Link to Prentice Hall Text: Chapters 22 & 23

Unit 13: Electrochemistry (Link to Prentice Hall Text: Chapters 22 & 23)

Name:______

Date Due Assignments Page Number: Problem Numbers

Assignment 1: Balancing and Identifying REDOX Reactions 673: 30, 32, 34

Assignment 2: Electrochemical Cells 701: 22, 34, 35, 38, 42

A. Atoms Compete for Electrons (Electrochemistry)

Defining Oxidation and Reduction

When chemical bonds form, electrons are either lost, gained or shared. In REDOX reactions, electrons are lost or gained.

Oxidation Is Loss, Reduction Is Gain, “OIL RIG”

Oxidation:

- Loss of electrons.

- Metals are more easily oxidized.

- They are termed “reducing agents.”

- If a substance is oxidized, its oxidation number increases.

Reduction:

- Gain of electrons.

- Non-metals are more easily reduced.

- They are termed “oxidizing agents.”

- If a substance is reduced, its oxidation number decreases. B. Oxidation Numbers

Rules for Assigning Oxidation Numbers:

(1) Oxidation numbers for atoms that are free elements are always ______.

(2) The oxidation number of monatomic ions are the same as the charge on the ______.

(3) The sum of the oxidation numbers in a compound is always ______.

(4) The sum of the oxidation numbers in a polyatomic ion is equal to ______.

(5) The oxidation number of Group 1 metals is always ______.

(6) The oxidation number of Group 2 metals is always ______.

(7) Oxygen almost always has a ______oxidation state.* Unless in a peroxide

(8) Halogens usually have a ______oxidation state.*

(9) Hydrogen usually has a ______oxidation state.*

* There are rare exceptions to these rules.

Practice

Assign oxidation numbers to each element in the following compounds or polyatomic ions. To really keep on your toes, see if you can name them as well.

1. MgBr2 9. CuSO4 17. K2Cr2O7

2. Cu 10. Cr 18. Al2O3

3. Fe2O3 11. H2CO3 19. Fe(NO3)3

4. AlN 12. Ba(NO3)2 20. SrCO3

5. SO3 13. NF3 21. Na2SO3

3- 6. PO4 14. CO 22. Ca(ClO)2

2- 7. Cr2O7 15. CO2 23. H2O

8. HClO2 16. CH4 24. K2S C. Identifying REDOX Reactions

Identifying REDOX Reactions Using Oxidation Numbers

When a REDOX reaction occurs, there must be an element that is reduced (gained electrons) and an element that is oxidized (lost electrons).

Are the following reactions REDOX reactions?

HINT: Find oxidation numbers of each element.

(1) 2H2 + O2  2H2O

(2) HCl + NH3  NH4Cl

Oxidation Is Loss, Reduction Is Gain, “OIL RIG”

Identifying Oxidation and Reduction Half Reactions

Oxidation and reduction occur simultaneously. Elements and charges must be balanced for each half-reaction.

1. Zn + CuCl2  Cu + ZnCl2

Oxidation Half:

Reduction Half:

2. Na + Cl2  NaCl

Oxidation Half:

Reduction Half:

3. Ca + O2  CaO

Oxidation Half:

Reduction Half:

4. MnO2 + HCl  MnCl2 + H2O + Cl2

Oxidation Half:

Reduction Half:

+ 2+ 5. Cu + Ag  Ag + Cu

Oxidation Half:

Reduction Half: D. Balancing REDOX Reactions Using Half-Reactions

Steps to Balancing REDOX Reactions

When a REDOX reaction occurs, there must be an element that is reduced (gained electrons) and an element that is oxidized (lost electrons). The number of electrons lost in the oxidation must equal the number of electrons gained in the reduction.

(1) Assign oxidation numbers. (2) Decide what is being oxidized and what is being reduced. (3) Write the half reaction for each oxidation and reduction. (4) Balance all atoms (5) Balance the charge with electrons in each half reaction. (6) Balance the number of electrons gained with the number of electrons lost. (7) Add the two half reactions. (8) Simplify each equation by canceling out things that are the same on both sides.

Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons), “OIL RIG”

Oxidation Step: Oxidation number of element increases. Reduction Step: Oxidation number of element decreases.

Hints: 1. There should be no electrons in your final equation. 2. The number of each element should be the same on both sides of the arrow when you are finished.

Balancing REDOX Reactions

Oxidation and reduction occur simultaneously. Elements and charges must be balanced for each half-reaction. Electrons must be the same on either sides of the arrow.

1. Zn + CuCl2  Cu + ZnCl2 Oxidizing Agent Reducing Agent

Oxidation Half:

Reduction Half:

Balanced Reaction:

4+ 2+ 2. Ni + Sn  Ni + Sn Oxidizing Agent Reducing Agent

Oxidation Half:

Reduction Half:

Balanced Reaction: 3. Ni + Sn+4  Ni+2 + Sn Oxidizing Agent Reducing Agent Oxidation Half:

Reduction Half:

Balanced Reaction:

4. Hg + Ag+1  Ag + Hg+2 Oxidizing Agent Reducing Agent Oxidation Half:

Reduction Half:

Balanced Reaction:

5. KClO3  KCl + O2 Oxidizing Agent Reducing Agent Oxidation Half:

Reduction Half:

Balanced Reaction:

6. H2S + O2  SO2 + H2O Oxidizing Agent Reducing Agent Oxidation Half:

Reduction Half:

Balanced Reaction:

+ +2 7. H + Sn  Sn + H2 Oxidizing Agent Reducing Agent Oxidation Half:

Reduction Half:

Balanced Reaction: E l e c t r o c h e m i s t r y P a g e | 6

E. Electrochemical Cells

The Basics of an Electrochemical Cell

Every electrochemical cell (contains two compartments where REDOX reactions can occur) contains 2 electrodes. An electrode is a strip of metal that facilitates the loss or gain of electrons. There are two types of electrodes, known as the cathode and the anode.

Anode Cathode

Where ______occurs. Where ______occurs.

Electrons always flow from the ______to the ______.

If you have 2 electrodes made of different metals connected, how can you tell which one is oxidized and which one is reduced? Table J!

Oxidation If the metal is ______on Table J. (More active = easier to lose electrons)

Reduction If the metal is ______on Table J. (Less active = harder to lose electrons)

Voltaic or Galvanic Electrochemical Cells

In a Galvanic/Voltaic cell, reactions occur ______; therefore, no additional energy must be added.

Spontaneous Reaction - Reactions that proceed without the addition of energy. - Predicting Spontaneous Reactions (Table J) - If the metal by itself if higher up on the chart than the ion in the compound. E l e c t r o c h e m i s t r y P a g e | 7

- In a voltaic cell, the anode is denoted with a NEGATIVE sign. - In a voltaic cell, the cathode is denoted with a POSTIVE sign.

- Each half cell is connected by a salt bridge. o What is a salt bridge?

- Negative ions always flow from the ______to the ______.

- Positive ions always flow from the ______to the ______.

- The electrodes are connected by a metal wire which allows for the flow/movement of electrons. Remember: Electrons always flow from the anode to the cathode.

- A voltmeter usually measures the electron potential in Volts. E l e c t r o c h e m i s t r y P a g e | 8

Determining Spontaneity

Which of the following reactions will take place spontaneously?

a. Ni (SO4) + Pb  Ni + Pb(SO4)c. Au (PO4) + Al  Au + Al (PO4) b. Sr(CO3) + Sn  Sr + Sn(CO3) d. Fe(OH) + Cu  Fe + Cu(OH)

Labeling Electrochemical Cell Processes

Between each of the following electrodes:

Species Oxidized Species Reduced Anode Metal Cathode Metal Direction of e- Direction of With Sign With Sign Flow Cation Flow in the Salt Bridge Calcium and Iron

Silver and Nickel

Magnesium and Lead

Copper and Silver

Electrolytic Cells E l e c t r o c h e m i s t r y P a g e | 9

Reactions that occur in an electrolytic cell occur ______. You must force the reaction to occur by adding electricity.

In Electrolytic Cells: - The anode is POSITIVE. - The cathode is NEGATIVE.

______

Why would you ever want to force a nonspontaneous reaction?

(a) To obtain pure metals Many metals are only found as compounds in nature. Electrolysis can lead to a deposit of the pure metal on the cathode.

(b) To recharge a battery A car battery powers the car through a spontaneous reaction, but what can you do if the battery dies?

(c) To coat one metal on top of another one, as with jewelry, or exhaust pipes. a. To make something look more expensive or shinier b. To improve corrosion resistance E l e c t r o c h e m i s t r y P a g e | 10

Using Electrolysis for Electroplating

An example of electrolytic cells in which an electric current is used to plate an object.

The substance to be plated is at the ______. The substance that does the plating is at the ______.

If I wanted to gold plate a necklace, where should I put the gold and where should I put the necklace in the electrolytic cell?

Why is it necessary to use an electrolytic cell to gold plate almost all other metals instead of a voltaic cell?

Would I need to use an electrolytic cell to plate nickel on top of an iron pipe? How do you know? E l e c t r o c h e m i s t r y P a g e | 11

Labeling Electrolytic Cells

Label the following for each of the electrolytic cells below. a. Anode e. Show electron flow using arrows. b. Cathode f. Label ions. c. Reduction Electrode g. Has the anode increased or decreased in mass? d. Oxidation Electrode h. Has the cathode increased or decreased in mass?

(a) Zn/Cu in CuSO4 (b) Ni/Ag in AgNO3

(c) Cu/Pb in CuCl2 (d) Ag and a ring in AgNO3 E l e c t r o c h e m i s t r y P a g e | 12

Calculating Cell Potentials

It is possible to calculate the potential difference between two half-cells using standard reduction potentials. All standard reduction potentials are tabulated relative to the Standard Hydrogen Electrode (SHE). The more positive the reduction potential value, the more likely the substance is to undergo reduction.

Electrochemical Potential or Electromotive Force (emf)

E°cell = E°red (cathode) - E°red (anode)

Determining Cell Potentials

Predict the voltages or emf produced by the following cells.

1. Zn/Zn+2//Fe+2/Fe

2. +2 - Mn/Mn //Br2/Br E l e c t r o c h e m i s t r y P a g e | 13

3. +2 +2 Ni/Ni //Hg2 /Hg

4. Cu/Cu+2//Ag+/Ag

5. Using standard reduction potentials, calculate the standard emf for each of the following reactions:

(a) + - H2 (g) + I2 (s)  2H + 2I

(b) Ni(s) + 2Ce+4 (aq)  Ni+2 (aq) + 2Ce3+(aq)

(c) Cr(s) + 2Cr+3 (aq)  3cr+2 (aq)

(d) 2Al3(aq)- + 3Cd(s)  2Al(s) + 3Cd2+(aq)]

6. A voltaic cell is constructed: one electrode compartment has an aluminum strip in contact with a solution of Al(NO 3)3, and the other is a standard hydrogen electrode.

(a) Write the half reactions involved and determine which electrode is the anode and which is the cathode.

(b) Will the aluminum strip gain or lose mass as the cell operates?

(c) Write a balanced equation for the overall cell reaction.

(d) What is the standard emf of the cell?