Chapter 21. the study of the production of electricity during chemical rxns and the changes produced by electrical current. Electrochemical reactions.

ElectrochemistryChapter 21

Electrochemistrythe study of the production of electricity during

chemical rxns and the changes produced by electrical current.

Electrochemical reactions are oxidation-reduction reactions.

Oilrig: oxidation loss of electrons, reduction gaining of electrons

1. Oxidation = loss of electronsa. the substance oxidized is the reducing agent

2. Reduction = gain of electronsa. the substance reduced is the oxidizing agent

3

Electrical Conduction Metals conduct electric currents well.

metallic conduction

Positively charged ions, cations, move toward the negative electrode.

Negatively charged ions, anions, move toward the positive electrode.

There are two kinds electrochemical cells.

Electrolytic cells - nonspontaneous chemical reactions

Voltaic or galvanic cells - spontaneous chemical reactions

The two parts of the reaction are physically separated.

–oxidation occurs at one cell–reduction occurs in the other cell

Conventions for electrodes:

Cathode - electrode at which reduction occurs (red cat)

Anode - electrode at which oxidation occurs (an ox)

Inert electrodes do not react with the liquids or products of the electrochemical reaction. Graphite and Platinum are common inert electrodes.

Electrolytic CellsUse electrical energy to force

nonspontaneous (non thermodynamically favored) chemical reactions to occur.

Process called electrolysis. Used in:–plating of jewelry and auto parts–electrolysis of chemical compounds

Electrolytic cells consist of a:–container for reaction mixture–electrodes immersed in the reaction mixture–source of direct current

In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode).

In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

Faraday’s Law of Electrolysis

The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell. During electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent.

Corresponds to the passage of one mole of electrons through the electrolytic cell.

FaradayAmount of electricity that reduces one

equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.

1 faraday of electricity = 6.022x1023 e-

1 faraday = 6.022x1023 e- = 96487 coulombs

1 eq. of oxidizing agent= gain of 6.022x1023 e-

1 eq. of reducing agent = loss of 6.022x1023 e-

Ex. 1) Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes. (1 ampere = 1 coulomb per second)

11

Ex. 2) Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in Ex. 1.

Voltaic or Galvanic Cells •Electrochemical cells in which a spontaneous

chemical reaction produces electrical energy.

•In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

•Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference Examples: Car & flashlight batteries

13

The Construction of Simple Voltaic Cells

Half-cell contains the oxidized and reduced forms of an element (or other chemical species) in contact with each other.

Simple cells consist of: two pieces of metal immersed in solutions of their

ions wire to connect the two half-cells salt bridge to

▪ complete circuit▪ maintain neutrality▪ prevent solution mixing

14

The Zinc-Copper Cell

Cell components:Cu strip immersed in 1.0 M copper (II) sulfateZn strip immersed in 1.0 M zinc (II) sulfatewire and a salt bridge to complete circuit

Initial voltage is 1.10 volts

15

The Zinc-Copper Cell

16

The Zinc-Copper Cell

Anode Zn Zn e

Cathode Cu e Cu

Cell rxn. Zn Cu Zn Cu

Spontaneous rxn. E V

0 2+

2+ 0

0 2+ 2+ 0

cell0

2

2

110.

In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

Zn/Zn2+(1.0 M) || Cu2+(1.0 M)/Cu

species (and concentrations)in contact with electrode surfaces

electrode surfaces

salt bridge

The Zinc-Copper Cell

Short hand notation for voltaic cellsZn-Cu cell example

18

The Copper - Silver Cell

Cell components:Cu strip immersed in 1.0 M copper (II)

sulfateAg strip immersed in 1.0 M silver (I)

nitratewire and a salt bridge to complete circuit

Initial voltage is 0.46 volts

19

The Copper - Silver Cell

20

The Copper - Silver Cell

Anode Cu Cu 2e

Cathode 2 Ag + e Ag

Cell rxn. Cu + 2 Ag Cu + 2 Ag

spontaneous rxn. E V

0 2

+ - 0

0 + 2+ 0

cell0

0 46. Compare the Zn-Cu cell to the Cu-Ag cell

Cu electrode is cathode in Zn-Cu cellCu electrode is anode in Cu-Ag cell

Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.

21

The Copper - Silver Cell

Demonstrates that Cu2+ is a stronger oxidizing agent than Zn2+

Cu2+ oxidizes metallic Zn to Zn2+

Ag+ is is a stronger oxidizing agent than Cu2+

Ag+ oxidizes metallic Cu to Cu2+

Arrange these species in order of increasing strengths

Zn < Cu < Ag Zn > Cu > Ag

strength as oxidizing agent strength as reducing agent

2+ 2+ + 0 0 0

22

Standard Electrode Potential Establish an arbitrary standard to

measure potentials of a variety of electrodes

Standard Hydrogen Electrode (SHE) assigned an arbitrary voltage of 0.000000…

V

23

The Electromotive (Activity) Series of the Elements

Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials.

Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials.

Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.

24

Uses of the Electromotive Series

Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.

For example, the half-reaction for the standard potassium electrode is:

The large negative value tells us that this reaction will occur only under extreme conditions.

K e K E = -2.925 V0 0

25

Uses of the Electromotive Series

Compare the potassium half-reaction to fluorine’s half-reaction:

The large positive value denotes that this reaction occurs readily as written.

Positive E0 values tell us that the reaction tends to occur to the right larger the value, greater tendency to occur

to the right Opposite for negative values

F + 2 e 2 F E = +2.87 V20 - - 0

Uses of the Electromotive Series

1. Choose the appropriate half-reactions from a table of standard reduction potentials.

2. Write the equation for the half-reaction with the more positive E0 value first, along with its E0 value.

3. Write the eqn for the other half-reaction as an oxidation with its oxidation potential, reverse the tabulated reduction half-reaction and change the sign of the tabulated E0.

4. Balance the electron transfer.

5. Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous.

27

Uses of the Electromotive Series

Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously.

Ex. 3) Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn2+ ions oxidize metallic Ag to Ag+ ions? What is the overall value for Eo ?

E0 values are not multiplied by any stoichiometric relationships in this procedure.

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Electrode Potentials for Other Half-Reactions

Ex.4) Will tin(IV) ions oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize tin(II) ions to tin(IV) ions in acidic solution? What is the overall value for rxn?

29

Effect of Conc. (or Partial Pressures) on Electrode Potentials - Nernst Equation

Standard electrode potentials are determined at thermodynamic standard conditions.1 M solutions1 atm of pressure for gasesliquids and solids in their standard statestemperature of 250 C

Potentials change if conditions are nonstandard.

Nernst equation describes the electrode potentials at nonstandard conditions.

30

Effect of Conc. (or Partial Pressures) on Electrode Potentials - Nernst Equation

E = E -2.303RT

nF log Q

E = potential under condition of interest

E potential under standard conditions

R = universal gas constant = 8.314 Jmol K

T = temperature in K

n = number of electrons transferred

F = the faraday = (96,487 C / mol e 1 JC V

J / V mol e

Q = reaction quotient

0

0

-.

-

)

,96 487

31

2.303 RTF

so

E = E -0.0592

n log Q0

0 0592.

Substitution of the values of the constants into the Nernst equation at 250 C gives:

The Nernst Equation used to calculate electrode potentials

and cell potentials for concentrations other than standard –state values. (pg 877)

Ex. 5) Calculate the cell potential for the following electrochemical cell. If the [Sn2+] = 4.5 x 10-1 M and [Ag+] = 0.110 M

Sn(s) + 2Ag+ Sn2+ + 2Ag(s)

33

The Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes.

Ex. 6) Calculate the initial potential of a cell that consists of an Fe3+/Fe2+ electrode in which [Fe3+]=1.0 x 10-2 M and [Fe2+]=0.1 M connected to a Sn4+/Sn2+ electrode in which [Sn4+]=1.0 M and [Sn2+]=0.10 M . A wire and salt bridge complete the circuit.

34

Relationship of E0cell to DG0

and K From previous chapters we know the

relationship of DG0 and K for a reaction.

K log RT 303.2Gor lnK -RTG 00

35

Relationship of E0cell to DG0

and K The relationship between DG0 and E0

cell is also a simple one.

e ofnumber n

e mol J/V 96,487 F where

E F-n G

-

-

0cell

0

36

Relationship of E0cell to DG0

and K You can combine these two

relationships into a single relationship to relate E0

cell to K.

RT

E Fn K ln

or

lnK RT -E Fn -

0cell

0cell

37

Relationship of E0cell to DG0

and K Ex. 7) Calculate the standard Gibbs free

energy change, DG0 , at 250C for the following reaction.

Cu PbCuPb 22

38

Relationship of E0cell to DG0

and K Calculate E0

cell using the appropriate half-reactions.

39

Relationship of E0cell to DG0

and K Now that we know E0

cell , we can calculate DG0 . The negative value tells us that the reaction is

spontaneous as written.

40

Relationship of E0cell to DG0

and K Ex. 8) Calculate the thermodynamic

equilibrium constant for the reaction in Ex. 7 at 250C.

41

Corrosion

Metallic corrosion is the oxidation-reduction reactions of a metal with atmospheric components such as CO2, O2, and H2O.

4 Fe + 3 O 2 Fe O overall rxn.

Rxn. occurs rapidly at exposed points.

020

2 3

Corrosion = Oxidation of a metal

The oxidation of most metals by oxygen is spontaneous. Many metals develop a thin coating of metal oxide on the outside that prevents further oxidation

The presence of a salt accelerates the corrosion process by increasing the ease with which electrons are conducted from anodic to cathodic regions

43

Corrosion Protection Some examples of corrosion protection.

1 Plate a metal with a thin layer of a less active (less easily oxidized) metal.

2. Connect the metal to a sacrificial anode, a piece of a more active metal.

"Tin plate" for steel.

Soil pipes and ship hulls

Mg & Zn are sacrificial anodes

44

Corrosion Protection

3 Allow a protective film to form naturally.

4 Al + 3 O 2 Al O

hard, transparent film

020

2 3

45

Corrosion Protection4 Galvanizing, coating steel with zinc, a

more active metal.

5. aint or coat with a polymeric material such as plastic or ceramic.

thin coat of Zn which must

oxidize before Fe begins to rust

Steel bathtubs are coated with ceramic.

46

Primary Voltaic Cells

As a voltaic cell discharges, its chemicals are consumed.

Once chemicals are consumed, further chemical action is impossible.

Electrodes and electrolytes cannot be regenerated by reversing current flow through cell.

47

The Dry Cell (LeClanche’ Cell)

One example is flashlight, radio, etc. batteries. Container is made of zinc

acts as an electrode Graphite rod is in center of cell

acts as the other electrode Space between electrodes is filled with a

mixture of: ammonium chloride, NH4Cl

manganese (IV) oxide, MnO2

zinc chloride, ZnCl2 porous inactive solid

48

The Dry Cell (LeClanche’ Cell) As current is produced, Zn dissolves

and goes into solution as Zn2+ ions. Zn electrode is negative (anode).

49

Secondary Voltaic Cells

Secondary cells are reversible, rechargeable.

Electrodes can be regenerated One example is the lead storage or car

battery.

50

Lead Storage Battery Electrodes are two sets of lead alloy grids (plates). Holes in one grid are filled with lead (IV) oxide,

PbO2. Other holes are filled with spongy lead. Electrolyte is dilute sulfuric acid. When battery is discharging, spongy lead is

oxidized to lead ions and the plate becomes negatively charged.

Pb Pb 2 e

anode - negative during discharge

s0 2+ -

51

Lead Storage Battery

Cell reaction for a discharging lead storage battery is

Anode Pb SO PbSO 2 e

Cathode PbO + 4 H + SO + 2 e PbSO + 2 H O

Cell rxn Pb PbO + 4 H + SO 2 PbSO + 2 H O

s 42-

4 s-

2 s+

42- -

4 s 2

s 2 s+

42-

4 s 2

52

Lead Storage Battery This battery can be recharged by passing

electricity through the cell in the opposite direction.

Voltaic cell is converted to galvanic cell.

At lead electrode, lead ions are reduced to lead atoms.

PbSO 2 e Pb SO

cathode - negative during charging

4 s-

s 42-

53

The Nickel-Cadmium (Nicad) Cell

Rechargeable cell used in calculators, cameras, watches, etc.

Discharge half-reactions are:

V4.1E

Ni(OH)+Cd(OH) OH 2 + NiO Cdrxn Cell

OH 2+Ni(OH)e 2+OH 2+NiO Cathode

e 2 Cd(OH) OH 2 Cd Anode

0

s2s22s2s

-s2

-2s2

-2

-s

54

The Hydrogen-Oxygen Fuel Cell

Fuel cells are batteries that must have their reactants continuously supplied in the presence of appropriate catalysts.

Hydrogen-oxygen fuel cell is used in the space shuttle. Hydrogen is oxidized at anode Oxygen is reduced at cathode

55

The Hydrogen-Oxygen Fuel Cell

The reaction combines hydrogen and oxygen to form water. Drinking water supply for astronauts.

Very efficient energy conversion 60-70%

l2g2g2

--2g2

-2

-g2

OH 2 O H 2 rxn Cell

OH 4e 4OH 2O1 Cathode

e 2OH 2OH 2 H2 Anode

56

Bonus Questions

1. What are the explosive chemicals in the fuel cell that were aboard Apollo 13? Which tank exploded?

2. Some of the deadliest snakes in the world, for example the cobra, have venoms that are neurotoxins. Neurotoxins have an electrochemical basis. How do neurotoxins disrupt normal chemistry and eventually kill their prey?