Download - PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 18

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


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)


Half-Reaction Method

Basic Solutions


- 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


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)


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)


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


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


For the overall reactionCu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)


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)


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 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)