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Essentially dealing with OXIDATION-REDUCTION REACTIONS

Electron transfer reaction

SPONTANEOUS REACTIONS: Examples: voltaic cells, batteries.

NON-SPONTANEOUS REACTIONS:

Examples: electrolysis, electrolytic cells. QUANTITATE REACTIONS

How much current flows 6(,7&8%'98"&:;<%'=(6(

Electrochemistry involves the relationship between electrical energy and chemical energy.

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Ionic equation: Net ionic equation:( What is oxidized? What is reduced? What is the oxidizing agent? What is the reducing agent?

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

Look at this oxidation reduction reaction

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VA9C(%"&('I7&<(E9'(G&%&':;A;AM(9F;G#H9A(4%#%&<(1.  Oxidation state of atom in elemental form is zero.

e.g. Cl2 O2 P4 C(s) S8 2. The oxidation number of a monatomic ion equals its charge. 3. Some elements have “common” oxidation numbers that can be used as reference in

determining the oxidation numbers of other atoms in the compound.(Alkali metals +1 Alkaline earth metals +2 Fluorine –1 H usually +1 (Hydrides: metal-H compounds (–1)) O usually –2 (peroxides (-1) & superoxides possible) Cl, Br, I almost always –1

4. Sum of oxidation numbers is equal to overall charge of molecule or ion: •  For a neutral compound the sum of oxidation numbers equals zero. •  For a polyatomic ion, the sum of the oxidation numbers is equal to the charge on the

ion. 5. Shared electrons are assigned to the more electronegative atom of the pair:

•  more electronegative atom will have a negative oxidation number.

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For main group elements (s and p block) The highest possible positive oxidation state is

equal to the group number!

(We won’t worry about the

transition metals right now)

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Driving force: atoms tend to lose or gain electrons to achieve:

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It helps to know the Periodic trends in oxidation states

Can we explain these common oxidation states?

H -1 +1

Li +1

Be +2

F -1

Na +1

Mg +2

Al +3

P +3 +5

Cl -1

+1,3,5,7

K +1

Ca +2

Ga +3

Ge +2 +4

As +3 +5

Br -1

+1,3,5,7

Rb +1

Sr +2

In +1 +3

Sn +2 +4

Sb +3 +5

I -1

+1,3,5,7

Ba +2

Tl +1 +3

Pb +2 +4

Bi +3 +5

O -2,-1

S -2

+2,4,6 Se -2

+2,4,6 Te -2

+2,4,6

Blue: most common oxidation states in Group 5 Z(,7&8%'98"&:;<%'=(6(2>(?9@#A(

1.  Assign Oxidation Numbers.

2.  Write incomplete half-reactions.

3.  Balance each half-reaction separately. a. Balance atoms undergoing redox. b. Balance remaining atoms.

i. Add H2O to balance oxygens. ii. Add H+ to balance hydrogens.

4.  Balance charges by adding electrons. 5.  Multiply each half-reaction so that the same number of electrons

are involved in the reduction and the oxidation.

6.  Add the half-reactions.

7.  In basic solutions, add OH! to neutralize H+

This is the method for BALANCING REDOX REACTIONS

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Example: Balancing Redox Reactions in acid(

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Example: Balancing Redox Reactions in base(

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Example: Balancing Redox Reactions (

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SPONTANEOUS ELECTROCHEMICAL REACTIONS can do work.

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Voltaic Cells Consist of: (OJ  ,7&8%'9G&<+(((

((((((((."&(%C9(<97;G(:&%#7<(K8#%"9G&(#AG(#A9G&L(((

((((((\"#%("#$$&A<(#%(%"&(OA9G&f((((

((((((\"#%("#$$&A<(#%(%"&(!#%"9G&f(((?J((,7&8%'97=%&(<97IH9A<((!J((\"#%(G9&<(%"&(<#7%(D';GM&(9'($9'9I<(G;B;G&'(G9f(((

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Which electrode will increase in mass? Which will decrease?

Which direction do the electrons flow?

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Cell voltage (EMF or Ecell) is the measure of reaction spontaneity (

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Cell voltage depends on:

1)

2)

3)

The more spontaneous a reaction, !  !  ! 

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The standard potential for an electrochemical cell is the potential (voltage) generated when reactants and

products of a redox reaction are in their standard states.(

It is convenient to tabulate redox reactions as half reactions. Standard half-cell potentials When all substances are in standard state:

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Standard State defined so potentials can be tabulated: T = 25°C. Gases, p = 1 atm. [Solutions] = 1M

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Reduction Half-reaction Eo1/2(V)

F2 (g) + 2e– ! 2F– (aq) +2.87 H2O2 (aq) + 2H+(aq) + 2e– !2H2O(!) +1.776

Au+(aq) + e– ! Au(s) +1.69 PbO2(s) + 4H+(aq) + SO4

2– (aq) + 2e– ! PbSO4(s) + 2H2O(!) +1.685

Cl2 (g) + 2e– ! 2Cl–(aq) +1.359 O2 (g) + 4H+(aq) + 4e– ! 2H2O(!) +1.23

ClO4– (aq) + 2H+(aq) + 2e– ! ClO3

– (aq) + H2O(!) +1.23 Pt2+(aq) + 2e– ! Pt(s) +1.20 2IO3

–(aq) + 12H+(aq) + 10e– ! I2(s) + 6H2O(!) +1.195 Br2 (!) + 2e– ! 2Br–(aq) +1.065

NO3– (aq) + 4H+(aq) + 3e– ! NO(s) + 2H2O(!) +0.96

Ag+(aq) + e– ! Ag(s) +0.799 Fe3+(aq) + e– ! Fe2+(aq) +0.771 O2 (g) + 2H+(aq) + 2e– ! H2O2(aq) +0.68 I2 (s) + 2e– ! 2I–(aq) +0.536 I3

– (aq) + 2e– ! 3I–(aq) +0.53

Cu+(aq) + e– ! Cu(s) +0.521 Cu2+(aq) + 2e– ! Cu(s) +0.337 Sn4+(aq) + 2e–! Sn2+(aq) +0.154 Cu2+(aq) + e– ! Cu+(aq) +0.153 2H+ (aq) + 2e– ! H2(g) +0.00 Pb2+(aq) + 2e– ! Pb(s) –0.126 Sn2+(aq) + 2e– ! Sn(s) –0.136 Co+2(aq) + 2e–! Co(s) "0.277 Ni2+(aq) + 2e– ! Ni(s) –0.28 Cd2+(aq) + 2e– ! Cd(s) –0.403 Fe2+(aq) + 2e– ! Fe(s) –0.440 Cr3+(aq) + 3e– ! Cr(s) –0.74 Zn2+(aq) + 2e– ! Zn(s) –0.763 2H2O (!) + 2e– ! H2(g) + 2OH– (aq) –0.83 SO4

2– (aq) + H2O(!) + 2e– ! SO32– (aq) + 2OH– (aq) –0.93

Mn2+(aq) + 2e– ! Mn(s) –1.18 Al3+(aq) + 3e– ! Al(s) –1.66 Mg2+(aq) + 2e– ! Mg(s) –2.37 Na+(aq) + e– ! Na(s) –2.71 Ca2+(aq) + 2e– ! Ca(s) –2.87 Li+(aq) + e– ! Li(s) –3.05

The 1O-_]!,--(b0.,a.3O-(is the potential associated with the half-reaction.(

Rules for half-cell potentials: 1.  The sum of two half-cells potentials in a cell equals the overall cell potential:

E°cell = E°1/2(oxid) + E°1/2(reduc)

2.  For any half-reaction: E°1/2(oxid) = ! E°1/2(reduc)

3. Standard half-cell is a hydrogen electrode:((

H2(g,1atm) " 2H+ (aq, 1M) + 2e!

E°1/2(oxid) = E°1/2(reduc) = 0 V

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E0cell = E°1/2(reduc) + E°1/2(oxid) =

What is E°1/2(reduc) ? What is E°1/2(oxid)?

What is E0cell for Cu/Zn cell?

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E0cell = E°1/2(reduc) + E°1/2(oxid) =

What is E°1/2(reduc) ? What is E°1/2 (oxid)?

How are half cell potentials measured? A definition of zero potential is needed.

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E0cell = E°1/2(reduc) + E°1/2(oxid) =

What is E°1/2(reduc) ? What is E°1/2 (oxid)?

What is E°cell for the following reaction?

Al(s) + Cu+2(aq) R Al+3(aq) + Cu(s)

Use the table of half cell potentials to find the CELL POTENTIAL for any redox reaction.

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Br2 O2 Fe2+ Na+

Which one of these is the best oxidizing agent? Strategy: write out half cell potentials for REDUCTION

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Which is the best reducing agent of the following? Cl!(aq) Fe(s) Fe+3(aq) Fe+2(aq) Cl2(g)

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Which is the best oxidizing agent of the following? Cl!(aq) Fe(s) Fe+3(aq) Fe+2(aq) Cl2(g)

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What is the value of "G° for the reaction?

The standard cell potential (E°cell) for the reaction below is 0.89V.

2Cr(s) + 3 Sn4+ (aq) R 2 Cr3+ (aq) + 3 Sn2+ (aq)

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Corresponds to Standard conditions 1M solution or 1 atm gas pressure T = 298K

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2>(?9@#A( ,7&8%'98"&:;<%'=(6( QW(

Replace "G with !nF((((E and "G° with !nF((((E°

cAF(((,((t(cAF(((,q((^(/.(+)(x((((((

For electrochemical cell at equilibrium: "G = 0, E cell = 0

Nernst equation gives the concentration dependence of cell potential Ecell.(

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!

E = Eo "RTn#

lnQ

!

E 12

= E 12

0 " RTn#

ln C[ ]c D[ ]d

A[ ]a B[ ]b

!

E 12

= E 12

o " 2.303 RT

n#log

C[ ]c D[ ]d

A[ ]a B[ ]b

!

E 12

= E 12

o " 0.059

nlog

C[ ]c D[ ]d

A[ ] a B[ ]b at 298K

So for a half reaction: aA + bB + n e! R cC + dD

Sample problem using the Nernst equation for a half cell potential.(

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What is the half cell potential of the Ag/Ag+ redox couple (E0 = +0.799 V) in a 1 M NaCl solution that contains solid

AgCl (Ksp = 1.1 x 10!10)?

For an overall cell reaction: lL + mM R pP + qQ

!

Ecell = E cell0 "

RTn#

ln P[ ]p Q[ ]q

L[ ]l M[ ]m

!

Ecell = E cell0 "

2.303 RTn#

logP[ ] p Q[ ]q

L[ ]l M[ ]m

!

Ecell = E cell0 "

0.0592n

logP[ ]p Q[ ]q

L[ ]l M[ ]m

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Nernst equation gives the concentration dependence of cell potential Ecell.(

4#:$7&(b'9D7&:(using the Nernst equation (

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