Transcript Electrochemistry - Faculty

PRINCIPLES OF CHEMISTRY II
CHEM 1212
CHAPTER 18
DR. AUGUSTINE OFORI AGYEMAN
Assistant professor of chemistry
Department of natural sciences
Clayton state university
CHAPTER 18
ELECTROCHEMISTRY
ELECTROCHEMISTRY
- Is the study of the relations between chemical
reactions and electricity
- Electrochemical processes involve the transfer of electrons
from one substance to another
OXIDATION-REDUCTION REACTIONS
- Also called redox reactions
- Involve transfer of electrons from one species to another
Oxidation - loss of electrons
Reduction - gain of electrons
- Ionic solid sodium chloride (Na+ and Cl- ions) formed from
solid sodium and chlorine gas
2Na(s) + Cl2(g) → 2NaCl(s)
- The oxidation (rusting) of iron by reaction with moist air
4Fe(s) + 3O2(g) → 2Fe2O3(s)
NONREDOX REACTIONS
- There is no transfer of electrons from one reactant
to another reactant
Examples
BaCO3(s) → BaO(s) + CO2(g)
Double-replacement reactions
OXIDATION NUMBER (STATE)
The concept of oxidation number
- provides a way to keep track of electrons in redox reactions
- not necessarily ionic charges
Conventionally
- actual charges on ions are written as n+ or n- oxidation numbers are written as +n or -n
Oxidation - increase in oxidation number (loss of electrons)
Reduction - decrease in oxidation number (gain of electrons)
OXIDATION NUMBERS
1. Oxidation number of uncombined elements = 0
Na(s), O2(g), H2(g), Hg(l)
2. Oxidation number of a monatomic ion = charge
Na+ = +1, Cl- = -1, Ca2+ = +2, Al3+ = +3
3. Oxygen is usually assigned -2
H2O, CO2, SO2, SO3
Exceptions: H2O2 (oxygen = -1) and OF2 (oxygen = +2)
OXIDATION NUMBERS
4. Hydrogen is usually assigned +1
Exceptions: -1 when bonded to metals
(+1in HCl, NH3, H2O and -1in CaH2, NaH)
5. Halogens are usually assigned -1 (F, Cl, Br, I)
Exceptions: when Cl, Br, and I are bonded to oxygen or a more
electronegative halogen
(Cl2O: O = -2 and Cl = +1)
6. The sum of oxidation numbers for
- neutral compound = 0
- polyatomic ion = charge
(H2O = 0, CO32- = -2, NH4+ = +1)
OXIDATION NUMBERS
CO2
The oxidation state of oxygen is -2
CO2 has no charge
The sum of oxidation states of carbon and oxygen = 0
1 carbon atom and 2 oxygen atoms
1(x) + 2(-2) = 0
x = +4
CO2
x
-2 for each oxygen
OXIDATION NUMBERS
CH4
x
+1
1(x) + 4(+1) = 0
x = -4
OXIDATION NUMBERS
NO3x
-2
1(x) + 3(-2) = -1
x = +5
HALF-REACTIONS
- Just the oxidation or the reduction is given
- The transferred electrons are shown
Oxidation Half-Reaction
- Electrons are on the product side of the equation
Reduction Half-Reaction
- Electrons are on the reactant side of the equation
HALF-REACTIONS
For the redox reaction
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Zn is oxidized (oxidation number changes from 0 to +2)
Cu is reduced (oxidation number changes from +2 to 0)
The oxidation half-reaction is: Zn(s) → Zn2+(aq) + 2eThe reduction half-reaction is: Cu2+(aq) + 2e- → Cu(s)
OXIDIZING AND REDUCING AGENTS
Oxidizing Agent
- Is the reduced species (accepts electrons from another species)
Reducing Agents
- Is the oxidized species (donates electrons to another species)
For the redox reaction
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Cu is reduced so is the oxidizing agent
Zn is oxidized so is the reducing agent
BALANCING REDOX EQUATIONS
Half-Reaction Method
Acidic Solutions
BALANCING REDOX EQUATIONS
- Write separate oxidation and reduction half-reactions
- Balance all the elements except hydrogen and oxygen in each
- Balance oxygen using H2O(l), hydrogen using H+(aq), and
charge using electrons (e-)
- Multiply both half-reactions by suitable factors to equalize
electron count
- Combine the balanced half-reactions
BALANCING REDOX EQUATIONS
Balance the following redox reaction in acid medium
MnO4-(aq) + Fe2+(aq) → Fe3+(aq) + Mn2+(aq)
Answer
MnO4-(aq) + 5Fe2+(aq) + 8H+(aq)
→
5Fe3+(aq) + Mn2+(aq) + 4H2O(l)
BALANCING REDOX EQUATIONS
Half-Reaction Method
Basic Solutions
BALANCING REDOX EQUATIONS
- Balance equation as if it were acidic
- Note H+ ions and add same number of OH- ions to both sides
- Cancel H+ and OH- (=H2O) with H2O on other side
BALANCING REDOX EQUATIONS
Balance the following redox reaction in basic medium
MnO4-(aq) + C2O42-(aq) → MnO2(s) + CO32-(aq)
Answer
2MnO4-(aq) + 3C2O42-(aq) + 4OH-(aq)
→
2MnO2(s) + 6CO32-(aq) + 2H2O(l)
ELECTRODE
- Conducts electrons into or out of a redox reaction system
Examples
platinum wire
carbon (glassy or graphite)
indium tin oxide (ITO)
Electroactive Species
- Donate or accept electrons at an electrode
ELECTRODE
Chemically Inert Electrodes
- Do not participate in the reaction
Examples
Carbon, Gold, Platinum, ITO
Reactive Electrodes
- Participate in the reaction
Examples
Silver, Copper, Iron, Zinc
CHEMICAL CHANGE
Spontaneous Process
- Takes place with no apparent cause
Nonspontaneous Process
- Requires something to be applied in order for it to occur
(usually in the form of energy)
VOLTAIC (GALVANIC) CELL
- Spontaneous reaction
- Produces electrical energy from chemical energy
- Can be reversed electrolytically for reversible cells
Example
Rechargeable batteries
Conditions for Non-Reversibility
- If one or more of the species decomposes
- If a gas is produced and escapes
VOLTAIC (GALVANIC) CELL
- A spontaneous redox reaction generates electricity
- One reagent is oxidized and the other is reduced
- The two reagents must be separated (cannot be in contact)
- Each is called a half-cell
- Electrons flow through a wire (external circuit)
VOLTAIC (GALVANIC) CELL
Oxidation Half reaction
- Loss of electrons
- Occurs at anode (negative electrode)
- The left half-cell by convention
Reduction Half Reaction
- Gain of electrons
- Occurs at cathode (positive electrode)
- The right half-cell by convention
VOLTAIC (GALVANIC) CELL
Salt Bridge
- Connects the two half-cells (anode and cathode)
- Filled with gel containing saturated aqueous salt solution (KCl)
- Ions migrate through to maintain electroneutrality
- Prevents charge buildup that may cease the reaction process
VOLTAIC (GALVANIC) CELL
For the overall reaction
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
e-
Voltmeter
-
Zn electrode
e-
+
Cu electrode
ClK+
Zn2+
Salt bridge (KCl)
Anode
Oxidation
Zn(s) → Zn2+(aq) + 2e-
Cu2+
Cathode
Reduction
Cu2+(aq) + 2e- → Cu(s)
POTENTIALS VOLTAIC CELL
Voltage or Potential Difference (E)
- Referred to as the electromotive force (emf)
- Is the voltage measured
- Measured by a voltmeter (potentiometer) connected to electrodes
Greater Voltage
- More favorable net cell reaction
- More work done by flowing electrons (larger emf)
POTENTIALS OF VOLTAIC CELL
Voltage or Potential Difference (E)
- Work done by or on electrons when they
move from one point to another
Units: volts (V or J/C)
Work (J) = E (V) x q (C)
POTENTIALS OF VOLTAIC CELL
Charge
Charge (q) of an electron = - 1.602 x 10-19 C
Charge (q) of a proton = + 1.602 x 10-19 C
C = coulombs
Charge of one mole of electrons
= (1.602 x 10-19 C)(6.022 x 1023/mol) = 9.6485 x 104 C/mol
= Faraday constant (F)
q=nxF
(n = number of moles)
POTENTIALS OF VOLTAIC CELL
Current
- The quantity of charge flowing past a point in an
electric circuit per second
Units
Ampere (A) = coulomb per second (C/s)
Charge (C) = current (A) x time (s)
STANDARD POTENTIALS
Electrode Potentials
- A measure of how willing a species is to gain or lose electrons
Positive Voltage (spontaneous process)
- Electrons flow into the negative terminal of voltmeter
(flow from negative electrode to positive electrode)
Negative Voltage (nonspontaneous process)
- Electrons flow into the positive terminal of voltmeter
(flow from positive electrode to negative electrode)
Conventionally
- Negative terminal is on the left of galvanic cells
STANDARD POTENTIALS
Standard Reduction Potential (Eo)
- Used to predict the voltage when different cells are connected
- Potential of a cell as cathode compared to
standard hydrogen electrode
- Species are solids or liquids
- Activities = 1
- We will use concentrations for simplicity
Concentrations = 1 M
Pressures = 1 atm
STANDARD POTENTIALS
Standard Hydrogen Electrode (SHE)
- Used to measure Eo for half-reactions
- Connected to negative terminal
- Pt electrode
- Acidic solution in which [H+] = 1 M
- H2 gas (1 bar) is bubbled past the electrode
H+(aq, 1 M) + e- ↔ 1/2H2 (g, 1 atm)
Conventionally, Eo = 0 for SHE
STANDARD POTENTIALS
The Eo for Ag+(aq) + e- ↔ Ag(s) is +0.799 V
Implies
- if a sample of silver metal is placed in a 1 M Ag+ solution, a
value of 0.799 V will be measured with S. H. E. as reference
STANDARD POTENTIALS
Silver does not react spontaneously with hydrogen
2H+(aq) + 2e- → H2(g)
Ag+(aq) + e- → Ag(s)
Eo = 0.000 V
Eo = +0.799 V
Reverse the second equation (sign changes)
Ag(s) → Ag+(aq) + eEo = -0.799 V
Multiply the second equation by 2 (Eo is intensive so remains)
2Ag(s) → 2Ag+(aq) + 2e- Eo = -0.799 V
Combine (electrons cancel)
2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g) Eo = -0.799 V
STANDARD POTENTIALS
Consider Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
Cu2+(aq) + 2e- → Cu(s)
Zn2+(aq) + 2e- → Zn(s)
Eo = +0.339 V
Eo = -0.762 V
Reverse the second equation (sign changes)
Zn(s) → Zn2+(aq) + 2eEo = +0.762 V
Combine (electrons cancel)
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = +1.101 V
Eo is positive so reaction is spontaneous
Reverse reaction is nonspontaneous
STANDARD POTENTIALS
- Half-reaction is more favorable for more positive Eo
- Refer to series and Eo values in textbook
- For combining two half reactions, the one higher in the series
proceeds spontaneously as reduction under standard conditions
- The one lower in the series proceeds spontaneously as
oxidation under standard conditions
Formal Potential
- The potential for a cell containing a specified concentration of
reagent other than 1 M
STANDARD POTENTIALS
- When a half reaction is multiplied by a factor
Eo remains the same
- For a complete reaction
Ecell = E+ - Eand
Eo = E+o - E-o
E+ = potential at positive terminal
E- = potential at negative terminal
STANDARD POTENTIALS
For the Cu – Fe cell at standard conditions
Cu2+ + 2e- ↔ Cu(s)
0.339 V
Fe2+ + 2e- ↔ Fe(s)
-0.440 V
Ecell = 0.779 V
Galvanic (overall) Reaction
Cu2+(aq) + Fe(s) ↔ Cu(s) + Fe2+(aq)
STANDARD POTENTIALS
- Positive E implies spontaneous forward cell reaction
- Negative E implies spontaneous reverse cell reaction
If cell runs for long
- Reactants are consumed
- Products are formed
- Equilibrium is reached
- E becomes 0
- Reason why batteries run down
∆G AND Keq
- Recall that spontaneous reaction has a negative value of ∆G
∆G = -nFE
n = number of moles of electrons transferred
F = Faraday constant = 9.6485 x 104 C/mol
E = cell potential (V or J/C)
Under Standard Conditions
∆Go = -nFEo
NERNST EQUATION
For the half reaction
aA + ne- ↔ bB
The half-cell potential (at 25 oC), E, is given by
RT  B b 

EE 
ln 
a 
nF  A 
O
b




0.0591
B

E  EO 
log 
a 
n
 A 
NERNST EQUATION
Eo = standard electrode potential
R = gas constant = 8.314 J/K-mol
T = absolute temperature
F = Faraday’s constant = 9.6485 x 104 C/mol
n = number of moles of electrons transferred
NERNST EQUATION
- The standard reduction potential (Eo)
is when [A] = [B] = 1M
- [B]b/[A]a = Q = reaction quotient
- Concentration for gases are expressed as pressures in atm
- Q = 1 for [ ] = 1 M and P = 1 atm
logQ = 0 and E = Eo
- Pure solids, liquids, and solvents do not appear in Q expression
NERNST EQUATION
At cell equilibrium at 25 oC
E = 0 and Q = Keq (the equilibrium constant)
0.0591
E 
logK eq
n
o
Or
Keq  10
nEo /0.05916
Positive Eo implies Keq > 1
Negative Eo implies Keq < 1
REFERENCE ELECTRODES
- Provide known and constant potential
Examples
Silver-silver chloride electrode (Ag/AgCl)
Saturated Calomel electrode (SCE)
INDICATOR ELECTRODES
- Respond directly to the analyte
- Two Classes of Indicator Electrodes
Metal Electrodes
- Surfaces on which redox reactions take place
Examples
Platinum
Silver
INDICATOR ELECTRODES
- Respond directly to the analyte
- Two Classes of Indicator Electrodes
- Ion-Selective Electrodes
- Selectively binds one ion (no redox chemistry)
Examples
pH (H+) electrode
Calcium (Ca2+) electrode
Chloride (Cl-) electrode
ELECTROLYSIS
- Voltage is applied to drive a redox reaction that
would not otherwise occur
Examples
- Production of aluminum metal from Al3+
- Production of Cl2 from Cl- Electroplating
ELECTROLYSIS CELL
- Consists of two electrodes in an electrolyte solution
- Nonspontaneous reaction
- Requires electrical energy to occur
CORROSION
- Oxidation of a metal to produce compounds of the metal
Examples
- Rusting of iron (Fe forms Fe2O3·xH2O
- Copper in bronze forming copper(II) compounds (green color)
Prevention
- Painting or coating
- Plating of iron with chromium
- Anodic protection (metal is oxidized under controlled conditions)
- Cathodic protection (a more reactive metal is placed in contact)