ebbing 20

7/29/2019 Ebbing 20

1/18

81

John A. SchreifelsChemistry 212

Chapter 20-1

Chapter 20

Electrochemistry

7/29/2019 Ebbing 20

2/18

82


Chapter 20-2

Overview

Half-Reactions Balancing oxidation reduction in acidic and basic solutions

Voltaic cells

Construction of voltaic cells

Notation for voltaic cells

Electromotive force (EMF)

Standard cell potentials

Equilibrium constants from EMFs

Concentration dependence of EMF

Electrolytic cells Aqueous electrolysis

Stoichiometry of electrolysis

7/29/2019 Ebbing 20

3/18

83


Chapter 20-3

Electrochemistry

Electrochemistry - Field of

Chemistry dealing with transfer of

electrons from one species to

another.

E.g. Zn in CuSO4(aq).

Electrochemical cell - combination of

two half reactions to produce

electricity from reaction.

E.g. Danielle cell: Zn and Cu

electrodes in salts of these ions.

7/29/2019 Ebbing 20

4/18

84


Chapter 20-4

Balancing Redox Reactions: Oxidation Number

Method

Determine oxidation # for each atom- both sides of equation.

Determine change in oxidation state for each atom.

Left side: make loss of electrons = gain.

Balance on other side.

Insert coefficients for atoms that don't change oxidation state.

E.g. Balance

FeS(s)+CaC2(s)+CaO(s)Fe(s)+CO(g)+CaS(s)

In acidic or basic solution balance as above, then balance

charge with H+ or OH on one side and water on other side.

Cancel waters that appear on both sides at end.

E.g. Balance which occurs in acidic solution:

242 CrOClCrOClO

7/29/2019 Ebbing 20

5/18

85


Chapter 20-5

Balancing: Half-Reaction Method

Write unbalanced half reactions for the oxidation and the

reduction Balance the number of elements except O and H for each.

Balance O's with H2O to the deficient side.

Balance H's with H+ to the hydrogen deficient side

Acidic: add H+

Basic: add H2O to the deficient side and OH to the other side.

Balance charge by adding e to the side that needs it.

Multiply each half-reaction by integers to make electronscancel.

Add the two half-reactions and simplify.

E.g. Balance: Acidic: Zn(s) + VO2+(aq) Zn2+(aq) + V3+(aq).

Basic: Ag(s) + HS(aq) + (aq) Ag2S(s) + Cr(OH)3(s).2

4CrO

7/29/2019 Ebbing 20

6/18

86


Chapter 20-6

Galvanic (Voltaic) and Electrolytic Cells

Cell reaction Redox reaction involved inelectrochemical cell.

Voltaic (galvanic) cell reaction is spontaneous and

generates electrical current.

Electrolytic non-spontaneous reaction occurs dueto passage of current from external power source.

E.g. charging of batteries.

7/29/2019 Ebbing 20

7/18

87


Chapter 20-7

Galvanic Cell 2

anode electrode whereoxidation occurs.

cathode electrode wherereduction occurs. salt bridge ionic solution connecting two half-

cells (half-reactions) to prevent solutions frommixing.

E.g. Which is the anode and cathode in the cellto the right? Write the halfreactions.

Cd(s) + 2Ag+(aq) 2Ag(s) + Cd2+(aq)

Sign of electrodes (current flows from anode tocathode):

E.g. determine direction of electron flow for thereaction for a galvanic cell made from Ni(s) andFe(s). The reaction is:

2Fe3+(aq) + 3Ni 2Fe(s) + 3Ni2+(aq)

Electrode Sign Description1. Cathode: + cations migrate to it.

cations reduced2. Anode: anions migrate to it.

cations made.

7/29/2019 Ebbing 20

8/18

88


Chapter 20-8

Shorthand Notation for Galvanic Cells

Shorthand way of portraying electrodes in a voltaic

electrochemical cell. Redox couple-oxidized and reduced forms of same element

when it is involved in electrochemical reaction. Shorthand:Ox/Red

E.g. Cu2+/Cu, Zn2+/Zn. 2 couples required for electrochemical reaction.

Shorthand rules: Anode reaction-left; reduced form first.

Cathode-right; oxidized form first.

Vertical line drawn between different phases including reaction ofgases at metal electrode.

Double vertical drawn where salt bridge separates two half-

reactions.E.g. draw cell diagram for

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s).

Fe3+(aq) + H2(g) Fe2+(aq) + 2H+(aq).

7/29/2019 Ebbing 20

9/18

89


Chapter 20-9

Electrical Work

Earlier w =

P

V In electrochemistry electrical pressure = potential

difference; w = Eq or charge times the electrical

pressure.

Units: Coulomb

Volts = Joules (SI Units 1J = 1C

V) ;

Also want to relate to # moles.

1 mole e = 1 Faraday = 1 F

F = qeN = 1.602x1019

C6.022x1023

/mol =9.65x104C/mol e.

7/29/2019 Ebbing 20

10/18

810


Chapter 20-10

Cell Potentials for Cell Reactions: Spontaneity of

Redox Reactions

G vs E:

G E G = nFE. Use this to calculate G forelectrochemical reaction when cell voltage known.

E.g. Determine G for Zn/Cu cell if E = 1.100V

G E

At Equilibrium 0 0Spontaneous +

Not spontaneous +

7/29/2019 Ebbing 20

11/18

811


Chapter 20-11

Standard Reduction Potentials

As with thermodynamic quantities, we list cell potentials at standard state = 1M

at 1 atm and usually 25C. Cell potential is the sum of half-cell potentials using Hess law.

Half-cell reactions for Daniell cell were

Potential at each electrode initially determined relative to SHE = Standardhydrogen electrode.

2H+(aq)+2e H2(g); [H+] = 1M and PH2=1atm.

Other reaction run to determine if the SHE reaction proceeds spontaneously indirection written when connected to other half-cells.

E.g. the cell potential of copper at standard state conditions relative toSHE(acting as the anode) was 0.340 V; determine the halfcell potential for Zn Zn2+ if the potential for the Daniell cell (standard state conditions) was 1.100V.

Half-cell reactions reported as reductions.

Zn(s) Zn2+

(aq) + 2e Anode

Cu

2+

(aq) + 2e

Cu(s) Cathode

Zn(s) +Cu2+

(aq) Zn2+

(aq) + Cu(s)

oanode

ocathode

ocell

EEE

7/29/2019 Ebbing 20

12/18

812


Chapter 20-12

Using Standard Reduction Potentials

Largenegative value meansoxidation strongly favored; strong reducing

agent. Largepositive value meansreduction strongly favored; strong oxidizingagent.

Relative values in table give an indication that one half-reaction favored overother. Summing half-cell reactions allow determination of standard cell potential.

Half-cell potential intensive property independent of amount of material wedont use stoichiometric coefficients for determining standard cellpotentials.

E.g. determine the cell potential ofBr2(l) + 2I

(aq) I2(l) + 2Br(aq)

E.g.2 determine the cell potential of

2Ag+(aq) + Cu(s) 2Ag(s) + Cu2+(aq)

E.g.3 Determine cell potential:

(aq) + Fe(s) Fe2+(aq) + Mn2+(aq) (balanced?).

when it is operated galvanically. Which is the oxidizing agent? reducing agent? E.g. 4 Determine if the reaction below is spontaneous in the direction written.

Fe3+(aq) + Ag(s) ?

4MnO

7/29/2019 Ebbing 20

13/18

813


Chapter 20-13

Spontaneity of Redox Reactions

G vs E:

G E G = nFE. Use this to calculate G forelectrochemical reaction when cell voltage known.

E.g. Determine G for Zn/Cu cell if E = 1.100V

G E

At Equilibrium 0 0Spontaneous +

Not spontaneous +

7/29/2019 Ebbing 20

14/18

814


Chapter 20-14

Effect of Concentration on Cell EMF: The Nernst

Equation

Recall that G = Go + RTlnQ where

or at 25C

which is called Nernst Equation.

E.g. Determine potential of Daniell cell at 25C if

[Zn2+] = 0.100 M and [Cu2+] = 0.00100 M.

ba

nm

]B[]A[

]N[]M[Q

QlnnF

RTEE o

Qlogn

0592.0EE o

7/29/2019 Ebbing 20

15/18

815


Chapter 20-15

Electrochemical Determination of pH

Electrodes can be used to determine acidity ofsolution by using hydrogen electrode with anotherone e.g. Hg2Cl2 half-cell.

E.g. determine the pH of a solution that develops acell potential of 0.280 V (at 25C) given the cell below

Pt(s) | H2(g) (1atm) | H+(? M) || Pb2+(1 M) | Pb(s)

E.g. 2 determine the pH of a solution that develops a

cell potential of0.200 V (at 25C) given the cellbelow (called a concentration cell).

Pt(s)|H2(g)(1atm)|H+(1.00M)||H+(?M)|H2(g)(1atm)|Pt(s)

pH.

05920E'

]log[H0.0592'EE

7/29/2019 Ebbing 20

16/18

816


Chapter 20-16

Standard Cell Potentials and Equilibrium

Constants

Free energy and equilibrium constant for reaction studied can

be determined from cell voltage.

Cell potential can be determined from G or from equilibrium

constant.

Recall: G = nFE = RTlnK.

E.g. Determine free energy and equilibrium constant for reaction

below (unbalanced).

(aq) + Fe(s) Fe2+(aq) + Mn2+(aq)

E.g.2 Determine cell potential and equilibrium constant of Cl2/Br2

cell.

E.g.3 The following cell has a potential of 0.578 V at 25C;determine Ksp.

Ag(s)|AgCl(s)|Cl(1.0 M)||Ag+(1.0 M)|Ag(s).

4MnO

7/29/2019 Ebbing 20

17/18

817


Chapter 20-17

Quantitative Aspects of Electrolysis

Current, i, measured units: 1 ampere = 1 coulomb per s (1 A = 1 Cs1).

Time of electrolysis, t (s), also measured. Total charge, Q, calculated from the product:

Q = it

Charge on 1 mol of e:

mol of e in an electrolysis obtained from balanced cell reaction:

E.g. determine # mol of electron involved in the electrolysis of the following: Ag+(aq) + e Ag(s)

2Cl(aq) + 2e Cl2(g)

Amount deposited given by:

E.g. Determine amount of Cu2+ electrolyzed from solution at constant current of6.00 A for period of 1.00 hour.

eofmolC96500

e

C10x602.1

eofmol

e10x02.6

eNF

1923

nF

ti=

nF

Q=moles

7/29/2019 Ebbing 20

18/18

818


Chapter 20-18

Electrical Work

Maximum electrical work: G = wmax = nFE

n = mol of electrons

F = Faradays constant

E = cell potential

Units Joules

Electrical Power: 1 watt = 1 J/s. Energy often expressed as kilowatthr

1 kW*hr = 1000 W*3600 s = 3.6x106 J

E.g. determine the maximum work in kW*hr required to produce

1.00 kg of Zn from Zn2+ in a Daniell cell where the cell potential

is 1.100 V for the production of Zn metal.Strategy:

Determine n: mol of Zn2+ times 2.

Calculate work.

ebbing 20

Documents