Electrochemistry

ElectrochemistryChapter 19

산화환원 반응과 전기화학

1. 산화환원 반응과 산화수2. 갈바니 전지3. 표준환원전위4. 산화환원 반응의 자발성 . 전위차와 자유에너지5. 전위차와 농도 , Nernst 방정식6. 전지7. 부식8. 전기분해9. 전기야금

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

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Oxidation half-reaction (lose e-)

Reduction half-reaction (gain e-)

전기화학 반응oxidation-reduction reactions in which:

• 자유에너지의 감소 => electrical energy : 갈바니 전지

• electrical energy => 자유에너지의 증가 : 전해 전지

0 0 2+ 2-

산화수 (Oxidation number)

The charge the atom would have in a molecule (or anionic compound) if electrons were completely transferred.

1. Free elements (uncombined state) have an oxidation number of zero.

Na, Be, K, Pb, H2, O2, P4 = 0

2. In monatomic ions, the oxidation number is equal to the charge on the ion.

Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2

3. The oxidation number of oxygen is usually –2. In H2O2

and O22- it is –1.

4. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.

6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.

5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.

HCO3-

O = -2 H = +1

3x(-2) + 1 + ? = -1

C = +4

Oxidation numbers of all the atoms in HCO3

- ?

산화 - 환원 반응식의 균형 맞추기

1. 균형되지 않은 반응식을 이온형으로 쓴다 .

산성 용액에서 Fe2+ 가 Cr2O72- 에 의해 Fe3+ 로 산화되는

반응

Fe2+ + Cr2O72- Fe3+ + Cr3+

2. 반응식을 두 개의 반쪽반응으로 나눈다 .

Oxidation:

Cr2O72- Cr3+

+6 +3

Reduction:

Fe2+ Fe3++2 +3

3. 각 반쪽반응에 대하여 O 와 H 를 제한 나머지의 계수를 맞춘다 .

Cr2O72- 2Cr3+

4. 산성 용액의 반응에서 , O 원자의 균형을 맞추기 위해 H2O 을 첨가하고 , H 원자의 균형을 맞추기 위해 H+ 를 첨가한다 .

Cr2O72- 2Cr3+ + 7H2O

14H+ + Cr2O72- 2Cr3+ + 7H2O

5. 전자를 첨가하여 각 반쪽 반응의 전하 균형을 맞춘다 .

Fe2+ Fe3+ + 1e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

6. 필요하면 각각의 반쪽반응에 적절한 계수를 곱하여 두 반쪽 반응에서의 전자의 수를 같도록 한다 .

6Fe2+ 6Fe3+ + 6e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O


7. 두 반쪽반응을 더하여 최종적으로 균형을 맞춘다 . 양변의 전자수가 상쇄되어야 한다 .

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

6Fe2+ 6Fe3+ + 6e-Oxidation:

Reduction:

14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O

8. 균형을 맞춘 반응식의 각 변의 원자와 전자의 수가 같은지 확인한다 .

14x1 – 2 + 6x2 = 24 = 6x3 + 2x3

9. 염기성 용액에 대하여는 , 최종 식에 나오는 H+ 마다 OH- 를 양변에 모두 추가한다 ..


산화 - 환원 반응식의 균형 맞추기 - 예제

Cu(s) + HNO3(aq) Cu2+(aq) + NO(g)

산화 : Cu Cu2+

환원 : NO3– NO

2. 반응식을 두 개의 반쪽반응으로 나눈다 .

3. 각 반쪽반응에 대하여 O 와 H 를 제한 나머지의 계수를 맞춘다 .

4. 산성 용액의 반응에서 , O 원자의 균형을 맞추기 위해 H2O 을 첨가하고 , H 원자의 균형을 맞추기 위해 H+ 를 첨가한다 .

환원 : 4H+ + NO3– NO + 2H2O

5. 전자를 첨가하여 각 반쪽 반응의 전하 균형을 맞춘다 .

산화 : Cu Cu2+ + 2e–

환원 : 4H+ + NO3– + 3e– NO + 2H2O

구리를 질산에 넣으면 산화질소 (NO) 가 발생한다 .

산화 - 환원 반응식의 균형 맞추기 - 예제

6. 필요하면 각각의 반쪽반응에 적절한 계수를 곱하여 두 반쪽 반응에서의 전자의 수를 같도록 한다 .

산화 : 3Cu 3Cu2+ + 6e–

환원 : 8H+ + 2NO3– + 6e– 2NO + 4H2O

7. 두 반쪽반응을 더하여 최종적으로 균형을 맞춘다 . 양변의 전자수가 상쇄되어야 한다 .

3Cu + 8H+ + 2NO3– 3Cu2+ + 2NO + 4H2O

or 3Cu + 6H+ + 2HNO3 3Cu2+ + 2NO + 4H2O

or 3Cu + 8HNO3 3Cu(NO3)2 + 2NO + 4H2O

산화 - 환원 반응식의 균형 맞추기 - 예제1) H2O2 + Fe2+ → Fe3+ + H2O (acidic)

2) Cu + HNO3 → Cu2+ + NO + H2O (acidic)

3) CN- + MnO4- → CNO- + MnO2 (basic, 힌트 C 의 산화수는 양쪽에서 모두

+2)

4) Br2 → BrO3- + Br- (basic, 힌트 Br2 이 산화와 환원이 함께 일어난다 )

5) Cr(OH)3(s) + ClO3−(aq) → CrO4

2−(aq) + Cl−(aq) (basic)

갈바니 전지 (Galvanic Cells)

spontaneousredox reaction

anodeoxidation

cathodereduction

+-

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


E0 = 0.76 Vcell

Standard emf (E0 )cell

0.76 V = 0 - EZn /Zn 0

2+

EZn /Zn = -0.76 V02+

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

E0 = EH /H - EZn /Zn cell0 0

+ 2+2

E0 = Ecathode - Eanodecell0 0

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


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 = 0.34 Vcell

Ecell = ECu /Cu – EH /H 2+ +2

0 0 0

0.34 = ECu /Cu - 00 2+

ECu /Cu = 0.34 V2+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) 0.00H2O2(aq) + 2 H+(aq) + 2 e- → 2 H2O 1.77 Pb2+(aq) + 2 e- → Pb(s) -0.13PbO2(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) -0.25MnO4

-(aq) + 8 H+(aq) + 5 e- → Mn2+(aq) + 4 H2O 1.51 Co2+(aq) + 2 e- → Co(s) -0.28

Au3+(aq) + 3 e- → Au(s) 1.50PbSO4(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) -0.40Cr2O7

2-(aq) + 14 H+(aq) + 6 e- → 2 Cr3+(aq) + 7 H2O 1.33 Fe2+(aq) + 2 e- → Fe(s) -0.44MnO2(s) + 4 H+(aq) + 2 e- → Mn2+(aq) + 2 H2O 1.23 Cr3+(aq) + 3 e- → Cr(s) -0.74O2(g) + 4 H+(aq) + 4 e- → 2 H2O 1.23 Zn2+(aq) + 2 e- → Zn(s) -0.76Br2(l) + 2 e- → 2 Br-(aq) 1.07 2 H2O + 2 e- → H2(g) + 2 OH-(aq) -0.83NO3

-(aq) + 4 H+(aq) + 3 e- → NO(g) + 2 H2O 0.96 Mn2+(aq) + 2 e- → Mn(s) -1.182 Hg2+(aq) + 2 e- → Hg2

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

2+(aq) + 2 e- → 2 Hg(l) 0.85 Be2+(aq) + 2 e- → Be(s) -1.85Ag+(aq) + e- → Ag(s) 0.80 Mg2+(aq) + 2 e- → Mg(s) -2.37Fe3+(aq) + e- → Fe2+(aq) 0.77 Na+(aq) + e- → Na(s) -2.71O2(g) + 2 H+(aq) + 2 e- → H2O2(aq) 0.68 Ca2+(aq) + 2 e- → Ca(s) -2.87MnO4

-(aq) + 2 H2O + 3 e- → MnO2(s) + 4 OH-(aq) 0.59 Sr2+(aq) + 2 e- → Sr(s) -2.89I2(s) + 2 e- → 2 I-(aq) 0.53 Ba2+(aq) + 2 e- → Ba(s) -2.90O2(g) + 2 H2 + 4 e- → 4 OH-(aq) 0.40 K+(aq) + e- → K(s) -2.93Cu2+(aq) + 2 e- → Cu(s) 0.34 Li+(aq) + e- → Li(s) -3.05AgCl(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 　　

Easier to reduce. S

tronger oxidizer Eas

ier

to o

xidi

ze.

Str

onge

r re

duce

r

갈바니 전지에 대한 전위 계산

E0 = Ecathode - Eanodecell0 01.

2. 열역학 방법 , using G = -nFE

What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr electrode in a 1.0 M Cr(NO3)3 solution?

Cd2+ (aq) + 2e- Cd (s) E0 = -0.40 V

Cr3+ (aq) + 3e- Cr (s) E0 = -0.74 V

Cd is the stronger oxidizer

Cd will oxidize Cr

Anode (oxidation):

Cathode (reduction): 2e- + Cd2+ (1 M) Cd (s)

Cr (s) Cr3+ (1 M) + 3e-

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

x 2

x 3


E0 = -0.40 – (-0.74) cell

E0 = 0.34 V cell

3Cd2+ (aq) + 6e- → 3Cd (s)

2Cr3+ (aq) + 6e- → 2Cr (s)

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

G10 = -nF Eh

0 = -6F(-0.40) V

G20 = -nF Eh

0 = -6F(-0.74) V

G0 = G10 - G2

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 = 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 = ?

G10 = -nF Eh

0 = -2F(0.340) V

G20 = -nF Eh

0 = -F(0.522) V

G0 = G10 - G2

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

Eh0 = 0.158 V

Cu2+ → Cu+ (aq) → Cu0.158 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

G = -nFEcell

G0 = -nFEcell0

n = number of moles of electrons in reaction

F = 96,485J

V • mol = 96,485 C/mol

G0 = -RT ln K = -nFEcell0

Ecell0 =

RTnF

ln K(8.314 J/K•mol)(298 K)

n (96,485 J/V•mol)ln K=

=0.0257 V

nln KEcell

0

=0.0592 V

nlog KEcell

0

산화환원 반응의 열역학

Oxidation:

Reduction:

What is the equilibrium constant for the following reaction at 250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq)

=0.0257 V

nln KEcell

0

2e- + Fe2+ Fe

2Ag 2Ag+ + 2e-

n = 2

0.0257 Vx nE0 cellexpK =

0.0257 Vx 2-1.24 V

= exp

K = 1.23 x 10-42

E0 = -0.44 – (0.80)

E0 = -1.24 V

E0 = EFe /Fe – EAg /Ag0 0

2+ +

EFe /Fe = -0.44 VEAg /Ag = 0.80 V

0

02+

+

전지의 전위에 미치는 농도의 효과

G = G0 + RT ln Q G = -nFE G0 = -nFE 0

-nFE = -nFE0 + RT ln Q

E = E0 - ln QRTnF

Nernst equation

At 298

-0.0257 V

nln QE0E = -

0.0592 Vn

log QE0E =

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 membranethat is permeable to H+ ions.

A potential difference developsbetween 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

+ -

Will the following reaction occur spontaneously at 250C if [Fe2+] = 0.60 M and [Cd2+] = 0.010 M? Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)

Oxidation:

Reduction: 2e- + Fe2+ 2Fe

Cd Cd2+ + 2e-

n = 2

E0 = -0.44 – (-0.40)

E0 = -0.04 V

E0 = EFe /Fe – ECd /Cd0 0

2+ 2+ -0.0257 V

nln QE0E =

-0.0257 V

2ln -0.04 VE =

0.0100.60

E = 0.013

E > 0 Spontaneous

EFe /Fe = -0.44 VECd /Cd = -0.40 V

0

02+

2+

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) Zn2+ (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

OverallZn + HgO + H2O > Zn(OH)2 + Hg

1.35 V

Anode:

Cathode:

Lead storagebattery

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-acidbattery

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.html




부동화 (Passivation)

galvanized steel

Fe

Zn

Fe2+ + 2e- Fe Eo = -0.44 V

Sn2+ + 2e- Sn Eo = -0.14 V

Zn2+ + 2e- Zn Eo = -0.76 V

tin steel

Fe

Sn

Al

Al2O3

1994 년 10 월 21 일 아침 7 시 38 분

성수대교 붕괴

부식 (Corrosion)

Fe2O3·xH2O and (FeO(OH), Fe(OH)3).

O2 + 4 e- + 2 H2O → 4 OH- E° = 1.23 V

Because it forms OH-, it is strongly affected by acid. Fe → Fe2+ + 2 e− E° = 0.44 VFe2+ → Fe3+ + e− E° = -0.771 V

Additional reactions involved:Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2 H+

Fe3+ + 3 H2O ⇌ Fe(OH)3 + 3 H+

Fe(OH)2 ⇌ FeO + H2O

Fe(OH)3 ⇌ FeO(OH) + H2O

2 FeO(OH) ⇌ Fe2O3 + H2

Conditions for CorrosionConditions for Iron Oxidation:

Iron will oxidize in acidic medium

SO2 H2SO4 H+ + HSO4+

Anions improve conductivity for oxidation.

Cl- from seawater or NaCl (snow melting) enhances rusting

Conditions for Prevention:

Iron will not rust in dry air; moisture must be present

Iron will not rust in air-free water; oxygen must be present

Iron rusts most rapidly in ionic solution and low pH (high H+)

The loss of iron and deposit of rust occur at different placm on object

Iron rust faster in contact with a less active metal (Cu)

Iron rust slower in contact with a more active metal (Zn)

Corrosion Prevention

Impressed current cathodic protection system in seawater

Sacrificial anode system in seawater

from, Richard Baxter, Jim Britton, www.stoprust.com

Cathodic Protection of an Iron Storage Tank

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) G0 = 474.4 kJ/mol

Anode:2 H2O(l) → O2 + 4 H+(aq) + 4 e- Eo

ox = -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- Eo

ox = -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

How much Ca will be produced in an electrolytic cell of molten CaCl2 if a current of 0.452 A is passed through the cell for 1.5 hours?

Anode:

Cathode: Ca2+ (l) + 2e- Ca (s)

2Cl- (l) Cl2 (g) + 2e-

Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g)

2 mole e- = 1 mole Ca

mol Ca = 0.452Cs

x 1.5 hr x 3600shr 96,500 C

1 mol e-

x2 mol e-

1 mol Cax

= 0.0126 mol Ca

= 0.50 g Ca

Chemistry In Action: Dental Filling Discomfort

Hg2 /Ag2Hg3 0.85 V2+

Sn /Ag3Sn -0.05 V2+

Dental amalgam consists of three solid phases having approximately corresponding to Ag2Hg3, Ag3Sn, and Sn8Hg.

Ag3Sn + Ag2Hg3 + Sn7-8Hg

Gamma Gamma-1 Gamma-2

Sn /Sn8Hg -0.05 V2+

Dry Cell or LeClanche CellDry CellsInvented 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) 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- 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) Zn(NH3)2Cl2(s)

Overall Cell reaction:Overall Cell reaction:2MnO2 (s) + 2NH4Cl(aq) + Zn(s) Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s) Ecell = 1.5 V

Uses: common household items, such as portable radios, toys, flashlights,Advantage;Advantage; Inexpensive, safe, available in many sizesDisadvantages: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 CellInvented 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.


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

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

Advantage:Inexpensive, safe, many sizes

Disadvantage:High current drain, NH3(g) build up, short shelf life

Alkaline BatteryAlkaline BatteryThe 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) ZnO(s) + H2O (l) + 2e-Cathode (reduction). 2MnO2 (s) + 2H2O (l) + 2e- Mn(OH)2(s) + 2OH-

(aq)

Overall Cell reaction::2MnO2 (s) + H2O (l) + Zn(s) ZnO(s) + Mn(OH)2(s) Ecell = 1.5 V

Uses:Uses: Same as for dry cell.Advantages:Advantages: No voltage drop and longer shell life than dry cell because of alkaline electrolyte; sale ,amu sizes.Disadvantages;Disadvantages; More expensive than common dry cell.

Alkaline BatteryLeclanche 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) ZnO(s) + H2O(l) + 2e-

Cathode: MnO2 (s) + H2O(l) + 2e- 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 BatteryMercury 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) ZnO(s) + H2O(l) + 2e-

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

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


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

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

PbSO4 (s) + 2 H2O E° = 1.685V

Net: PbO2 (s) + Pb(s) + 2H2SO4 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 SO4

2- to form PbSO4(s)

Nickel-Cadmium BatteryBattery for the Technological AgeRechargeable, 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.


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

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

보통 anode 를 " 양극 ", cathode 를 " 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 cathode 이고 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 anode 에서 전자가 나와서 환원이 일어나는 cathode로 흘러 들어가게 되므로 , 전류는 cathode 에서 anode 로 흐른다 . 그래서 cathode 의 전위가 더 높다 . cathode 는 플러스극이고 anode 는 마이너스 극이 된다 . 전해 전지에서도 anode 에서 전자가 나오고 , cathode 로 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 cathode 로 들어가고 , 전위가 더 높은 쪽에서는 anode 에서 전자를 강제로 빼앗아 온다 . 따라서 이 경우에는 cathode 의 전위가 anode 의 전위보다 더 낮다 . 대한화학회에서는 anode 를 " 산화극 ", cathode 를 " 환원극 " 이라고도 번역한다 .

Electrochemistry

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