electrochemistry chapter 4 4.4 and 4.8 chapter 19 19.1-19.5 and 19.8

Electrochemistry

Chapter 44.4 and 4.8Chapter 1919.1-19.5 and 19.8

Oxidation-Reduction Reactions

Electron Transfer Reactions Many redox reactions take place in water. Half-reactions OILRIG


2Mg (s) + O2 (g) 2MgO (s)

2Mg 2Mg2+ + 4e- Oxidation half-reaction (lose e-)

O2 + 4e- 2O2- Reduction half-reaction (gain e-)

2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e-

2Mg + O2 2MgO


Oxidizing Agent- a substance that accepts electrons from another substance, causing the other substance to be oxidized.

Reducing Agent- a substance that donates electrons to another substance, causing it to become reduced.


Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)


Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)

Zn Zn2+ + 2e- Zn is oxidized

Zn is the reducing agent

Cu2+ + 2e- Cu Cu2+ is reduced

Cu2+ is the oxidizing agent


Copper wire reacts with silver nitrate to form silver metal.What is the oxidizing agent in the reaction?

Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)

Cu Cu2+ + 2e-

Ag+ + 1e- Ag Ag+ is reduced

Ag+ is the oxidizing agent

Oxidation Number

Oxidation Number- the number of charges the atom would have in a molecule (or an ionic compound) if electrons were transferred completely.

H2(g) + Cl2(g) → 2HCl(g)

0 0 +1 -1

Assigning Oxidation Numbers

1. Free elements (uncombined state) have an oxidation number of zero.

Na, Be, K, Pb, H2, O2, P4 = 0

2. In monatomic ions, the oxidation number is equal to the charge on the ion.

Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2

3. The oxidation number of oxygen is usually –2. In H2O2

and O22- it is –1.


4. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.

5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.

6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.


Oxidation numbers of all the elements in HCO3

- ?

HCO3-

O = -2 H = +1

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

C = +4

Redox Titrations

Oxidizing agent/Reducing Agent Equivalence point reached when reducing

agent is completely oxidized by the oxidizing agent.

Indicator is needed

Redox Titrations

Electrochemistry

Branch of Chemistry that deals with interconversion of electrical and chemical energy.

ImportanceUse of electricity in everyday life Interconversion of energy Study of electrochemical processes

ElectrochemistryElectrochemical processes are oxidation-reduction reactions

in which:

• the energy released by a spontaneous reaction is converted to electricity or

• electrical energy is used to cause a nonspontaneous reaction to occur

Balancing Redox Equations

Simple reactions easy to balance Complex reactions

Usually encountered in laboratoryCan include CrO4

2-, Cr2O72-, MnO4

-, NO3- and

SO42-

Follow balancing guidelines

Balancing Redox Reactions

Guidelines Step 1: Write the unbalanced equation for the reaction in ionic

form. Step 2: Separate the equation into two half-reactions Step 3: Balance each half-reaction for number and type of atoms

and charges. (Look at medium) Step 4: Add the two half-reactions together and balance the final

equation by inspection. The electrons on both sides must cancel. (Be sure they are equal)

Step 5: Verify that the equation contains the same type and numbers of atoms and the same charges on both sides of the equation.

Balancing Redox Equations

1. Write the unbalanced equation for the reaction ion ionic form.

The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?

Fe2+ + Cr2O72- Fe3+ + Cr3+

2. Separate the equation into two half-reactions.

Oxidation:

Cr2O72- Cr3+

+6 +3

Reduction:

Fe2+ Fe3++2 +3

3. Balance the atoms other than O and H in each half-reaction.

Cr2O72- 2Cr3+

Balancing Redox Reactions4. For reactions in acid, add H2O to balance O atoms and H+ to

balance H atoms.Cr2O7

2- 2Cr3+ + 7H2O14H+ + Cr2O7

2- 2Cr3+ + 7H2O

5. Add electrons to one side of each half-reaction to balance the charges on the half-reaction.

Fe2+ Fe3+ + 1e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

6. If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.

6Fe2+ 6Fe3+ + 6e-

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

Balancing Redox Reactions7. Add the two half-reactions together and balance the final

equation by inspection. The number of electrons on both sides must cancel.

6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O

6Fe2+ 6Fe3+ + 6e-Oxidation:

Reduction:

14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O

8. Verify that the number of atoms and the charges are balanced.

14x1 – 2 + 6x2 = 24 = 6x3 + 2x3

9. For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation.

Galvanic Cells

Oxidation-Reduction Reaction Oxidizing agent and Reducing agent are not in

contact External conducting medium Galvanic Cell/Volatic Cell- the experimental

apparatus for generating electricity through the use of a spontaneous reaction.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

e– e–

Zincanode

Coppercathode

Saltbridge

Cottonplugs

ZnSO4

solution

Zn is oxidized toZn2+ at anode.

2e– + Cu2+ (aq) Cu(s)

CuSO4

solution

Zn2+

SO 42–

Cl– K+

Voltmeter

(aq) + Cu(s)

Zn2+

2e–Cu2+

Cu

SO42–

Zn(s) Zn2+ (aq) + 2 e–

Net reaction

Cu

Cu2+ is reducedto Cu at cathode.

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

2+

Galvanic Cells

Cell Voltage- the difference in electrical potential between the anode and the cathode. Measured by a voltmeter.

Electromotive force (emf) and Cell Potential Voltage of a cell depends on:

Composition of electrode and ions Concentration of ions Temperature

Galvanic Cells

Cell Diagram

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

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)

anode cathode

Standard Reduction Potentials

Used to calculate the emf of a galvanic cell.

Use the values from anode and cathode

Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.

Standard emf (E0 )cell

E0 = Ecathode - Eanodecell0 0


Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)

Anode (oxidation): Zn (s) Zn2+ (1 M) + 2e-

Cathode (reduction): 2e- + 2H+ (1 M) H2 (1 atm)

Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)

Standard Reduction PotentialsE0 = 0.76 Vcell

Standard emf (E0 )cell

0.76 V = 0 - EZn /Zn 0

2+

EZn /Zn = -0.76 V02+

Zn2+ (1 M) + 2e- Zn E0 = -0.76 V

E0 = EH /H - EZn /Zn cell0 0

+ 2+2


Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)


19.3

Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)

2e- + Cu2+ (1 M) Cu (s)

H2 (1 atm) 2H+ (1 M) + 2e-Anode (oxidation):

Cathode (reduction):

H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)


E0 = 0.34 Vcell

Ecell = ECu /Cu – EH /H 2+ +2

0 0 0

0.34 = ECu /Cu - 00 2+

ECu /Cu = 0.34 V2+0

Tips

E0 is for the reaction as written The more positive E0 the greater the tendency

for the substance to be reduced The half-cell reactions are reversible The sign of E0 changes when the reaction is

reversed Changing the stoichiometric coefficients of a

half-cell reaction does not change the value of E0

Spontaneity of Redox Reactions

Relationship between Eºcell, ΔGº and K.

G0 = -nFEcell0 n = number of moles of electrons in reaction

F = 96,500J

V • mol = 96,500 C/mol

=0.0257 V

nln KEcell

0 =0.0592 V

nlog KEcell

0

Spontaneity of Redox Reactions

Relationships between Eºcell, ΔGº and K

The Effect of Concentration on Cell Emf Not all reactions in galvanic cells can

occur under standard state conditions. Nernst Equation- equation used to

calculate cell potential under nonstandard state conditions.

-0.0257 V

nln QE0E = -

0.0592 Vn

log QE0E =

Will the following reaction occur spontaneously at 250C if [Fe2+] = 0.60 M and [Cd2+] = 0.010 M? Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)

2e- + Fe2+ 2Fe

Cd Cd2+ + 2e-Oxidation:

Reduction: n = 2

E0 = -0.44 – (-0.40)

E0 = -0.04 V

E0 = EFe /Fe – ECd /Cd0 0

2+ 2+

-0.0257 V

nln QE0E =

-0.0257 V

2ln -0.04 VE =

0.0100.60

E = 0.013

E > 0 Spontaneous

19.5

Concentration Cells

Galvanic cells consisting of the same type of electrodes in the same solution. Each solution is a different concentration.

Zn(s) | Zn2+(0.10M) || Zn2+(1.0M) | Zn(s)

E = Eº - 0.0257V/ 2 x ln [Zn2+] dil / [Zn2+]conc

E = 0- 0.0257V/2 ln (0.10/1.0) E = 0.0296 V

Concentration Cells

ImportanceUnequal ion concentrationsMembrane potentialPropagation

Electrolysis

Electrolysis- is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur.

Electrolytic Cell- an apparatus for carrying out electrolysis.

electrochemistry chapter 4 4.4 and 4.8 chapter 19 19.1-19.5 and 19.8

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