Redox Reviewed

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Steps in Balancing Equations

The half-equation method separates the oxidation and reduction of a redox reaction in half reactions. Overall scheme for the half reaction method:

Step 1: Split reaction into half-reactions (reduction and oxidation) Step 2: Balance the charge or oxidation number with electrons Step 3: Balance O by adding H₂O

Step 4: Balance H by adding H+

Step 5: Multiply by some integer to make electrons (lost) = electrons (gained)

Step 6: Add half equations and cancel substances on both sides Step 7: (only in basic solution): add OH-_{and cancel H}

20

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Mg + HCl



MgCl

₂

+ H

₂

Mg is Oxidized and is the reducing agent H is Reduced and is the oxidation agent

Mg  Mg2+

+2e-H+ +e-  H

2 becomes 2H+ +2e- H2

---Mg + 2H+ +2e- _Mg2++2e- + H 2

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Fe + V

₂

O

₃



Fe

₂

O

₃

+ VO

Fe is Oxidized and is the reducing agent V is Reduced and is the oxidation agent

V₂O₃ VO 6e- + 6H++ 3V

2O3 6VO +3H2O Fe  Fe2O3

---2Fe + 3V₂O₃ Fe₂O₃+ 6VO

+ H₂O 2H+ ₊ ₂

2e- + (Multiple by 3)

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6e-KMnO₄ + KNO₂ + H₂SO₄  MnSO₄+ H₂O + KNO₃ + K₂SO₄

MnO

₄

- + NO

₂

-+ H+ Mn

2+

+ NO

3

-+ H

2

O

NO₂- is Oxidized and is the reducing agent

MnO₄- is Reduced and is the oxidation agent

MnO₄- Mn2+

NO₂- NO₃

-

---2MnO₄- + 5NO₂- 6H+5NO

3- 2Mn2++3H2O

+ 4H₂O 8H+ +

5e- +

H₂O + + 2H+ +

2e-Multiple by 2

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Electron Transfer Reactions

• _{Electron transfer reactions are}

oxidation-reduction

or

redox

reactions.

• _{Results in the generation of an electric}

current (electricity) or be caused by

imposing an electric current.

• _{Therefore, this field of chemistry is often}

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Review:Terminology for Redox

Reactions

Review:Terminology for Redox

Reactions

What is OXIDATION?

loss of electron(s) by a species; increase in oxidation number; increase in oxygen.

What is REDUCTION?

gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.

What is an OXIDIZING AGENT?

electron acceptor; species is reduced. What is a REDUCING AGENT?

electron donor; species

is oxidized.

What is OXIDATION?

loss of electron(s) by a species; increase in oxidation number; increase in oxygen.

What is REDUCTION?

gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.

What is an OXIDIZING AGENT?

electron acceptor; species is reduced.

What is a REDUCING AGENT?

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REMEMBER:

You can

’

t have one… without the other!

• _{Reduction (gaining electrons) can’t happen without an}

oxidation to provide the electrons.

• _{You can’t have 2 oxidations or 2 reductions in the same}

equation. Reduction has to occur at the cost of oxidation

LEO

the lion says

GER

!

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Again…. Another way to remember

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OXIDATION-REDUCTION

REACTIONS

OXIDATION-REDUCTION

REACTIONS

Direct Redox Reaction

Oxidizing and reducing agents in direct

contact.

Cu(s) + 2 Ag+(aq) --->

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OXIDATION-REDUCTION

REACTIONS

OXIDATION-REDUCTION

REACTIONS

Indirect Redox Reaction

A battery functions by transferring electrons through an external wire from the reducing

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Why Study Electrochemistry?

• _Batteries

• Corrosion

• Industrial production

of chemicals

such

as Cl

₂

, NaOH, F

₂

and Al

• Biological redox

reactions

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Electrochemical Cells

• _{An apparatus that allows a}

redox reaction to occur by

transferring electrons

through an external

connector.

• _{Product favored reaction}

--->

voltaic or galvanic cell

----> electric current

• _{Reactant favored reaction}

--->

electrolytic cell

--->

electric current used to

cause chemical change.

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Anode

Cathode

Basic Concepts

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CHEMICAL CHANGE --->

ELECTRIC CURRENT

CHEMICAL CHANGE --->

ELECTRIC CURRENT

With time, Cu plates out

onto Zn metal strip, and

Zn strip “disappears.”

With time, Cu plates out

onto Zn metal strip, and

Zn strip “disappears.”

• _{Zn is oxidized}

_{and is the reducing agent}

Zn(s) ---> Zn

2+

(aq) +

2e-•

_Cu

2+

is reduced

and is the oxidizing agent

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• _{To obtain a useful current,}

we separate the oxidizing

and reducing agents so that

electron transfer occurs thru

an external wire.

CHEMICAL CHANGE --->

ELECTRIC CURRENT

CHEMICAL CHANGE --->

ELECTRIC CURRENT

This is accomplished in a

GALVANIC

or

VOLTAIC

cell.

A group of such cells is called a

battery

.

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•Electrons travel thru external wire.

• _{Salt bridge}

_{allows anions and cations to move}

between electrode compartments.

•Electrons travel thru external wire.

• _{Salt bridge}

_{allows anions and cations to move} between electrode compartments.

Zn --> Zn2+_{+ 2e-} _Cu2+_{+ 2e- --> Cu}

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CELL POTENTIAL, E

• _{For Zn/Cu cell,}

_potential

_is

_{+1.10 V}

_{at 25 ˚C and}

when [Zn

2+

] and [Cu

2+

] = 1.0 M.

• _{This is the}

STANDARD CELL

POTENTIAL, E

o

• _{—a quantitative measure of the tendency of}

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Calculating Cell Voltage

• _{Balanced half-reactions can be added}

together to get overall, balanced

equation.

Zn(s) ---> Zn

2+

(aq) +

2e-Cu

2+

(aq) + 2e- ---> Cu(s)

---Cu

2+

(aq) + Zn(s) ---> Zn

2+

(aq) + Cu(s)

Zn(s) ---> Zn

2+

(aq) +

REDUCTION POTENTIALS

2 Eo (V)

Cu2+ + 2e- Cu

+0.34

2 H+ + 2e- H

0.00 Zn 2+ + 2e- Zn

-0.76

oxidizing

ability of ion

reducing ability

+

2e-or

Cd

2+

+ 2e- --> Cd

Fe --> Fe

2+

+

2e-or

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More About

Calculating Cell Voltage

Assume I

-

ion can reduce water.

2 H

₂

O + 2e-

--->

H

₂

+ 2 OH

-

Cathode

2 I

-

_--->

I

2

+ 2e-

Anode

E˚ = E˚

_cat

+ E˚

_an

= (-0.828 V) - (- +0.535 V) =

-1.363 V

Minus E˚ means rxn. occurs in opposite direction

(the connection is backwards or you are

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Charging a Battery

When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.

In your car, the battery charger is called an alternator. If you have a dead battery, it

could be the battery needs to be replaced OR the alternator is not charging the battery

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Dry Cell Battery

Anode (-)

Zn ---> Zn

2+

+

2e-Cathode (+)

2 NH

₄+

+ 2e- --->

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Alkaline Battery

Nearly same reactions as in common dry cell, but

under basic conditions.

Anode (-):

Zn + 2 OH

-

---> ZnO + H

2

O +

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-Mercury Battery

Anode:

Zn is reducing agent under basic conditions

Cathode:

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-Lead Storage Battery

Anode (-) Eo = +0.36 V

Pb + HSO₄- ---> PbSO

4 + H+ +

2e-Cathode (+) Eo = +1.68 V

PbO₂ + HSO₄- + 3 H+ + 2e-

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Ni-Cad Battery

Anode (-)

Cd + 2 OH- ---> Cd(OH)

2 +

2e-Cathode (+)

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-H

₂

as a Fuel

Cars can use electricity generated by H

₂

/O

₂

fuel cells.

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Balancing Equations

for Redox Reactions

Some redox reactions have equations that must be balanced by special techniques.

MnO₄- + 5 Fe2+ + 8 H+ ---> Mn2+ + 5

Fe3+ + 4 H 2O

Mn = +7 Fe = +2

Fe = +3

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Balancing Equations

Consider the reduction of Ag+ ions with

copper metal.

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Balancing Equations

Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction.

Ox Cu ---> Cu2+

Red Ag+ ---> Ag

Step 2: Balance each element for mass. Already done in this case.

Step 3: Balance each half-reaction for charge by adding electrons.

Ox Cu ---> Cu2+ +

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Balancing Equations

Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the

oxidizing agent requires.

Reducing agent Cu ---> Cu2+ +

2e-Oxidizing agent 2 Ag+ + 2 e- ---> 2 Ag

Step 5: Add half-reactions to give the overall equation. Cu + 2 Ag+ ---> Cu2+ + 2Ag