Unit 12 Electrochemistry - PowerPoint Presentation

Slide1

Unit 12ElectrochemistryDay 2

Lecture Presentation

Slide2

Warm Up

SHOW ME: Cornell Notes, Video Notes, and Practice ProblemsFIND: Balancing Redox Reactions Worksheet 1 at desk and take out notebook paper to work on TIME: 4 minutesWHEN DONE: Get calculator from blue bin

Slide3

Agenda

Review Balancing Redox

Rxn

Practice (12 min)

Voltaic Cells (12 min)

Cell Potential (9 min)

Multiple Choice Questions

Pass Back Work and Work Time:

ChemActivities

and Set Up Lab

Slide4

Voltaic Cells

In spontaneous redox reactions, electrons are transferred and energy is released.That energy can do work if the electrons flow through an external device.This is a voltaic cell.

Slide5

Voltaic Cells

The oxidation occurs at the anode.The reduction occurs at the cathode.When electrons flow, charges aren’t balanced. So, a salt bridge, usually a U-shaped tube that contains a salt/agar solution, is used to keep the charges balanced.

Slide6

Voltaic Cells

In the cell, electrons leave the anode and flow through the wire to the cathode.Cations are formed in the anode compartment.As the electrons reach the cathode, cations in solution are attracted to the now negative cathode.The cations gain electrons and are deposited as metal on the cathode.

Slide7

Electromotive Force (emf)

Water flows spontaneously one way in a waterfall.Comparably, electrons flow spontaneously one way in a redox reaction, from high to low potential energy.

Slide8

Electromotive Force (emf)

The potential difference between the anode and cathode in a cell is called the

electromotive force (emf)

.

It is also called the

cell potential

and is designated

E

cell

.

It is measured in volts (V). One volt is one joule per coulomb (1 V = 1 J/C).

Slide9

Standard Reduction Potentials

Reduction potentials for many electrodes have been measured and tabulated.The values are compared to the reduction of hydrogen as a standard.

Slide10

Standard Hydrogen Electrode

Their reference is called the standard hydrogen electrode (SHE).By definition as the standard, the reduction potential for hydrogen is 0 V:

2 H+(aq, 1M) + 2e–  H2(g, 1 atm)

Slide11

Standard Cell Potentials

The cell potential at standard conditions can be found through this equation:

E

cell

°

=

Ered (cathode) – Ered (anode)

°

°

Because cell potential is based on the potential energy per unit of charge, it is an intensive property.

Slide12

Cell Potentials

For the anode in this cell, E°red = –0.76 VFor the cathode, E°red = +0.34 VSo, for the cell, E°cell = E°red (anode) – E°red (cathode) = +0.34 V – (–0.76 V) = +1.10 V

Slide13

Oxidizing and Reducing Agents

The more positive the value of E°red, the greater the tendency for reduction under standard conditions.The strongest oxidizers have the most positive reduction potentials.The strongest reducers have the most negative reduction potentials.

Slide14

Free Energy and Redox

Spontaneous redox reactions produce a positive cell potential, or emf.

E

°

=

E

°

red

(reduction) –

E

°

red

(oxidation)

Note that this is true for ALL redox reactions, not only for voltaic cells.

Since Gibbs free energy is the measure of spontaneity, positive emf corresponds to

negative

Δ

G

.

How do they relate?

Δ

G

= –

nFE

(

F

is the Faraday constant, 96,485 C/mol.)

Slide15

Free Energy, Redox, and K

How is everything related?ΔG° = –nFE° = –RT ln K

Slide16

Nernst Equation

Remember,

Δ

G

=

Δ

G

°

+

RT

ln

Q

So,

–

nFE

=

nFE

°

+

RT

ln

Q

Dividing both sides by

–

nF

, we get the

Nernst equation

:

E

=

E

°

– (

RT

/

nF

) ln

Q

OR

E

=

E

°

– (2.303

RT

/

nF

) log

Q

Using standard thermodynamic temperature and the constants

R

and

F

,

E

=

E

°

– (0.0592/n) log

Q

Slide17

Concentration Cells

Notice that the Nernst equation implies that a cell could be created that has the same substance at both electrodes, called a concentration cell.

For such a cell,

would be 0, but

Q would not.

E

cell

°

Therefore, as long as the concentrations are different, E will not be 0.

Slide18

Some Applications of Cells

Electrochemistry can be applied as follows:Batteries: a portable, self-contained electrochemical power source that consists of one or more voltaic cells.Batteries can be primary cells (cannot be recharged when “dead”—the reaction is complete) or secondary cells (can be recharged).Prevention of corrosion (“rust-proofing”)Electrolysis

Slide19

Some Examples of Batteries

Lead–acid battery: reactants and products are solids, so Q is 1 and the potential is independent of concentrations; however, made with lead and sulfuric acid (hazards).Alkaline battery: most common primary battery.Ni–Cd and Ni–metal hydride batteries: lightweight, rechargeable; Cd is toxic and heavy, so hydrides are replacing it.Lithium-ion batteries: rechargeable, light; produce more voltage than Ni-based batteries.

Slide20

Some Batteries

Lead–Acid Battery

Alkaline Battery

Slide21

Lithium-Ion Battery

Slide22

Fuel Cells

When a fuel is burned, the energy created can be converted to electrical energy.

Usually, this conversion is only 40% efficient, with the remainder lost as heat.

The direct conversion of chemical to electrical energy is expected to be more efficient and is the basis for

fuel cells

.

Fuel cells are NOT batteries; the source of energy must be continuously provided.

Slide23

Hydrogen Fuel Cells

In this cell, hydrogen and oxygen form water.The cells are twice as efficient as combustion.The cells use hydrogen gas as the fuel and oxygen from the air.

Slide24

Corrosion

Corrosion is oxidation.Its common name is rusting.

Slide25

Preventing Corrosion

Corrosion is prevented by coating iron with a metal that is more readily oxidized.Cathodic protection occurs when zinc is more easily oxidized, so that metal is sacrificed to keep the iron from rusting.

Slide26

Preventing Corrosion

Another method to prevent corrosion is used for underground pipes.A sacrificial anode is attached to the pipe. The anode is oxidized before the pipe.

Slide27

Electrolysis

Nonspontaneous reactions can occur in electrochemistry IF outside electricity is used to drive the reaction.Use of electrical energy to create chemical reactions is called electrolysis.

Slide28

Electrolysis and “Stoichiometry”

1 coulomb = 1 ampere × 1 secondQ = It = nFQ = charge (C)I = current (A)t = time (s)n = moles of electrons thattravel through the wire inthe given timeF = Faraday’s constantNOTE: n is different than thatfor the Nernst equation!

Slide29

The oxidation state of nitrogen in the ammonium ion (NH41+) is _______.

+1

0

−

1

−

3

Slide30

The oxidation state of manganese in the permanganate ion (MnO41−) is _______.

a

.

−

1

+2

+4

d

. +

7

Slide31

Zn + Cu2+  Zn2+ + CuThe reducing agent in the reaction above is _______.

Zn

Cu

2+

Zn

2+

d. Cu

Slide32

Zn + Cu2+  Zn2+ + CuThe oxidizing agent in the reaction above is _______.

Zn

Cu

2+

Zn

2+

d. Cu

Slide33

To balance the half-reaction MnO4−  Mn2+ in acidic solution, ___ electrons must be added on the ___ side.

5; product

2; product

7; reactant

5; reactant

Slide34

MnO4−  Mn2+ To balance this reaction in acidic solution, ___ H+ must be added on the ___ side.

8; product

4; product

8; reactant

4; reactant

Slide35

If the value of the standard cell potential for a reaction is large and positive, then the reaction is

at equilibrium.

spontaneous.

nonspontaneous

.

d

. very

fast.

Slide36

The purpose of the salt bridge in a voltaic cell is to

provide H

+

ions needed to balance charges.

maintain neutrality by allowing

the flow

of ions.

serve as the site for oxidation

to

occur.

serve as the site for

reduction

to

occur.

Slide37

Which of the following is the strongest oxidizing agent? (If necessary, consult a table of standard reduction potentials.)

F

2

Cl

2

Br

2

d. I

2

Slide38

Which of the following is the strongest reducing agent? (If necessary, consult a table of standard reduction potentials.)

Zn

Al

Na

d. Li

Slide39

Cl2  2 Cl− εred = + 1.36 VI2  2 I− εred = + 0.54 VSelect the true statement.

Cl

2

will reduce

I

−

to

I

2

.

Cl

2

will oxidize

I

−

to

I

2

.

I

2

will reduce

Cl

−

to Cl

2

.

I

2

will oxidize Cl

2

to

Cl

−

.

Slide40

G = Gibbs free energy. n = moles of electrons. F = Faraday constant. E = cell potential. Change in G = ?

F +

nE

F

−

nE

c.

−

nFE

d

.

nF

−

E

Slide41

The Nernst Equation is most useful for determining cell potentials when _______ are nonstandard.

oxidizing agents

reducing agents

ion concentrations

d. temperatures

Slide42

In a concentration cell, the half-reactions are

the same.

at equilibrium.

acid–base

reactions.

d

. different

colors.

Slide43

The electrolyte in a lead-acid automobile battery is _______.

Pb

PbO

2

PbSO

4

d. H

2

SO

4

Slide44

A cell that uses external energy to produce an oxidation–reduction reaction is called _______ cell.

a galvanic

a voltaic

an electrolytic

d. a

prison

Slide45

Reduction occurs at the

anode, in both galvanic and electrolytic cells.

cathode, in both galvanic and electrolytic cells.

anode in galvanic cells and cathode in electrolytic cells.

cathode in galvanic cells and anode in electrolytic cells.

Slide46

Oxidation occurs at the

anode, in both galvanic and electrolytic cells.

cathode, in both galvanic and electrolytic cells.

anode in galvanic cells and cathode in electrolytic cells.

d

. cathode

in galvanic cells and anode in electrolytic cells.

Slide47

Q = charge in coulombsI = current in amperest = time in secondsWhich is true?

Q = I + t

Q = I

−

t

Q = It

d

. Q

= I / t

Slide48

Corrosion of metals can be prevented by all of the following methods except

a sacrificial anode.

a salt bridge.

formation of an oxide coating.

paint.