Electrochemistry

Chapter 19

산화환원 반응과 전기화학

1. 산화환원 반응과 산화수

2. 갈바니 전지

3. 표준환원전위

4. 산화환원 반응의 자발성. 전위차와 자유에너지

5. 전위차와 농도, Nernst 방정식

6. 전지

7. 부식

8. 전기분해

9. 전기야금

2Mg (s) + O2 (g) 2MgO (s)

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Oxidation half-reaction (lose e-)

Reduction half-reaction (gain e-)

전기화학 반응

oxidation-reduction reactions in which:

• 자유에너지의 감소 => electrical energy : 갈바니 전지

• electrical energy => 자유에너지의 증가 : 전해 전지

0 0 2+ 2-

갈바니 전지(Galvanic Cells)

spontaneous

redox reaction

anode

oxidation

cathode

reduction

+ -

H+

MnO4-

Fe+2

Galvanic Cell

Salt

Bridge

allows

current

to flow

H+

MnO4-

Fe+2 e-

Electricity travels in a complete circuit

H+

MnO4-

Fe+2

Porous

Disk

Instead of a salt bridge

../../Desktop/simulations/voltaicCell20.html

Reducing

Agent

Oxidizing

Agent

e-

e-

e- e-

e-

e-

Anode Cathode

The difference in electrical

potential between the anode

and cathode is called:

• cell voltage

• electromotive force (emf)

• cell potential

Cell Diagram

Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)

[Cu2+] = 1 M & [Zn2+] = 1 M

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)

anode cathode

갈바니 전지(Galvanic Cells)

산화 환원

Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

2e- + 2H+ (1 M) H2 (1 atm)

Zn (s) Zn2+ (1 M) + 2e- Anode (oxidation):

Cathode (reduction):

Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

Anode: 산화극

Cathode: 환원극

+ -

표준전극전위(Standard Electrode Potentials)

Standard reduction potential (E0) is the voltage associated

with a reduction reaction at an electrode when all solutes

are 1 M and all gases are at 1 atm.

E0 = 0 V

Standard hydrogen electrode (SHE)

2e- + 2H+ (1 M) H2 (1 atm)

Reduction Reaction

표준전극전위(Standard Electrode Potentials)

E0 = 0.76 V cell

Standard emf (E0 ) cell

0.76 V = 0 - EZn /Zn 0

2+

EZn /Zn = -0.76 V 0

2+

Zn2+ (1 M) + 2e- Zn E0 = -0.76 V

E0 = EH /H - EZn /Zn cell 0 0

+ 2+ 2

E0 = Ecathode - Eanode cell 0 0

Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

표준전극전위(Standard Electrode Potentials)

Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)

2e- + Cu2+ (1 M) Cu (s)

H2 (1 atm) 2H+ (1 M) + 2e- Anode (oxidation):

Cathode (reduction):

H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)

E0 = Ecathode - Eanode cell 0 0

E0 = 0.34 V cell

Ecell = ECu /Cu – EH /H 2+ + 2 0 0 0

0.34 = ECu /Cu - 0 0

2+

ECu /Cu = 0.34 V 2+ 0

표준전극전위(Standard Electrode Potentials)

표준환원전극전위

Standard Reduction Potentials at 25oC* Half-Reaction E0 (V) Half-Reaction E0 (V)

F2(g) + 2 e- → F-(aq) 2.87

O3(g) + 2 H+(aq) + 2 e- → O2(g) + H2O 2.07

Co3+(aq) + e- → Co2+(aq) 1.82 2 H+(aq) + 2 e- → H2(g) 0.00

H2O2(aq) + 2 H+(aq) + 2 e- → 2 H2O 1.77 Pb

2+(aq) + 2 e- → Pb(s) -0.13

PbO2(s) + 4 H+(aq) + SO4

2-(aq) + 2 e- → PbSO4(s) + 2 H2O 1.70 Sn2+(aq) + 2 e- → Sn(s) -0.14

Ce4+(aq) + e- → Ce3+(aq) 1.61 Ni2+(aq) + 2 e- → Ni(s) -0.25

MnO4-(aq) + 8 H+(aq) + 5 e- → Mn2+(aq) + 4 H2O 1.51 Co

2+(aq) + 2 e- → Co(s) -0.28

Au3+(aq) + 3 e- → Au(s) 1.50 PbSO4(s) + 2 e- → Pb(s) + SO4

2-(aq) -0.31

Cl2(g) + 2 e- → 2 Cl-(aq) 1.36 Cd2+(aq) + 2 e- → Cd(s) -0.40

Cr2O72-(aq) + 14 H+(aq) + 6 e- → 2 Cr3+(aq) + 7 H2O 1.33 Fe

2+(aq) + 2 e- → Fe(s) -0.44

MnO2(s) + 4 H+(aq) + 2 e- → Mn2+(aq) + 2 H2O 1.23 Cr

3+(aq) + 3 e- → Cr(s) -0.74

O2(g) + 4 H+(aq) + 4 e- → 2 H2O 1.23 Zn

2+(aq) + 2 e- → Zn(s) -0.76

Br2(l) + 2 e- → 2 Br-(aq) 1.07 2 H2O + 2 e

- → H2(g) + 2 OH-(aq) -0.83

NO3-(aq) + 4 H+(aq) + 3 e- → NO(g) + 2 H2O 0.96 Mn

2+(aq) + 2 e- → Mn(s) -1.18

2 Hg2+(aq) + 2 e- → Hg22+(aq) 0.92 Al3+(aq) + 3 e- → Al(s) -1.66

Hg22+(aq) + 2 e- → 2 Hg(l) 0.85 Be2+(aq) + 2 e- → Be(s) -1.85

Ag+(aq) + e- → Ag(s) 0.80 Mg2+(aq) + 2 e- → Mg(s) -2.37

Fe3+(aq) + e- → Fe2+(aq) 0.77 Na+(aq) + e- → Na(s) -2.71

O2(g) + 2 H+(aq) + 2 e- → H2O2(aq) 0.68 Ca

2+(aq) + 2 e- → Ca(s) -2.87 MnO4

-(aq) + 2 H2O + 3 e- → MnO2(s) + 4 OH

-(aq) 0.59 Sr2+(aq) + 2 e- → Sr(s) -2.89 I2(s) + 2 e

- → 2 I-(aq) 0.53 Ba2+(aq) + 2 e- → Ba(s) -2.90 O2(g) + 2 H2 + 4 e

- → 4 OH-(aq) 0.40 K+(aq) + e- → K(s) -2.93 Cu2+(aq) + 2 e- → Cu(s) 0.34 Li+(aq) + e- → Li(s) -3.05 AgCl(s) + e- → Ag(s) + Cl-(aq) 0.22 SO4

2-(aq) + 4 H+(aq) + 2 e- → SO2(g) + 2 H2O 0.20

Cu2+(aq) + e- → Cu+(aq) 0.13

Sn4+(aq) + 2 e- → Sn2+(aq) 0.13

Easie

r to re

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. Stro

ng

er o

xid

izer E

asie

r to

oxid

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Str

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reducer

갈바니 전지에 대한 전위 계산

E0 = Ecathode - Eanode cell 0 0 1.

2. 열역학 방법, using DG = -nFE

Potential, Work and DG

• DGº = -nFEº

• if Eº > 0, then DGº < 0 spontaneous

• if Eº< 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

3Cd2+ (aq) + 6e- → 3Cd (s)

2Cr3+ (aq) + 6e- → 2Cr (s)

2Cr (s) + 3Cd2+ (1 M) → 3Cd (s) + 2Cr3+ (1 M)

another approach

DG10 = -nF Eh

0 = -6F(-0.40) V

DG20 = -nF Eh

0 = -6F(-0.74) V

DG0 = DG10 - DG2

0 = -6F(-0.40+0.74) V =-6F E0

E0 = 0.34 V cell

cell

Cu2+ (aq) + 2e- Cu (s) E0 = 0.340 V

Cu+ (aq) + e- Cu (s) E0 = 0.522 V

Anode (oxidation):

Cathode (reduction): 2e- + Cu2+ Cu (s)

Cu (s) Cu+ + e-

Cu2+ + e- Cu+

E0 = Ecathode - Eanode cell 0 0

E0 = 0.340– (0.522) = -0.184 V ? hcell

Cu2+ (aq) + e- → Cu+ (aq) E0 = ?

Cu2+ (aq) + 2e- Cu (s)

Cu+ (aq) + e- Cu (s)

Cu2+ (aq) + e- → Cu+ (aq) Eh0 = ?

DG10 = -nF Eh

0 = -2F(0.340) V

DG20 = -nF Eh

0 = -F(0.522) V

DG0 = DG10 - DG2

0 = -F(2(0.340)-0.522) V =-F Eh0

Eh0 = 0.158 V

Cu2+ → Cu+ (aq) → Cu 0.158 V 0.522 V

0.340 V

Fe2+ (aq) + 2e- Fe (s)

Fe3+ (aq) + e- Fe2+(aq)

Fe3+ (aq) + 3e- → Fe (s) Eh0 = ?

Eh0 = ?

Fe3+ → Fe2+ (aq) → Fe +0.771 V -0.440 V

?

Eh0 = -0.440 V

Eh0 = +0.771 V

산화환원 반응의 열역학

DG = -nFEcell

DG0 = -nFEcell 0

n = number of moles of electrons in reaction

F = 96,485 J

V • mol = 96,485 C/mol

DG0 = -RT ln K = -nFEcell 0

Ecell 0 =

RT

nF ln K

(8.314 J/K•mol)(298 K)

n (96,485 J/V•mol) ln K =

= 0.0257 V

n ln K Ecell

0

= 0.0592 V

n log K Ecell

0

산화환원 반응의 열역학

전지의 전위에 미치는 농도의 효과

DG = DG0 + RT ln Q DG = -nFE DG0 = -nFE 0

-nFE = -nFE0 + RT ln Q

E = E0 - ln Q RT

nF Nernst equation At 298

- 0.0257 V

n ln Q E 0 E = -

0.0592 V n

log Q E 0 E =

The Nernst Equation

DG = DGº +RTln(Q)

• -nFE = -nFEº + RTln(Q)

• E = Eº - RT ln(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 emf of the cell made up of the glass electrode and

the reference electrode is measured with a voltmeter that

is calibrated in pH units.

Measurement of pH: the glass electrode

Very thin glass membrane

that is permeable to H+ ions.

A potential difference develops

between the two sides of the

membrane.

Glass pH Electrodes

1. a sensing part of electrode, a bulb made from a specific glass

2. sometimes the electrode contains a small amount of AgCl precipitate inside the glass electrode

3. internal solution, usually 0.1M HCl for pH electrodes

4. internal electrode, usually silver chloride electrode or calomel electrode

5. body of electrode, made from non-conductive glass or plastics.

6. reference electrode, usually the same type as 4

7. junction with studied solution, usually made from ceramics or capillary with asbestos or quartz fiber.

Ag/AgCl reference electrode Glass pH electrode

+ -

전지

Leclanché cell

Dry cell

Zn (s) Zn2+ (aq) + 2e- Anode:

Cathode: 2NH4 (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)

+

Zn (s) + 2NH4+ (aq) + 2MnO2 (s) Zn

2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e- Anode:

Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)

Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

Mercury Battery

전지

Location Reaction Potential

Anode Zn + 2OH- > Zn(OH)2 + 2e- 1.25 V

Cathode HgO +H2O + 2e- > Hg + 2OH- 0.098 V

Overall Zn + HgO + H2O > Zn(OH)2 + Hg 1.35 V

Anode:

Cathode:

Lead storage

battery

PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e- PbSO4 (s) + 2H2O (l) 4

Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4

Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l) 4

전지

Lead-acid

battery

Solid State Lithium Battery

전지

연료전지(Fuel Cell ) converts chemical energy into electricity.

In contrast to storage battery, fuel cell does not need to involve a

reversible reaction since the reactant are supplied to the cell as

needed from an external source. This technology has been used in

the Gemini, Apollo and Space Shuttle program.

Half reactions: E°Cell = 0.9 V

Advantage: Clean, portable and product is water. Efficient (75%) contrast to 20-25% car, 35-40% from coal electrical plant Disadvantage: Needs continuous flow of reactant, Electrodes are short lived and expensive.

Anode:

Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)

2H2 (g) + 4OH- (aq) 4H2O (l) + 4e

-

2H2 (g) + O2 (g) 2H2O (l)

Relative Energy Density of Some Common

Secondary Cell Chemistries

Current Science: microbial fuel cell (MFC) technology

From: Andrew Kato Marcus, arizona State Univ.,

http://researchstories.asu.edu/2008/01/post_1.htmlhttp://researchstories.asu.edu/2008/01/post_1.htmlhttp://researchstories.asu.edu/2008/01/post_1.htmlhttp://researchstories.asu.edu/2008/01/post_1.html

Electrolysis is the process in which electrical energy is used

to cause a nonspontaneous chemical reaction to occur.

전기분해

Electrolysis of Water

2H2O(l) → 2 H2(g) + O2(g)

Electrolysis of Water

2H2O(l) → 2 H2(g) + O2(g) DG0 = 474.4 kJ/mol

Anode:

2 H2O(l) → O2 + 4 H+(aq) + 4 e- Eoox = -1.23 V

2 SO42- → S2O8

2- + 2 e- Eoox = -2.05 V

39 kWh of electricity and 8.9 liters of water are required

to produce 1 kg of hydrogen at 25°C and 1 atm. In reality, 50.3-70.1 kWh of electricity is required.

Cathode:

2 H+(aq) + 2 e- → H2(g) Eo

red = 0.00 V

Electrolysis of NaCl(aq)

2H2O(l) + 2 Cl-(aq) → H2(g) + Cl2(g) + 2 OH

-(aq)

Electrolysis of NaCl(aq)

Anode:

(1) 2 H2O(l) → O2 + 4 H+(aq) + 4 e- Eoox = -1.23 V

(2) 2 Cl- → Cl2 + 2 e- Eoox = -1.36 V

Cathode:

(1) 2 H+(aq) + 2 e- → H2(g) Eo

red = 0.00 V

(2) Na+ + e- → Na Eored = -2.71 V

(3) 2 H2O + 2 e- → H2 + 2 OH

- Eored = -0.83 V

(1) is favored at standard state, but [H+] is so low in the

solution, so (3) is the reaction taking place.

(1) is favored if ideal, but overvoltage for (1) is so high

in reality, so (2) is the reaction taking place.

Electrolysis and Mass Changes

charge (C) = current (A) x time (s)

1 mole e- = 96,485.34 C

Dry Cell or LeClanche Cell Dry Cells

Invented in the 1860’s the common dry cell or LeClanche cell, has become a familiar household

item. An active zinc anode in the form of a can house a mixture of MnO2 and an acidic electrolytic

paste, consisting of NH4Cl, ZnCl2, H2O and starch powdered graphite improves conductivity. The

inactive cathode is a graphite rod.

Anode (oxidation)

Zn(s) g Zn2+

(aq) + 2e-

Cathode (reduction). The cathodic half-reaction is complex and even today, is

still being studied. MnO2(s) is reduced to Mn2O3(s) through a series of steps that

may involve the presence of Mn2+ and an acid-base reaction between NH4+ and

OH- :

2MnO2 (s) + 2NH4+

(aq) + 2e- g Mn2O3(s) + 2NH3(aq) + H2O (l)

The ammonia, some of which may be gaseous, forms a complex ion with Zn2+,

which crystallize in contact Cl- ion:

Zn2+(aq) + 2NH3 (aq) + 2Cl

-(aq) g Zn(NH3)2Cl2(s)

Overall Cell reaction:

2MnO2 (s) + 2NH4Cl(aq) + Zn(s) g Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s) Ecell = 1.5 V

Uses: common household items, such as portable radios, toys, flashlights,

Advantage; Inexpensive, safe, available in many sizes

Disadvantages: At high current drain, NH3(g) builds up causing drop in voltage,

short shelf life because zinc anode reacts with the acidic NH4+ ions.

Dry Cell or LeClanche Cell

Invented by George Leclanche, a French Chemist.

Acid version: Zinc inner case that acts as the anode and a carbon rod in contact with a moist paste of solid MnO2 , solid NH4Cl, and carbon

that acts as the cathode. As battery wear down, Conc. of Zn+2 and NH3 (aq)

increases thereby decreasing the voltage.

Half reactions: E°Cell = 1.5 V

Anode: Zn(s) g Zn+2(aq) + 2e-

Cathode: 2NH4+(aq) + MnO2(s) + 2e

- g Mn2O3(s) + 2NH3(aq) + H2O(l)

Advantage: Inexpensive, safe, many sizes Disadvantage: High current drain, NH3(g) build up, short shelf life

Alkaline Battery Alkaline Battery

The alkaline battery is an improved dry cell. The half-reactions are similar, but the

electrolyte is a basic KOH paste, which eliminates the buildup of gases and maintains the

Zn electrode.

Anode (oxidation)

Zn(s) + 2OH- (aq) g ZnO(s) + H2O (l) + 2e-

Cathode (reduction).

2MnO2 (s) + 2H2O (l) + 2e- g Mn(OH)2(s) + 2OH

-(aq)

Overall Cell reaction:

2MnO2 (s) + H2O (l) + Zn(s) g ZnO(s) + Mn(OH)2(s) Ecell = 1.5 V

Uses: Same as for dry cell.

Advantages: No voltage drop and longer shell life than dry cell

because of alkaline electrolyte; sale ,amu sizes.

Disadvantages; More expensive than common dry cell.

Alkaline Battery

Leclanche Battery: Alkaline Version

In alkaline version; solid NH4Cl is replaced with KOH or NaOH. This

makes cell last longer mainly because the zinc anode corrodes less

rapidly under basic conditions versus acidic conditions.

E°Cell = 1.5 V

Anode: Zn(s) + 2OH-(aq) g ZnO(s) + H2O(l) + 2e

-

Cathode: MnO2 (s) + H2O(l) + 2e- g MnO3 (s) + 2OH

-(aq)

Nernst equation: E = E° - [(0.592/n)log Q], Q is constant !!

Advantage: No voltage drop, longer shelf life. Disadvantage: More expensive

Mercury Button Battery

Mercury and Silver batteries are similar.

Like the alkaline dry cell, both of these batteries use zinc in a basic

medium as the anode. The solid reactants are each compressed with

KOH, and moist paper acts as a salt bridge.

E°Cell = 1.6 V

Anode: Zn(s) + 2OH-(aq) g ZnO(s) + H2O(l) + 2e

-

Cathode (Hg): HgO (s) + 2H2O(l) + 2e- g Hg(s) + 2OH

-(aq)

Cathode (Ag): Ag2O (s) + H2O(l) + 2e- g 2Ag(s) + 2OH

-(aq)

Advantage: Small, large potential, silver is nontoxic. Disadvantage: Mercury is toxic, silver is expensive.

Lead Storage Battery

•Lead-Acid Battery. A typical 12-V lead-acid battery has six cells connected in series, each of which delivers about 2 V. Each cell contains

two lead grids packed with the electrode material: the anode is spongy Pb,

and the cathode is powered PbO2. The grids are immersed in an

electrolyte solution of 4.5 M H2SO4. Fiberglass sheets between the grids

prevents shorting by accidental physical contact. When the cell

discharges, it generates electrical energy as a voltaic cell.

Half reactions: E°Cell = 2.0 V

Anode: Pb(s) + SO42- g PbSO4 (s) +2 e- E° = 0.356

Cathode (Hg): PbO2 (s) + SO42- + 4H+ + 2e- g

PbSO4 (s) + 2 H2O E° = 1.685V

Net: PbO2 (s) + Pb(s) + 2H2SO4 g PbSO4 (s) + 2 H2O E°Cell = 2.0 V

Note hat both half-reaction produce Pb2+ ion, one through

oxidation of Pb, the other through reduction of PbO2. At both

electrodes, the Pb2+ react with SO42- to form PbSO4(s)

Nickel-Cadmium Battery

Battery for the Technological Age

Rechargeable, lightweight “ni-cad” are used for variety of cordless appliances.

Main advantage is that the oxidizing and reducing agent can be regenerated

easily when recharged. These produce constant potential.

Half reactions: E°Cell = 1.4 V

Anode: Cd(s) + 2OH-(aq) g Cd(OH)2 (s) + 2e

- Cathode: 2Ni(OH) (s) + 2H2O(l) + 2e

- g Ni(OH)2 (s) + 2 OH-(aq)

보통 anode를 "양극", cathode를 "음극"이라고 번역한다. 우리가 흔히 사용하는 CRT (cathode ray tube)를 음극관이라고 한다. 그러나 이러한 명칭이 혼돈을 주는 경우도 있다. 우리가 일상에서 접하는 배터리의 플러스 극은 cathode이고 마이너스 극은 anode이다.

다음을 명심하면 혼돈을 피할 수 있다.

1) 갈바니 전지와 전해전지에서 모두 산화가 일어나는 전극은 "anode"이고, 환원이 일어나는 전극은 "cathode"이다.

2) 갈바니 전지는 자발적인 반응이고, 전해전지는 비자발적인 변화를 강제로 일으키는 것이다.

3)전류는 전위가 높은 곳에서 낮은 곳으로 흐르며, 전자의 이동 방향은 전류의 방향과는 반대이다.

갈바니 전지에서는 산화가 일어나는 anode에서 전자가 나와서 환원이 일어나는 cathode로 흘러 들어가게 되므로, 전류는 cathode에서 anode로 흐른다. 그래서 cathode의 전위가 더 높다. cathode는 플러스극이고 anode는 마이너스 극이 된다.

전해 전지에서도 anode에서 전자가 나오고, cathode로 전자가 들어간다. 그러나 이 때의 전자는 자발적으로 흐르는 것이 아니라 외부에 설치한 전류 공급원에서 강제로 흘려주는 것이다. 그러므로 외부의 전류 공급원의 전위가 낮은 쪽에서 흘러나온 전자가 cathode로 들어가고, 전위가 더 높은 쪽에서는 anode에서 전자를 강제로 빼앗아 온다. 따라서 이 경우에는 cathode의 전위가 anode의 전위보다 더 낮다.

대한화학회에서는 anode를 "산화극", cathode를 "환원극"이라고도 번역한다.