds 4 M ammonia to the copper
  sulfate solution, producing the complex ion
  Cu(NH3) (aq). The student remeasures the cell
  potential and discovers the voltage to be 0.88
  volt. What is the Cu2 (aq) concentration in the
  cell after the ammonia has been added?