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AP Chemistry Chapter 20 Notes. Electrochemistry. Applications of Redox. 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. Applications.

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Electrochemistry


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electrochemistry

AP Chemistry

Chapter 20 Notes

Electrochemistry

Applications of Redox

review
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
applications
Applications
  • Moving electrons is electric current.
  • 8H++MnO4-+ 5Fe+2 +5e-® Mn+2 + 5Fe+3 +4H2O
  • Helps to break the reactions into half rxns.
  • 8H++MnO4-+5e-® Mn+2 +4H2O
  • 5Fe+2® 5Fe+3 + 5e- )
  • In the same mixture it happens without doing useful work, but if separate
slide4
Connected this way the reaction starts
  • Stops immediately because charge builds up.

H+

MnO4-

Fe+2

galvanic cell
Galvanic Cell

Salt Bridge allows current to flow

H+

MnO4-

Fe+2

slide6

e-

  • Electricity travels in a complete circuit
  • Instead of a salt bridge

H+

MnO4-

Fe+2

slide7

Porous Disk

H+

MnO4-

Fe+2

slide8

e-

e-

e-

e-

Anode

Cathode

e-

e-

Reducing Agent

Oxidizing Agent

cell potential
Cell Potential
  • Oxidizing agent pushes the electron.
  • Reducing agent pulls 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
slide10

0.76

H2 in

Cathode

Anode

H+ Cl-

Zn+2 SO4-2

1 M ZnSO4

1 M HCl

standard hydrogen electrode
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

cell potential1
Cell Potential
  • Zn(s) + Cu+2 (aq)® Zn+2(aq) + Cu(s)
  • The total cell potential is the sum of the potential at each electrode.
  • Eºcell = EºZn® Zn+2 + EºCu+2® Cu
  • We can look up reduction potentials in a table.
  • One of the reactions must be reversed, so change it sign.
cell potential2
Cell Potential
  • Determine the cell potential for a galvanic cell based on the redox reaction.
  • Cu(s) + Fe+3(aq)® Cu+2(aq) + Fe+2(aq)
  • Fe+3(aq)+ e-® Fe+2(aq) Eº = 0.77 V
  • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V
  • Cu(s) ® Cu+2(aq)+2e-Eº = -0.34 V
  • 2Fe+3(aq)+ 2e-® 2Fe+2(aq) Eº = 0.77 V
line notation
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.
  • For the last reaction
  • Cu(s)½Cu+2(aq)½½Fe+2(aq),Fe+3(aq)½Pt(s)
galvanic cell1
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
practice
Practice
  • Completely describe the galvanic cell based on the following half-reactions under standard conditions.
  • MnO4- + 8 H+ +5e-® Mn+2 + 4H2O Eº=1.51
  • Fe+3 +3e-® Fe(s) Eº=0.036V
potential work and d g
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 = qE
  • Faraday = 96,485 C/mol e-
  • q = nF = moles of e- x charge/mole e-
  • w = -qE = -nFE = DG
potential work and d g1
Potential, Work and DG
  • DGº = -nFE º
  • if E º < 0, then DGº > 0 spontaneous
  • if E º > 0, then DGº < 0 nonspontaneous
  • In fact, reverse is spontaneous.
  • Calculate DGº for the following reaction:
  • Cu+2(aq)+ Fe(s) ® Cu(s)+ Fe+2(aq)
  • Fe+2(aq)+ e-® Fe(s) Eº = 0.44 V
  • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V
cell potential and concentration
Cell Potential and Concentration
  • Qualitatively - Can predict direction of change in E from LeChâtelier.
  • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s)
  • Predict if Ecell will be greater or less than Eºcell if [Al+3] = 1.5 M and [Mn+2] = 1.0 M
  • if [Al+3] = 1.0 M and [Mn+2] = 1.5M
  • if [Al+3] = 1.5 M and [Mn+2] = 1.5 M
the nernst equation
The Nernst Equation
  • DG = DGº +RTln(Q)
  • -nFE = -nFEº + RTln(Q)
  • E = Eº - RTln(Q) nF
  • 2Al(s) + 3Mn+2(aq) ® 2Al+3(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.
the nernst equation1
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
  • nFEº = ln(K) RT
batteries are galvanic cells
Batteries are Galvanic Cells
  • Car batteries are lead storage batteries.
  • Pb +PbO2 +H2SO4®PbSO4(s) +H2O
  • Dry Cell Zn + NH4+ +MnO2 ® Zn+2 + NH3 + H2O
  • Alkaline Zn +MnO2 ® ZnO+ Mn2O3 (in base)
  • NiCad
  • NiO2 + Cd + 2H2O ® Cd(OH)2 +Ni(OH)2
corrosion
Corrosion
  • Rusting - spontaneous oxidation.
  • Most structural metals have reduction potentials that are less positive than O2 .
  • Fe ® Fe+2+2e-Eº= 0.44 V
  • O2 + 2H2O + 4e- ® 4OH- Eº= 0.40 V
  • Fe+2 + O2 + H2O ® Fe2 O3 + H+
  • Reaction happens in two places.
slide24

Salt speeds up process by increasing conductivity

Water

Rust

e-

Iron Dissolves- Fe ® Fe+2

preventing corrosion
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.
electrolysis
Electrolysis
  • Running a galvanic cell backwards.
  • Put a voltage bigger than the potential and reverse the direction of the redox reaction.
  • Used for electroplating.
slide27

1.10

e-

e-

Zn

Cu

1.0 M Cu+2

1.0 M Zn+2

Cathode

Anode

slide28

A battery >1.10V

e-

e-

Zn

Cu

1.0 M Cu+2

1.0 M Zn+2

Cathode

Anode

calculating plating
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+
other uses
Other uses
  • Electroysis of water.
  • Seperating 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.
slide31

Balancing

Redox

Equations

slide32

2. in base

Am3+(aq) + S2O82-(aq) ---->

AmO2+(aq) + SO42-(aq)

slide34

4. Bi(OH)3 + SnO22-

Bi(s) + SnO32-

slide36

Electrolytic Cell

a cell that uses electrical energy to produce a chemical change that would otherwise NOT occur spontaneously

slide38

(+)

(-)

M+(aq)

M

M

X-(aq)

slide39

e-

e-

(+)

(-)

M+(aq)

M

M

X-(aq)

Anode

M M+ + e-

oxidation

Cathode

M+ + e- M

reduction

slide40

Ampere

a unit of electrical current equal to one coulomb of charge per second

coul

sec

1 amp = 1

slide41

Coulomb

a unit of electric charge equal to the quantity of charge in about 6 x 1019 electrons

slide42

Faraday

a constant representing the charge on one mole of electrons

1 F = 96,485 C

96,500 C

slide43

3: It is necessary to replate a silver teapot with 15.0 g of silver. If the electrolytic cell runs at 2.00 amps, how long will it take to plate the teapot?

slide44

4: Sodium metal and chlorine gas are prepared industrially in a Down’s Cell from the electrolysis of molten NaCl. What mass of metal and volume of gas can be made per day if the cell operates at 7.0 volts and 4.0 x 104 amps if the cell is 75% efficient?

slide45

5: At what current must a cell be run in order to produce 5.0 kg of aluminum in 8.0 hours if the cell produces solid aluminum from molten aluminum chloride?

slide46

ELECTROCHEMISTRY,

FREE ENERGY,

& EQUILIBRIUM

slide48

but: wmax = G

and q = nF

thus if: wmax = - q . Emax

then G = - nFE

slide49

G = G0 + RT ln Q

G = - nFE

- nFE = - nFE0 + RT ln Q

slide52

IF T = 250C = 298.15 K

ln Q = 2.303 log Q

R = 8.314 J/mol.K

F = 96,485 C/mol

slide53

what if : Q = Keq ?

then: E = 0.0 V

slide57

7a: Calculate the standard free energy for the cell:

Cr(s) Cr3+ (1M) Fe2+ (1M) Fe(s)

7b: What will be the voltage if [Fe2+] = 0.50M and [Cr3+] = 0.30M at 200C?

slide58

8: Through electrochemical calculations, determine the Ksp for silver bromide.

AgBr + e- Ag + Br-

E0 = 0.10 V

slide60

Review

Oxidation: loss of e-

[increase in ox #]

[reducing agent]

Reduction: gain of e-

[decrease in ox #]

[oxidizing agent]

slide61

Reduction Potential

The ease with which a chemical species can be reduced

slide62

Standard Reduction

Potential

Appendix M

Table 20.1 in text

slide65

2a. Choosing from among the reactants in the given half reactions, identify the strongest and weakest oxidizing agents.

anode and cathode
Anode and Cathode
  • OXIDATION occurs at the ANODE.
  • REDuction occurs at the CAThode.
slide67

Electrochemical Cell

device in which chemical energy is spontaneously changed to electrical energy

slide68

battery

voltaic cell

galvanic cell

slide70

M2+(aq)

M1+(aq)

M1

X-(aq)

X-(aq)

M2

slide71

M2+(aq)

M1+(aq)

M1

X-(aq)

X-(aq)

M2

Anode

M1 M1+ + e-

Cathode

M2+ + e- M2

slide72

K+(aq) NO3-(aq)

M2+(aq)

M1+(aq)

M1

X-(aq)

X-(aq)

M2

Cathode

M2+ + e- M2

Anode

M1 M1+ + e-

slide73

e- flow is from source of high “concentration” to source of low “concentration”

slide74

e-

e-

K+(aq) NO3-(aq)

M2+(aq)

M1+(aq)

M1

X-(aq)

X-(aq)

M2

Cathode

M2+ + e- M2

Anode

M1 M1+ + e-

slide75

shorthand notation

oxidation reduction

M1 |M1+ || M2+ |M2

anode  cathode

 e- flow 

slide77

Electrochemical Standard State Conditions

[ions] = 1 M

T = 250C

Pgas = 1 atm

slide78

An electrochemical cell is spontaneous if:

Oxidation-reduction occurs

Ered + Eox > 0

slide81

Line Notation:

ANODE CATHODE

Ni(s)|Ni2+ (aq, 1 M)||Au3+(aq, 1 M)|Au(s)

oxidation reduction

slide82

Line Notation:

ANODE CATHODE

Al(s) | Al3+(aq, 1 M) || Ni2+(aq, 1 M) | Ni (s)

oxidation reduction