chapter 19 2018. 4. 3. · 2 (g) 2mgo (s) 2mg 2mg2+ + 4e-o 2-+ 4e 2o2-oxidation half-reaction (lose...
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Electrochemistry
Chapter 19
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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
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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
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Reducing
Agent
Oxidizing
Agent
e-
e-
e- e-
e-
e-
Anode Cathode
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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
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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
du
ce
. Stro
ng
er o
xid
izer E
asie
r to
oxid
ize.
Str
onger
reducer
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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
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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
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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
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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.
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