PRINCIPLES OF CHEMISTRY II

CHEM 1212

CHAPTER 18

DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences

Clayton state university

CHAPTER 18

ELECTROCHEMISTRY

- Is the study of the relations between chemical reactions and electricity

- Electrochemical processes involve the transfer of electrons from one substance to another

ELECTROCHEMISTRY

- Also called redox reactions- Involve transfer of electrons from one species to another

Oxidation - loss of electronsReduction - 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 air4Fe(s) + 3O2(g) → 2Fe2O3(s)

OXIDATION-REDUCTION REACTIONS

- There is no transfer of electrons from one reactant to another reactant

ExamplesBaCO3(s) → BaO(s) + CO2(g)

Double-replacement reactions

NONREDOX 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 = 0Na(s), O2(g), H2(g), Hg(l)

2. Oxidation number of a monatomic ion = chargeNa+ = +1, Cl- = -1, Ca2+ = +2, Al3+ = +3

3. Oxygen is usually assigned -2H2O, CO2, SO2, SO3

Exceptions: H2O2 (oxygen = -1) and OF2 (oxygen = +2)

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, CO3

2- = -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 = 01 carbon atom and 2 oxygen atoms

1(x) + 2(-2) = 0x = +4

CO2

x -2 for each oxygen

OXIDATION NUMBERS

CH4

x +1

1(x) + 4(+1) = 0x = -4

OXIDATION NUMBERS

NO3-

x -2

1(x) + 3(-2) = -1x = +5

OXIDATION NUMBERS

- 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) + 2e-

The reduction half-reaction is: Cu2+(aq) + 2e- → Cu(s)

HALF-REACTIONS

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

OXIDIZING AND REDUCING AGENTS

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) + C2O4

2-(aq) → MnO2(s) + CO32-(aq)

Answer

2MnO4-(aq) + 3C2O4

2-(aq) + 4OH-(aq)

→

2MnO2(s) + 6CO32-(aq) + 2H2O(l)

BALANCING REDOX EQUATIONS

ELECTRODE

- Conducts electrons into or out of a redox reaction system

Examplesplatinum 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

ExamplesCarbon, Gold, Platinum, ITO

Reactive Electrodes- Participate in the reaction

ExamplesSilver, 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

ExampleRechargeable batteries

Conditions for Non-Reversibility- If one or more of the species decomposes

- If a gas is produced and escapes

- 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 reactionCu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)

VOLTAIC (GALVANIC) CELL

Anode Oxidation

Zn(s) → Zn2+(aq) + 2e-

Cathode Reduction

Cu2+(aq) + 2e- → Cu(s)

Salt bridge (KCl)

Cl-

K+

Voltmeter

- +

e- e-

Cu electrodeZn electrode

Cu2+Zn2+

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 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 = n x F (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)

POTENTIALS OF VOLTAIC CELL

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 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 simplicityConcentrations = 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) Eo = 0.000 VAg+(aq) + e- → Ag(s) Eo = +0.799 V

Reverse the second equation (sign changes)Ag(s) → Ag+(aq) + e- Eo = -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) Eo = +0.339 VZn2+(aq) + 2e- → Zn(s) Eo = -0.762 V

Reverse the second equation (sign changes)Zn(s) → Zn2+(aq) + 2e- Eo = +0.762 V

Combine (electrons cancel)Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq) Eo = +1.101 V

Eo is positive so reaction is spontaneousReverse 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 factorEo remains the same

- For a complete reaction

Ecell = E+ - E-

and

Eo = E+o - E-

o

E+ = potential at positive terminalE- = 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) ReactionCu2+(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

STANDARD POTENTIALS

∆G AND Keq

- Recall that spontaneous reaction has a negative value of ∆G

∆G = -nFE

n = number of moles of electrons transferredF = Faraday constant = 9.6485 x 104 C/mol

E = cell potential (V or J/C)

Under Standard Conditions∆Go = -nFEo

NERNST EQUATION

a

bO

A

Blog

n

0.0591EE

For the half reactionaA + ne- ↔ bB

The half-cell potential (at 25 oC), E, is given by

a

bO

A

Bln

nF

RTEE

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 atmlogQ = 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)

eqo logK

n

0.0591E

Or

/0.05916nEeq

o

10K

Positive Eo implies Keq > 1Negative 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

ExamplesPlatinum

Silver

INDICATOR ELECTRODES

- Respond directly to the analyte

- Two Classes of Indicator Electrodes

- Ion-Selective Electrodes

- Selectively binds one ion (no redox chemistry)

ExamplespH (H+) electrode

Calcium (Ca2+) electrodeChloride (Cl-) electrode

ELECTROLYSIS

- Voltage is applied to drive a redox reaction thatwould 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)