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Chapter 17: Electrochemistry. Review of Redox Reactions Galvanic Cells: Using spontaneous redox reactions to generate electrical energy. Galvanic Cells Cell potential D G and work Cell potential and concentration Applications: Batteries, fuel cells, corrosion

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Chapter 17: Electrochemistry

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Chapter 17: Electrochemistry

  • Review of Redox Reactions

  • Galvanic Cells: Using spontaneous redox reactions to generate electrical energy.

    • Galvanic Cells

    • Cell potential

    • DG and work

    • Cell potential and concentration

    • Applications: Batteries, fuel cells, corrosion

  • Electrolytic Cells: Using electricity to cause nonspontaneous redox reactions to occur.

    • Electrolytic Cells

    • Applications: Electrolysis of water, electrolysis of mixtures of ions, production of Al, electrorefining, metal plating, electrolysis of NaCl


Example 1

Balance the following redox reactions

using the half reaction method.

  • ClO3- + As2S3 Cl- + H2AsO4- + SO42-

  • K2S + KMnO4 S8 + MnO2 + KOH


Figure 17.6: Cartoon of atoms reacting


Figure 17.5: A standard hydrogen electrode


Example 2

If a standard Cu|Cu2+ electrode is

connected to a standard Al|Al3+

electrode, what reaction occurs? What

is °cell? Sketch the cell, label the

cathode and anode, show the direction

of electron flow, and give the line

notation for the cell.


Example 3

Consider the following species under

standard conditions:

Ce4+, Ce3+, Fe2+, Fe3+, Fe, Mg, Mg2+, Ni2+, Sn

  • Which is the strongest oxidizing agent?

  • Which is the strongest reducing agent?

  • Will Fe dissolve in 1.0 M Ce4+? If so, will Fe3+ or Fe2+ be formed?

  • Which can be oxidized by H+(aq)?

  • Which can be reduced by H2(g)?


Example 4

Select an oxidizing agent to oxidize Cl-

to Cl2 without oxidizing Br- to Br2.


Example 5

Calculate the maximum work available

from 25.0 g of aluminum in the

following galvanic cell for which the emf

is 1.15 V. Note that O2 is reduced to

H2O in this reaction.

Al(s)|Al3+(aq)||H+(aq)|O2(g)|Pt(s)


Example 6

Calculate the cell potential for the

following Galvanic cell at 25°C.

Ni(s)|Ni2+(1.0M)||Sn2+(1.0x10-4M)|Sn(s)


Example 7

Find the potential of a

Ag+(1.0x10-7M)|Ag(s) electrode at 25°C.


Example 8

Calculate the equilibrium constant for

the following reaction at 25°C.

Ag+(aq) + Fe2+(aq)  Ag(s) + Fe3+(aq)


Example 9

A concentration cell is made up of two

Ag/Ag+ half cells. In the first half cell,

[Ag+] = 0.010 M. In the second half

cell, [Ag+] = 4.0 x 10-4 M. What is the

cell potential? Which half cell functions

as the anode?


Figure 17.13: Lead storage battery


Figure 17.14: Common dry cell battery

KOH


Figure 17.16: hydrogen-oxygen fuel cell


Figure 17.17: Corrosion of Iron


Figure 17.18: Cathodic protection of an underground pipe


Example 10

Predict the products and calculate the

minimum voltage required for the

electrolysis of the following substances

using platinum electrodes.

  • MgBr2(l)

  • 1.0 M NiCl2(aq)


Example 11

Write the net ionic equation for the

reaction you expect to occur when the

electrolysis of NiSO4(aq) is conducted

using a nickel anode and an iron

cathode.


Example 12

How many grams of silver are deposited

at a platinum cathode in the

electrolysis of an aqueous solution of

AgNO3 by 1.73A of electric current in

2.5 hours?


Example 13

How long will it take to produce 10.0 g

of bismuth (Bi) by the electrolysis of a

BiO+ solution using a current of 25.0 A?


Figure 17.22: Production of Al


Figure 17.23: Refining of Copper


Figure 17.25: Downs cell for production of sodium and chlorine


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