Transcript Electronics

Basic Electrochemistry
What is electrochemistry?
• Electronics: the transport of electrons (or positive holes)
Optoelectronics: light + electronics
• Electrochemical systems (electrodics + ionics)
• Electrochemistry:
the coupling of chemical changes to the passage of electricity
 ionic conduction (flow of ions) + electronic conduction (flow of
electrons)
 Electrochemical devices & electrochemical technologies
 Materials & devices & processings
• Examples of Electrochemical devices/technologies
Battery or Fuel cell: chemical state changes(electrochemistry)  electric power
Photoelectrochemical cell (Solar cell): light + electrochemistry  electric power
Photocatalysis: light  hydrogen or chemical reaction
Electrochromic display: chemical state changes by electric signal  coloration
Sensors: chemical state changes by mass  electric signal
Electrolysis: electric power  chemical species by chemical state changes
Electrodeposition: electric power  chemical change: thin film, Cu metallization
Several distinct states may correspond to the same energy. That is, each energy level may be degenerate. Three energy levels are shown here, possessing one, three, and five distinct states.
Corrosion: potential difference  chemical change
Etching
• Solid State Electrochemistry
Solid electrolyte: solid substances which can conduct electric current by ionic
motion as do electrolyte solutions  “solid state electrochemistry” or “solid state
ionics”  “solid state device”
Basic concepts for electrochemistry
•
Electric charge & current
Electric charge (=amount of electricity) Q (unit: Coulomb, C), time t
Electric current (unit: ampere (A)):
I = dQ/dt
Q =  Idt
Current density (unit: A/m2): i = I/A,
A: surface of area
Ammeter: measuring current
Circuit: electric current flows in a closed path
Electrical potential & electric field
Electrical potential (unit; volts, V), : the pressure of the electric fluid
Voltage: the electrical potential difference ()
Voltmeter: measuring an electrical potential difference
Electric field strength (unit: V/m)
X = -d/dx
Ohm’s law: most conductors obey this law
Current density is proportional to the field strength
iX
i = X = - d/dx
; electrical conductivity (siemens/m, S = A/V), 1/; resistivity
 = -RI
R;resistance (unit of ohm), G; conductance,
G = 1/R = A/L = -I/
L; conductor length, A; cross section
Ohm’s law does not have universal validity. It does not apply to
electrochemical cells.
Resistor: a device that is fabricated to have a stable and known resistance
Power (watts) = I2R
Electrical quantities & their SI units
Quantity
Unit
Current (I)
Current density (i)
Charge (Q)
Charge density ()
Potential ()
Field strength (X)
Conductivity ()
Resistance (R)
Conductance (G)
Permittivity ()
Energy of work (w)
Power
Capacitance (C)
Ampere (A)
Ampere per square meter (A/m2)
Coulomb (C = As)
Coulomb per cubic meter (C/m3)
Volt (V = J/C)
Volt per meter (V/m)
Siemens per meter (S/m)
Ohm ( =1/S = V/A)
Siemens (S = A/V)
Farad per meter (F/m = C/Vm)
Joule (J = VC)
Watt (W = J/s = AV)
Farad (F = s/ = Ss), F = C/V
Classes of conductors
Materials 1.Conductors Electronic conductors
Ionic conductors
2. Insulators
Conductors: metals
Insulators: plastics, ceramics, gases
No clear cut distinction between conductor and insulator
Typical value of electrical conductivity
Material
/Sm-1
Ionic conductors
Ionic crystals
Solid electrolytes
Strong(liquid) electrolytes
10-16 – 10-2
10-1 – 103
10-1 – 103
Electronic conductors
Metals
Semiconductors
Insulators
103 – 107
10-3 – 104
<10-10
S/m  x10-2 for S/cm
Electrical conductivity of various materials (most at 298 K)
Material
/Sm-1
Superconductors (low temp)
Ag
Cu
Hg
C (graphite)
Doped polypyrrole
Molten KCl (at 1043 K)
5.2 M H2SO4 (battery acid)
Seawater
Ge
0.1 M KCl
H2O
Typical glass
Teflon, (CF2)n
Vacuum & most gases

6.3 x 107
6.0 x 107
1.0 x 106
4 x 104
6 x 103
217
82
5.2
2.2
1.3
5.7 x 10-6
3 x 10-10
10-15
0
Charge carriers
Electron pairs
Electrons
Electrons
Electrons
Pi electrons
Pi electrons
K+ and ClH+ and HSO4Cations & anions
Electrons and holes
K+ and ClH+ and OHUnivalent cations
?
Electronically conductive polymers
Mobilities: conduction from the standpoint of the charge carriers
Electric current = rate at which charge crosses any plane = [number of carriers
per unit volume][cross sectional area][charge on each carrier][average carrier
speed]
I = dQ/dt = (NAci)(A)(Qi)(i)
i: particular charge carrier, ci; concentration, Qi; charge, i; average velocity,
NA; Avogadro’s constant (6.0220 x 1023 mol-1), A; area
zi; charge number = Qi/Qe where Qe (1.6022 x 10-19 C),
e.g., electrons:-1, Mg2+; +2
i  fi  X  d/dx
fi; force exerted on the charge carrier, X; electric field strength
mobility of the carrier, ui (m2s-1V-1 unit) = velocity to field ratio (i / X)
i = uiX = - (zi/zi)uid/dx
zi: absolute value of the charge number
ue- of electrons: 6.7 x 10-3 m2s-1V-1 for Ag, less mobile in other metals
mobility of ions in aqueous solution: smaller than the factor of 105 (factor 105
slower); ucu2+o = 5.9 x 10-8 m2s-1V-1 in extremely diluted solution
Current I,
I = -A NAQeziuicid/dx
Faraday constant
F = NAQe = (6.02 x 1023 mol-1)(1.6022 x 10-19 C) = 96485 Cmol-1
is numerically equal to the charge carried by one mole of univalent cations.
(F is large. Small amount of chemicals
higher electricity)
If there are several kind of charge carriers,
I = -AFd/dxziuici
i = -Fd/dxziuici
Transport number ti; the fraction of the total current carried by one particular
charge carrier
ti = (ziuici )/(ziuici)
From i = X = -d/dx, conductivity 
 = Fziuici
molar ionic conductivity (i); Fui
Ion mobilities at extreme dilution in aqueous solution at 298 K
Ion
uo/m2s-1V-1
H+
K+
Ag+
Cu2+
Na+
Li+
OHSO42ClClO4C6H5COO-
362.5 x 10-9
76.2 x 10-9
64.2 x 10-9
58.6 x 10-9
51.9 x 10-9
40.1 x 10-9
204.8 x 10-9
82.7 x 10-9
79.1 x 10-9
69.8 x 10-9
33.5 x 10-9
Capacitance
parallel conducting plate separated by a narrow gap containing air or insulator
Idt = Q  E
Q = -CE
C; capacitance (unit; farads (F) = C/V)
C = -Q/E = A/L
A;cross-section area of the gap, L; width, ; permittivity of the insulator
• Relative permittivity (r) or dielectric constant (유전상수)
air: ~ 1
water: 78  Coulomb interaction energy is reduced by two orders of magnitudes
from its vacuum value
polar molecules: r
refractive index: nr = r1/2 at the frequency
Capacitor;  ; current integrator
Permittivity of various materials
Material
1012 /Fm-1
Material
1012 /Fm-1
vacuum (0)
N2(g)
Teflon(s), (CF2)n
CCl4(l)
Polyethene (s)
Mylar (s)
SiO2(s)
Typical glass (s)
C6H5Cl(l)
8.85419
8.85905
18
19.7
20
28
38.1
44
49.8
Neoprene
ClC2H4Cl(l)
CH3OH(l)
C6H5NO2(l)
CH3CN(l)
H2O(l)
HCONH2(l)
TiO2(s)
BaTiO3(s)
58
91.7
288.9
308.3
332
695.4
933
1500
110000
/0; relative permittivity or dielectric constant
mylar; poly(ethylene glycol terephthalate), (CH2OOCC6H4COOCH2)n
Liquid > solid: large capacitance in electrochemical capacitor (supercapacitor)
Summary
Electricity flows either by electron motion or ion motion
In both cases,
the intensity of the flow (= current density)  electric field strength
conductivity 
i = X = -d/dx
 = Fziuici
determined by the concentration of charge carriers and their mobilities
one form of Ohm’s law
E = -RI
potential difference across resistor to the current flowing through it
Resistor: dissipate energy
Capacitor: store energy
. Potential & Thermodynamics
Introduction
Electrochemistry: chemical change  electric force
Electrodics: in which the reactions at electrodes are considered
Ionics: in which the properties of electrolytes have the central attention 
concentration of ions, their mobilities, interactions etc
Basic laws were developed in systems with liquid electrolytes  “solid state”
(same and different features of solid electrolyte system)
Ionic solutions
Most important ionic conductor e.g., aqueous solution of electrolyte
Electrolyte; a substance that produces ions so enhance the electrical
conductivity
e.g., solid(NaCl), liquid(H2SO4), gas(NH3)
cf) solid electrolyte
Electrode
The junction between electronic conductor and ionic conductor that the
chemistry of electrochemistry occurs
Electrochemical cell
Basic unit: an ionic conductor sandwiched between two electronic conductors
e.g., aqueous solution of electrolyte between two pieces of metal, solid
electrolyte between two metals
Cell voltage (E) or emf(electromotive force)
electric potential difference between the two electronic conductors
voltameter
e.g., lead/acid cell (car battery)
Electronic conductors: PbO2, Pb
Ionic conductor: concentrated aqueous solution of sulfuric acid
Electrochemical reaction
Anode: Pb(s) + HSO4-(aq)  2e- + PbSO4(s) + H+(aq)
Cathode: PbO2(s) + HSO4-(aq) + 3H+(aq) + 2e-  PbSO4(s) + 2H2O(l)
Cell: PbO2(s) + Pb(s) + 2H+(aq) + 2HSO4-(aq)  2PbSO4(s) + 2H2O(l)
Right-hand electrode: electrons produced: oxidation, “anode”
Left-hand electrode: electrons consumed; reduction, “cathode”
Energy is delivered by the cell into the load; ex) car: starting engine, lighting
lamps
Galvanic cell: a cell which provides energy in this way, “discharge”(방전)
2.0 V without current flow, 1.8 V with current flow (load); “polarization”;
voltages decrease in magnitude when energy is taken from them. the effect
becomes greater if the current is increased.
“charge” (충전): current flow in the opposite direction by using an external
source (ex. Battery); Electrolytic cell; opposite direction to its spontaneous
motion
PbO2 : anode, Pb: cathode
2.0 V; perfect balance between the applied and cell voltages, no current flow 
equilibrium cell voltage or reversible cell voltage or null voltage or rest voltage
or “open-circuit voltage”(since no current flows, it makes no difference if the
circuit is interrupted, as by opening the switch)
Voltammogram
Plot of cell currents versus the cell voltages (volt + am(pere) + mogram)
Not linear  electrochemical cells do not obey Ohm’s law
Notation of the structure of cells
Zn/Zn2+, Cl-/AgCl/Ag
Hg/Hg2Cl2/Cl-(aq)//Zn2+(aq)/Zn
/: phase boundary, “,” or : two components in the same phase,
//: liquid junction (a salt bridge)
left: oxidation (anode), right: reduction(cathode)
Thermodynamics
Why is it that chemical reactions in electrochemical cells proceed
spontaneously in one direction and furnish current?
(thermodynamics: 평형상태에 대한 정보, kinetics: 전극반응속도에
대한 정보)
:
Cell potential of an electrochemical cell
Ecell = Eright – Eleft
or Ecell = Ecathode – Eanode
E obtained from the Nernst equation
oO + …+ ne- = rR + ….
pP + …. = qQ + … + neoO + pP + … = qQ + rR + …
(reduction)
(oxidation)
Ecell (cell reaction)
Ecell = E0 – (RT/nF)ln[(aQqaRr..)/(aOoaPp..)]
Gibbs free energy, G = -nFEcell
G <0  spontaneous
E0: standard electrode potential = Eright0 – Eleft0
Eright0, Eleft0,,: standard electrode potential of half reactions expresses as reductions
vs. NHE(normal hydrogen electrode) with all species at unit activity (ai =1)
(see the Table of Standard Potentials)
e.g., MnO2 + 4H+ + 2e-  Mn2+ + 2H2O
E0 = + 1.23 V
E = E0 –(RT/2F)ln[(aH+4)/aMn2+],
G = -nFE
aMnO2, aH2O = unity
cf. RT/2F = [(8.314 JK-1mol-1)(298 K)/2(96485 JV-1mol-1)] = 0.01285 V
Several distinct states may correspond to the same energy. That is, each energy level may be degenerate. Three energy levels are shown here, possessing one, three, and five distinct states.
Several distinct states may correspond to the same energy. That is, each energy level may be degenerate. Three energy levels are shown here, possessing one, three, and five distinct states.
e.g., Zn/Zn2+(aq), Cu2+(aq)/Cu
Cell: Zn + Cu2+  Zn2+ + Cu
Right: Cu2+ + 2e-  Cu E0 = +0.34 V
Left: Zn2+ + 2e-  Zn E0 = -0.76 V
Ecell0 = +0.34 – (-0.76) = +1.10 V
G0 = -2 x 1.10(V) x 96485 (JV-1mol-1) = -212 kJmol-1
reaction  spontaneous
EEcell = E0 – (RT/2F)ln(aZn2+/(aCu2+)
If we assume aZn2+= aCu2+, Ecell = 1.10 V
-----Hg/Hg2Cl2/Cl-(aq)//Zn2+(aq)/Zn
2Hg + Cl- + Zn2+  Hg2Cl2 + Zn
right: Zn2+ + 2e-  Zn E0 = -0.76 V
left: Hg2Cl2 + 2e-  2Hg + 2Cl- E0 = +0.27 V
Ecell0 = -0.76 –0.27 =-1.03 V, G0 = +199 kJmol-1, should be opposite direction
Measurement of E0:
(i) experiment
(ii) E0 = (RT/nF)lnK, K; equilibrium constant of cell  K = exp(-G0/RT)
(iii) E0 = Eright0 – Eleft0 or E0 = Ecathode0 – Eanode0 (from Table)
(iv) E0 = -G0/nF
Cell: PbO2(s) + Pb(s) + 2H+(aq) + 2HSO4-(aq)  2PbSO4(s) + 2H2O(l)
From thermodynamics Table,
Standard Gibbs Energy (kJmol-1): -813.76 (PbSO4(s)), -237.13 (H2O(l)), -218.96
(PbO2(s)), -755.91 (HSO4-(aq)), cf) G0 for element (Pb(s)) and H+(aq) = 0
G0 = 2G0 (PbSO4(s)) + 2G0 (H2O(l)) – [G0 (PbO2(s)) + 2G0 (HSO4-(aq))]
= -371 kJmol-1
 G0 = -nFE0
 E0 = 371000(Jmol-1)/[2 x 96485 (JV-1mol-1)] = 1.923 V
battery acid: 5.2 M
Ecell = 1.923 V – (RT/2F)ln[aH2O(l)2/(aH+(aq)2aHSO4-(aq)2)]
= 1.923 V – 0.01285ln [1/(5.2)2] = 2.008 V
activity term: minor contribution to the cell voltage
activity (a)  concentration (c); a = c, ; activity coefficient
ai  1(solvent, pure solid, ideal solution)
(Examples)
1. Indicate in the following reactions which are reductions and which are oxidations:
(1) Fe2+ + 2e-  Fe
(2) Cl-  1/2Cl2 + e- (3) Fe2+  Fe3+ + e(4) CrO42- + 3e-  Cr3+ (5) O2 + 4e-  2O2(6) Br2 + 2e-  2Br2. A Galvanic cell is constructed from a Cu2+/Cu electrode and an Ag+/Ag electrode.
(1) Make a schematic drawing of the cell
(2) Write the reactions at the electrode
(3) Indicate the anode and the cathode
3. Assuming standard states for all reactants and products, determine the spontaneous
direction of the following reactions by calculating the cell potential:
(1) Cu + 2HCl = CuCl2 + H2
(2) Ag + FeCl3 = FeCl2 + AgCl
Principles of electrochemistry: Definitions
Two equal electrodes (Lecture #4)  interest in one electrode only (Lecture #5)
Electrodes
Working electrode(WE): electrode of interest
Reference electrode(RE): second electrode, measure potential of WE with respect
to RE
Electrode potential E = Ework –Eref
Reference electrodes
SHE (standard hydrogen electrode) or NHE(normal hydrogen electrode):
universally accepted standard
H+(aq, a=1) + e- = 1/2H2(g, 105 Pa) E = 0 V
SCE (saturated calomel electrode)
Hg2Cl2(s) + 2e- = 2Hg + Cl- Eref = 0.244 V vs. NHE
Ag/AgCl
AgCl(s) + e- = Ag(s) + Cl-(aq) Eref = 0.199 V with saturated KCl
Potentials of reference electrodes
E(RHE) = E(NHE) + 0.05916pH
E(SCE) = E(NHE) – 0.2444
E(Ag/AgCl) = E(NHE) – 0.2223
E(Ag/AgCl, sat.KCl) = E(NHE) – 0.196
E(Hg/HgO 1M KOH) = E(NHE) – 0.1100 + 0.05946pH
E(Hg/Hg2SO4) = E(NHE) – 0.6152
Potential vs. energy (vs. vacuum)
예: Potential vs. energy (vs. vacuum)
Controlling potential of the working electrode with respect to the reference 
controlling the energy of the electrons within the working electrode
More negitive potential  energy of electrons is raised  reach a level to
occupy vacant states (LUMO) on species in the electrolyte  flow of electrons
from electrode to solution (a reduction current)
More positive potential  electron flow from solution (HOMO) to electrode
(oxidation current)
Working electrode can act (i) as only a source (for reduction) or a sink (for
oxidation) of electrons transferred to or from species in electrolyte (e.g., C, Au,
Pt, Hg) or can (ii) take part in the electrode reaction, as in dissolution of a metal
M (Zn  Zn2+ + 2e-)
Applying potential from its equilibrium (or its zero-current)
Polarization
Voltammogram: historical one vs. new one
E > 0  working electrode potential > 0 (positive: right of x-axis)
I > 0  oxidation at the working electrode
Polarization: the shift in the voltage across a cell caused by the passage of
current
Departure of the cell potential from the reversible(or equilibrium or
nernstian) potential
Ohmic polarization
Activation polarization
Concentration polarization
Overvoltage (): the voltage shift caused by each kind of polarization
Extent of potential measured by the overpotential:  = E - Eeq
E = En + ohm + act + conc
(i) ohmic polarization
ohm = IRsol, “IR drop”
If free of activation & concentration polarization, slope = 1/Rsol
Rsol = L/A
Rsol = L/A
If free of activation & concentration polarization, slope = 1/Rsol
Electrochemistry needs to minimize ohm
 (conductivity)  ohm  (by adding extra electrolyte: “supporting
electrolyte”)
three-electrode system
two-electrode cell vs. three-electrode cell
Eappl = E + iRs = Eeq +  + iRs
IRs: ohmic drop in the solution (ohmic polarization)  should be minimized 
short distance between working and reference electrode & three-electrode cell
Two-electrode cell: iRs problem due to high current flow
Three-electrode cell: current between WE and auxiliary electrode(or counter
electrode)
Potential measurement between WE and RE  almost no current
to reference electrode
 Potentiostat, etc electrochemical system: three electrode system
(ii) activation polarization
slow electrode reaction  activation polarization; slow kinetics  activation
energy
This can be overcome by increasing the temperature and
by applying extra voltage (activation overvoltage (act))
(iii) concentration polarization
from difference between the electrode surface and bulk concentration
R  O + neconc = E –En = (RT/nF)ln[(cRbcOs)/cRscOb]]
Limiting current
Ideal polarizable electrode (totally polarized electrode): a very large change in
potential upon small current
Ideal nonpolarizable electrode: potential does not change upon passage of current
(e.g., reference electrode)
Double layer
Electrode-solution interface  capacitor “double layer”
Same concept as capacitor (two metal sheets separated with q (coulomb)/E =
C(farad)) ))
Same concept as capacitor (two metal sheets separated with q (coulomb)/E = C(farad))
))
qM = -qS
qM: charge from electrons in metal electrode, qS: charge from ions in solution
charge density
M =qM/A (C/cm2)
double layer capacitance, cd: 10 ~ 40 F/cm2
several models: Helmholtz, Gouy-Chapman, Stern, Grahame model etc
Grahame model: IHP (inner Helmholtz plane, specifically adsorbed anion) +
OHP (outer HP, solvated cation) + diffuse layer
Electrochemical systems in terms of circuit elements
e.g.,) Hg/K+, Cl-/SCE, Hg: ideal polarized electrode
CSCE, Cd: capacitances of SCE and double layer, Rs: solution resistor
CT = CSCECd/(CSCE + Cd), CSCE  Cd  CT  Cd  RC circuit
e.g.,) applying voltage (or potential) step:
potential step: E, EC of capacitor, ER of resistor
q = CdEC
E = ER + EC = iRs + q/Cd
i = dq/dt
dq/dt = -q/(RsCd) + E/Rs
q =0 at t = 0  q = ECd[1 – exp(-t/RsCd)]
By differentiating,
I = (E/Rs)exp(-t/RsCd)
At time constant  = RsCd  current for charging the double layer capacitance
drops to
37 % at  = t, 5 % at  = 3t
e.g.,) Rs = 1 , Cd = 20 F,  = 20 sec  double layer charging is 95 %
complete in 60 sec
sec
sec
Double layer charging pr
Double layer charging process: “non-faradaic process”
Cf) oxidation /reduction  electron transfer ; governed by Faraday’s law (the
amount of chemical reaction caused by the flow of current is proportional to the
amount of electricity passed)  “faradaic process” or “charge transfer process”
Semiconductor electrode
Semiconductor/electrolyte  space charge region due to space charge capacity,
Csc, 0.001 ~ 1 Fcm-2, (cf; Cdl = 10 ~ 100 Fcm-2 )  band bending
n-type SC
when EF of SC lies above that in electrolyte  electron flow from SC (positively
charged) to electrolyte (negatively charged)  bent upward
by applying potential of bulk = surface, band bending & space charge region
disappear  “flat band potential (fb or Efb)”
space charge capacitance Csc  Mott-Schottly equation
1/Csc2 = (2/e0N)1/2(- - kT/e)
: dielectric constant, N: donor or acceptor densities, e: quantity of charge, - =
E-Efb
A plot of 1/Csc2 vs. potential E should be linear  Efb, doping level N
p-type
p-type