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Chemistry 30 Worksheets

Electrochemistry

Chemistry 30 Worksheets Introduction to Redox Chemistry

1. Describe the difference between an atom and an ion.

2. Write a chemical equation that shows the formation of the following ions. a. Bromide ions from a bromine molecule.

b. Copper (II) ions from a copper atom.

c. Phosphide ions from a phosphorous molecule.

3. When a metal atom forms an ion, the atom _________ electrons to form a __________________ charged

ion.

4. When a non-metal atom forms an ion, the atom _________ electrons to form a __________________

charged ion.

5. Write a chemical equation to show how the following substances behave when placed in water. a. Calcium iodide

b. Perchloric acid

c. Chlorous acid

d. Ammonia

e. Chlorine

f. Copper (II) sulfate

6. Predict the products of the following reactions, and write a net ionic equation. Then write the oxidation and reduction half reactions.

a. Zn(s) + AgNO3 (aq) à

Chemistry 30 Worksheets b. Cl2 (aq) + KI(aq) à

c. Al(s) + HCl (aq) à

d. NaCl(aq) + CuSO4(aq) à

Operational and Theoretical Definitions

1. Write an empirical definition for each of the following terms. a. Reduction

b. Oxidation

c. Metallurgy

d. Reducing Agent

e. Oxidizing Agent

2. For each of the following equations o Classify as oxidation or reducing according to the empirical (operational) definition. o Identify the oxidizing or reducing agent a. 4 Fe(s) + 3 O2 (g) à 2 Fe2O3 (s) b. 2 PbO(s) + C(s) à 2 Pb(s) + CO2(g) c. NiO(s) + H2(g) à Ni(s) + H2O(l) d. Sn(s) + Br2(l) à SnBr2(s)

3. What class of elements behaves as oxidizing agents for metals?

4. Write theoretical definitions for each of the following a. Redox reaction

b. Reduction

c. Oxidation

5. For each of the following equations, write the oxidation and reductions half reactions.

a. Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s) b. Mg(s) + 2 H+(aq) à Mg2+(aq) + H2(g)

Identifying Redox Reactions using Oxidation Numbers

1. H2SO4 + Ca(OH)2 à CaSO4 + 2H2O

2. 4 Fe(s) + 3 O2(g) à 2 Fe2O3(s)

3. C3H8(g) + 5 O2(g) à 3 CO2(g) + 4 H2O(g)

4. CO2(g) + H2O(l) à H2CO3(aq)

5. Cr2O72- + C2H5OHà 2 Cr3+ + 3 CH3COOH + H2O

6. 2 H2O2 (l) à 2 H2O (g) + O2 (g)

Using Oxidation Numbers to Balance Redox Reactions Use this reaction to walk through the steps in balancing by this method. Acidified dichromate ions react with a solution containing iodide ions to produce iodine and chromium (III) ions. Step 1: Write the skeleton equation, if not already given. Step 2: Assign oxidation numbers and note any changes. Step 3: Separate the equation into two half reactions. Balance the atoms, excepting O and H in each equation. Step 4: Balance electrons in each half reaction by looking at the change in oxidation numbers.

• in oxidation, electrons will be the product • in reduction, electrons will be the reactant.

Step 5: Equalize the electrons lost and gained by multiplying one or both equations by a coefficient. Then add the two reactions together. Step 6: Balance the oxygen and hydrogen.

*in an acidic solution add water to balance the oxygen and hydrogen ions to balance the hydrogen. *In a neutral solution, water only can be added

Step 7: Check your work to ensure that atoms and charges balance. Note: the same method can be used to balance half reactions, just omit steps 3 and 5. Example: MnO4

- à Mn2+ in an acidic solution. Step 1: skeleton equation already given. Step 2: Assign oxidation numbers and note any changes. Step 4: Balance electrons in each half reaction by looking at the change in oxidation numbers.

Step 6: Balance the oxygen and hydrogen. Step 7: Check your work to ensure that atoms and charges balance. Example 2 Hydrogen sulfide reacts with hydrogen peroxide to form sulfur and water. Step 1: Write the skeleton equation. Step 2: Assign oxidation numbers and note any changes. Step 3: Separate the equation into two half reactions. Balance the atoms, excepting O and H in each equation. Step 4: Balance electrons in each half reaction by looking at the change in oxidation numbers. Step 5: Equalize the electrons lost and gained by multiplying one or both equations by a coefficient. Then add the two reactions together. Step 6: Balance the oxygen and hydrogen.

Step 7: Check your work to ensure that atoms and charges balance. Example 3 (disproportionation) Nitrous acid reacts to form nitric acid, nitrogen monoxide and water. Step 1: Write the skeleton equation, if not already given. Step 2: Assign oxidation numbers and note any changes. Step 3: Separate the equation into two half reactions. Balance the atoms, excepting O and H in each equation. Step 4: Balance electrons in each half reaction by looking at the change in oxidation numbers. Step 5: Equalize the electrons lost and gained by multiplying one or both equations by a coefficient. Then add the two reactions together. Step 6: Balance the oxygen and hydrogen.

Step 7: Check your work to ensure that atoms and charges balance.

Assignment: Balancing Redox Reactions Using Oxidation Numbers. 1. MnO4

-(aq) + C2O42-(aq) à Mn2+ (aq) + CO2(g) (acidic solution)

2. CuO + NH3 à N2 + H2O + Cu (neutral solution)

3. CH3OH + MnO4– à Mn2+ + CH2O (acidic solution)

4. H2 + Fe2O3 à FeO + H2O ( neutral solution )

5. Fe2+ + Cr2O72– à Cr3+ + Fe3+ (acidic solution)

6. H2O2 + Fe2+ à Fe3+ + H20 (acidic solution)

7. Photosynthesis

Disproportionation Reactions

8. PbSO4 à Pb + PbO2 + SO42—

(acidic solution)

9. Cl2 + H2O à HOCl + H+ + Cl– ANSWERS 1. 2 MnO4

-(aq) + 16 H+ + 5 C2O42-(aq) à 2 Mn2+(aq) +10 CO2(g) + 8 H2O

2. 3 CuO + 2 NH3 à N2 + 3 H2O + 3 Cu 3. 5 CH3OH + 2 MnO4

– + 6 H+ à 2 Mn2+ + 5 CH2O + 8 H2O 4. H2 + Fe2O3 à 2 FeO + H2O 5. 6 Fe2+ + Cr2O7

2– + 14 H+ à 2 Cr3+ + 6 Fe3+ + 7 H2O 6. H2O2 + 2 Fe2+ + 2 H+ à 2 Fe3+ + 2 H20 8. 2PbSO4 + 2H2O à Pb + PbO2 + 2 SO4

2— + 4H+ 9. Cl2 + H2O à HOCl + H+ + Cl–

Spontaneous Redox Reaction Assignment

1. A student is required to store an aqueous solution of iron (III) nitrate. She has a choice of a copper, tin, iron, or silver container. Use the redox table to predict which container would be the most suitable. Explain your reasoning.

2. An analytical chemist reacts an unknown metal, X, with a copper (II) sulfate solution. The copper forms a coating (plates) on the metal X. X does not react with aqueous zinc nitrate.

a. What is the order in which these metallic ions tend to react? Place the most reactive first.

b. What groups of metals are eliminated as a possible identity of the unknown metal X?

c. What other solutions might be chosen to help identify the metal?

3. Why is gold more commonly found as a pure solid and not as an ore?

a. Does this help explain why it is a good choice for jewelry?

4. Why is lithium metal so rare?

5. Write the reduction and oxidation half reactions for Sn2+ (aq), Fe2+ (aq), Cr2+(aq), and H2O(l). Worksheet: Building Redox Tables Examples A B Co (s) + Pd2+ (aq) à Co2+(aq) + Pd(s) Pd(s) + Pt2+(aq) à Pt (s) + Pd2+ (aq) Mg (s) + Co2+ (aq) à Mg2+(aq) + Co(s)

L4+ M3+ N2+ O+ L (s) X √ √ X M (s) X X X X N (s) X √ X X O (s) √ √ √ X

1. Construct a redox table for each of the following reactions:

Co2+(aq) + Zn(s) → Co(s) + Zn2+(aq) Mg2+ + Zn(s) → no evidence of reaction

2. Construct a redox table for each of the following reactions:

Be(s) + Cd2+(aq) → Be2+(aq) + Cd(s) Cd(s) + 2 H+(aq) → Cd2+(aq) + H2(g)

Ca2+(aq) + Be(s) → no evidence of reaction Cu(s) + H+(aq) → no evidence of reaction

3. Construct a redox table using the following evidence:

I2(aq) Cu2+(aq) Ag+(aq) Br2(aq) I–(aq) x x √ √ Cu(s) √ x √ √ Ag(s) x x x √

Br–(aq) x x x x

4. Construct a redox table for these reactions. Cl2 (g) + 2Br - (aq) à 2Cl- (aq) + Br2(aq) 2 Ag(s) + Br2(aq)à 2 Ag+ (aq) + 2Br - (aq) 2 Ag(s) + I2(aq)à 2 Ag+ (aq) + 2I - (aq)

5. Construct a redox table for these reactions. Four non-metallic oxidizing agents X2, Y2, Z2, and W2 are combined with solutions of ions, X-, Y-, Z-, W-. The following results were recorded; X2 reacts with W- and Y-, only Y- will reduce W2

• metal ions, nonmetals and solutions containing H+(aq) ions are OA

• nometal ions, metals and solutions containing OH–(aq) ions are RA

• some species are an OA or RA combination; for example MnO4–(aq) and H+(aq)

• some species are on both sides of the redox table and may act as either OA or RA: o H2O(l) o Sn 2+(aq) o Cr2+(aq) o Fe2+(aq)

Worksheet: predicting redox reactions using the half-reaction table 1. Write the redox reaction and predict the spontaneity for the following:

a) Concentrated nitrous acid is poured on to a strip of zinc.

b) Hydrochloric acid is poured onto a gold ring.

c) Aluminum lawn furniture is exposed to the action of wind (O2)and rain (H2O).

d) Tin (II) bromide solution is added to acidic potassium permanganate solution.

e) An aqueous solution of potassium permanganate was reacted with an acidic solution of sodium bromide.

f) A strip of silver metal is placed in solution of nickel (II) chloride.

Generalizations in identifying OA and RA using the half-reaction table

g) Liquid mercury is mixed with an acidified paste of manganese (IV) oxide.

h) Hydrogen peroxide and silver nitrate solutions are mixed.

i) Potassium metal is placed in water.

j) In a car battery, lead and lead (IV) oxide electrodes are exposed to a sulfuric acid electrolyte.

2. Use the following hypothetical reaction to answer the next question:

Q2(g) + 2R– (aq) à R2(l) + 2Q–(aq)

a) write the half reaction for the species that gains electrons. Is this oxidation or reduction? b) write the half reaction for the species that loses electrons. Is this oxidation or reduction? c) which species is the oxidizing agent:_________

d) which species is the reducing agent:_________

e) identify the species that has the greatest strength of attraction of electrons:_______

f ) identify the species that has the least strength of attraction of electrons:__________

Redox Stoichiometry: Show your work for each question.

1. A 10.0 mL acidified sample of a 0.0774 mol/L solution of FeSO4•(NH4)2SO4•6H2O (aq) is titrated with an average volume of 13.6 mL of KMnO4 (aq). Calculate the concentration of the KMnO4 (aq) (11.4 mmol/L)

2. If the concentration of bromide ions in seawater is 0.40 mol/L, what mass of chlorine gas would be required to oxidize all of the bromide ions in 3.00 kL of seawater? (43 kg)

3. A 0.125 mol/L of potassium dichromate solution is used to titrate 10.0 mL of chromium (II) sulfate in an acidic solution. What is the concentration of chromium (II) ions? (1.31mol/L)

Titration of Chromium (II) Sulfate Solution Trial 1 2 3 4

final buret reading (mL) 15.8 34.9 18.9 34.1 initial buret reading (mL) 3.4 17.5 1.5 11.0 volume of K2Cr2O7 (aq)

4. What volume of 0.10 mol/L sivler nitrate solution will react completely with 25.0g of nickel metal? (8.5L)

5. Fluoride treatments of children’s teeth have been found to significantly reduce tooth decay. When this was first discovered, toothpastes were produced containing tin (II) fluoride. 0.0832 mol/L of a potassium dichromate solution is used to titrate 10.0 mL sample of acidified tin (II) fluoride solution prepared for research on toothpaste. What is the concentration of the tin (II) fluoride solution? (0.310 mol/L)

Titration of Tin (II) Fluoride Solution Trial 1 2 3

final reading (mL) 15.8 28.1 40.6 initial reading (mL) 3.4 15.8 28.1

Volume of K2Cr2O7 (aq) Extra Practice: Redox Stoichiometry 1. An aluminum strip was placed into a solution of nickel (II) nitrate. Calculate the mass of nickel that forms onto

the strip if 40.0 mL of a 0.200 mol/L nickel (II) nitrate solution is used. 2. In a redox titration 30.0 mL of a 0.0500 mol/L solution of potassium dichromate was used to oxidize 10.0 mL of

an acidified solution containing Fe2+. What is the concentration of iron (II) ions in the solution? 3. A 25.0 mL of acidified 0.500 mol/L calcium iodide solution was titrated to the endpoint with 0.0500 mol/L

solution of potassium permanganate solution. a) What volume of permanganate ions are used? b) What colour would signal the endpoint of the titration? 4. A student poured 100 mL of a 0.125 mol/L solution of nitric acid into a copper can for storage. A little while

later, he noticed that the copper can was corroding. a) Calculate the mass of copper that was corroded by the nitric acid. b) What material would make a suitable storage container for nitric acid?

Unit 2: Electrochemistry Worksheets

Chapter 12 Review 1. Chlorine is bubbled through an aqueous solution of sodium bromide. Species List:

Reduction half reaction:

Oxidation half reaction:

Net redox reaction:

Is the reaction spontaneous? Explain why or why not.

Describe a diagnostic test to identify one of the products.

2. Aluminum is exposed to moist air. Species List:



Net redox reaction:


Describe a diagnostic test to identify one of the products.

3. Water is poured onto a gold ring. Species List:



Net redox reaction:



Use the following information to answer the question #4. Q2+(aq) + 2R (s) à Q (s) + 2R+(aq) Q2+(aq) + E (s) à no reaction 2P+(aq) + E (s) à 2P (s) + E2+(aq) 4a) Construct a redox table for the above four species. 4 b) Numerical Response : The order of oxidizing agents, from strongest to weakest is: ______, ______, ______, ______ Use the following hypothetical reaction to answer the next question: Q2(g) + 2R- (aq) à R2(l) + 2Q-(aq) 5. a) write the half reaction for the species that gains electrons. Is this oxidation or reduction? How do you know? b) Identify the oxidizing agent:_________ Identify the reducing agent:_________ c) identify the species that has the greatest strength of attraction of electrons:_________ d ) identify the species that has the least strength of attraction of electrons:__________ 6. Prove that the following is an acidic disproportionation reaction by providing both half rections.

Cl2 à HOCl + Cl- 7. Balance the following redox reaction that is reacting in an acidic solution.

CH3NO2 + Ti3+ à CH3NH2 + Ti4+


8. A redox titration was completed by titrating 10.0 mL of aqueous tin (II) nitrate with acidified 0.0955 mol/L potassium dichromate solution. If, on average, 12.4 mL of potassium dichromate solution were required for complete reaction what is the molar concentration of the tin (II) nirate solution? (0.355 mol/L)


How can you identify redox reactions?

1) In a redox reaction, oxidation and reduction are taking place simultaneously:

Oxidation Reduction

2) Oxidation Numbers

3) Use Oxidation Numbers of Identify 1) Disproportionation Reactions 2) Redox Reactions in Living Systems

4) Balancing Redox Reactions Using Oxidation Numbers


5) Spontaneous and Non-‐Spontaneous Reactions

6) Building Redox Tables

7) How do you predict the products of a redox reaction?

8) Redox Stoichiometry and Redox Titrations


Voltaic (Galvanic) Cells Any device that uses a _________________________________________________________________to transfer _______________________________________________________________ potential energy into _______________________________________________________ energy. Write the half reactions for; oxidation of zinc metal reduction of copper (II) ions. Voltage is a measure of ______________________________________________________________________ And it has the units of ____________________ or ____________________________. Voltage has the symbol _____ Charge is measured in Coulombs ________. Label the following diagram. _______________________ ____________________________


_____________________________________________ ________________________________________________ A cathode is _________________________________________________________________________________ and it

is the metal that is ____________________________________________.

An anode is _________________________________________________________________________________ and it

is the metal that is ____________________________________________.

A salt bridge is used to solve the problem of

_________________________________________________________________________________________________

Note: This reaction can only continue until one of the reactants is completely used up.

Note: The voltage produced by the cell decreases as the concentration of the reactants

decreases.

Half Cells

If we look at just the oxidation reaction at the ____________________________, then we are

looking at one-‐half of the cell. It is not desirable to have the half cell react

______________________________ because the electrons that are transferred cannot

____________________________________________. One way to prevent this is to use an electrolyte in


which the cations are those of _____________________________________.

No reaction can take place between the copper metal electrode and the copper (II) ions in

the solution.

In order to keep the reaction going, other ions must be available to keep the solution

________________________________________. A couple of options are __________________________________

or a ___________________________________________________________________.

Salt Bridge

This is a hollow tube filled with an _________________________(nonreactive) electrolyte such as

sodium nitrate, sodium sulfate or potassium chloride.


Porous Cup or membrane A material, such as unglazed ceramics, that allow _____________________________ to pass but

does not let the solutions mix. Again, ____________________________move to the cathode and

____________________________ to the anode.


Assignment For each of the following: Label the cathode, anode, electron movement, ion movement, and write the half reactions taking place at each half cell. Describe one observation that could be made at each half cell that would indicate the cell is functioning.

#1 #2


Draw a diagram of the following cells. Include the labels and equations as above. #3

Ag|Ag+ ||Fe2+|Fe3+ #4

C(s)|Cr2O72-‐ (aq), H+(aq)||Cu2+(aq)|Cu(s) For each of the following cells, use the given cell notation to identify the strongest


oxidizing and reducing agents. Write chemical equations to represent the cathode, anode, and net cell reactions. Label electrodes, electrolytes, electron flow, and ion movement. 1. Cd(s) | Cd(NO3)2(aq) || AgNO3(aq) | Ag(s)

2 . Pt(s) | H+,K2Cr2O7(aq) || PbSO4(aq) | Pb(s)


3. For the following cell, what is the cathode half reaction anode half reaction

• purpose of the KNO3 solution On the diagram, label the movement of electrons through the cell.

Is the Sn(s) necessary for the cell to function, or could it be replaced by an inert electrode? Explain. Is the Fe(s) necessary for the cell to function, or could it be replaced by an inert electrode? Explain. Is the Sn(NO3)(aq) necessary for the cell to function, or could it be replaced by an inert electrolyte? Explain.


Chemistry 30: Cell Potential Worksheet

1) a) Label the cathode and the anode. b) Calculate the cell potential. 2. Consider the following

voltaic cell. a) Describe the purpose of the salt bridge and identify one substance that might be used in it. b) Identify compound P and metal Q. c) Deduce the half-‐equation for the reaction in the left-‐hand cell.

d Calculate the cell potential 3.

a) Label the cathode and the anode. b) Using the reading on the voltmeter, calculate the reduction potential for the Pd/Pd2+ half cell. c )Using an arrow, show the movement of electrons through the cell.


4. Underground iron pipes in contact with moist soil are likely to corrode. This corrosion can be prevented by applying the principles of electrochemistry. Connecting an iron pipe to a magnesium block with a wire creates an electrochemical cell. The magnesium block acts as the anode and the iron pipe acts as the cathode. A diagram of this system is shown below.

a) Describe the movement of electrons in the voltaic cell produced when magnesium is connected to the iron pipe.

b) Explain how corrosion of the iron pipe is prevented using this protection system. c) Explain, in terms of reactivity, why magnesium is preferred over zinc to protect underground iron pipes. d) If the magnesium block is not attached to the iron pipe it will corrode. If the iron corrodes identify the oxidation reaction that takes place If the iron corrodes identify the reduction reaction that takes place 5. Calculate the cell potential of the following voltaic cells given the net cell reaction. Write the reaction

taking place at the cathode and anode. a) Sn4+(aq) + Co(s) -‐-‐-‐> Sn2+(aq) + Co2+(aq) b) S(s) + 2 H+(aq) + Pb(s) + SO42-‐(aq) -‐-‐-‐> H2S(aq) + PbSO4(s) c ) 2 AgBr(s) + Cd(s) -‐-‐-‐> Cd2+(aq) + 2 Ag+(aq) + 2 Br-‐(aq)


Assigning Reference Values In this Thought Lab, you will investigate what happens to calculated cell potentials when the reference half-‐cell is changed. Procedure 1. Choose the half reaction for Al3+ and Al as your reference point and assign a value of 0

V for this half-‐reaction. To make the standard cell potential for the Al3+/Al half-‐reaction equal to zero, you would have to add 1.66 V to the accepted standard reduction potential. To adjust all the reduction potentials to the new reference, you add 1.66 V to each value.

Reduction half-‐reaction Accepted E° (V)

Adjusted E°(V)

F2(g) + 2e− à 2F−(aq) +2.87 Fe3+(aq) + e− à Fe2+(aq) +0.77 2H+(aq) + 2e− à H2(g) 0.00 Al3+(aq) + 3e− à Al(s) −1.66 Li+(aq) + e− →àLi(s) −3.04

2. Use the given standard reduction potentials to calculate the standard cell potentials for the following redox reactions: (a) 2Li(s) + 2H+(aq) à 2Li+(aq) + H2(g) (b) 2Al(s) + 3F2(g) à 2Al3+(aq) + 6F–(aq) (c) 2FeCl3(aq) + H2(g) à 2FeCl2(aq) + 2HCl(aq) (d) Al(NO3)3(aq) + 3Li(s) à 3LiNO3(aq) + Al(s)

3. Repeat your calculations using the new, adjusted reduction potentials. 4. Compare your calculations from Procedure Steps 2 and 3. What effect does changing the zero on the scale of reduction potentials have on:

(a) reduction potentials?

(b) cell potentials?


Dry Cells Dry cells are voltaic cells where the electrolyte has been thickened into a paste

An Alkaline Battery

Cathode Reaction Anode Reaction

• Dry cells stop producing electrical energy when the _____________________ are used up • A battery is a set of voltaic cells connected in _____________

Example: a 9 volt battery is really six 1.5 volt dry cells connected in series (In a series connection, the negative electrode of one cell is connected to the positive electrode of another cell)

A primary cell A secondary cell

A lead-acid battery (car battery) is a secondary cell

Cathode Reaction


Anode Reaction When your car is running, an electric current reverses the cathode and anode reactions This replenishes the reactants so the battery does not go “dead”

Fuel Cells A battery that can be refueled They are designed so the reactants flow into the cell, and the products flow out Fuel cells are more efficient than combustion engines or generators and do not produce greenhouse gases or other polluting gases

Overall Cell Reaction Cell potential


The fuel cell provides a highly efficient conversion of the chemical energy in hydrogen, natural gas, or hydrocarbons into electrical energy, and because of their high energy density (energy per unit weight of the power source), fuel cells are superior to batteries in portable equipment. Corrosion: An Unwanted Voltaic cell

Metals can be oxidized by the oxygen in our atmosphere. Rust is produced when iron is oxidized to form Fe2O3 • x H2O The surface of a piece of iron acts like a voltaic cell

Anode Cathode

Simplified Reaction The Fe(OH)2(s) further reacts to form

Fe2O3 • x H2O Preventing Corrosion

• Paint or enamel coatings prevent air and water from reaching the metal • Galvanizing • Covering iron with zinc

Zinc is more reactive than iron (SRA) so it will be oxidized instead of iron

• Cathodic protection Attaching a more reactive metal to an iron object (Al, Mg, Zn) The more reactive metal is oxidized instead of the iron (sometimes called a sacrificial anode) Must be periodically replaced as they are used up.

Corrosion


Electrolytic Cell Worksheet

1. In an electrolytic cell, the cathode is (negative/positive). Chemicals that come into contact with the (-)

electrode will (gain/lose) electrons and be (oxidized/reduced). 2. Write the change that water goes through at the (-) electrode. 3. Chemicals that come into contact with the (+) electrode will (gain/lose) electrons and be

(oxidized/reduced). The (+) electrode in electrolysis is called the (cathode/anode). 4. Write the change that water goes through at the (+) electrode. 5. Add these two reactions together (make certain the electrons cancel) and write the overall reaction for

the electrolysis of water. 6. Consider electrolysis using an aqueous solution of sodium sulfate. Both the Na+ and H2O will be near the (-) electrode. Which chemical is more likely to be reduced? Both the SO4

2- and H2O will be near the (+) electrode. Which chemical will be oxidized? ____ 7. In the electrolysis of KI(aq)

Both the K+ and H2O will be near the (-) electrode. Which chemical is more likely to be reduced? ____ Both the I- and H2O will be near the (+) electrode. Which chemical is more likely to be oxidized? ____

Write the reactions at each electrode and the overall reaction: Cathode: Anode: Overall: Calculate the cell potential


8. In the electrolysis of CuSO4(aq) Both the Cu2+ and H2O will be near the (-) electrode. Which chemical will be reduced? ____ Both the SO4

2- and H2O will be near the (+) electrode. Which chemical will be oxidized? ____ Write the reactions at each electrode and the overall reaction: Cathode: Anode: Overall: Calculate the cell potential 9. Draw a diagram of an electrolytic cell containing a zinc iodide solution and inert carbon electrodes. • Label the power supply and electrodes, including signs, the electrolyte, and the directions of electron and ion movements. • Write half-reaction and net equations. • Calculate the cell potential, using standard values. 10. List the main similarities between a voltaic cell and an electrolytic cell. 11. What is the key difference between voltaic and electrolytic cells? 12. Explain why a power supply is necessary for an electrolytic cell. 13. A student brought in an old silver medal into the chemistry lab to plate it with copper. He set up the

following cell: What is object 1? What is object 2? Should the medal be placed at the cathode or the anode? Write the reaction taking place at the cathode.


14. Draw an electrolytic cell that could be used to plate an iron ring with gold. Be sure to include all of the necessary parts. In addition, label the anode, solution used and composition of the electrodes.

Chlor-‐alkali cell


Down’s Cell


Hall-Heroult Process


Electroplating


Refining Metals


Electrolysis worksheet 1. Calculate the mass of zinc plated onto the cathode of an electrolytic cell by a

current of 750 mA in 3.25 h. (2.97 g)

2. How many minutes does it take to plate 0.925 g of silver onto the cathode of an electrolytic cell using a current of 1.55 A? (8.90 min)


3. The nickel anode in an electrolytic cell decreases in mass by 1.20 g in 35.5 minutes. The oxidation half-‐reaction converts nickel atoms to nickel (II) ions. What is the average current? (1.85 A)

4. The following two half-‐reactions take place in an electrolytic cell with an iron anode and a chromium cathode:

Oxidation: Fe(s) à Fe2+ (aq) + 2e-‐ Reduction: Cr3+ (aq) + 3e-‐ à Cr(s)

During the process, the mass of the iron anode decreases by 1.75 g a. Find the change in mass of the chromium cathode. (1.09 g) b. Explain why you do not need to know the electric current or the time to

complete part a.


Building Voltaic Cells

Problem What is the measured cell potential of all the possible voltaic cells built from the following half -cells? Zn|Zn2+ Cu|Cu2+ Pb|Pb2+ Ag|Ag+

Experimental Design Six voltaic cells will be built using 0.10 mol/L solutions, metal strips and porous cups. The cell potential will be measured and compared to the predicted values. Manipulated variable: Responding variable: Three controlled variables: Procedure: See overhead Evidence Cathode Anode Predicted Cell

Potential (V) Measured Cell Potential (V)

1.

2.

3.

4.

5.

6.

Analysis Draw a diagram of the cell that had the greatest measured potential. Label the cathode and anode,show the direction of ion movement and electron movement. Write the cathode reaction, anode reaction and net cell reaction. For each half cell describe a piece of qualitative evidence that would indicate that the cell was working.


Conclusion

1. Did the measured cell potentials match the predicted cell potentials? Was there a general trend in the discrepancies between them?

2. Were the cells you built standard cells?

3. Explain why there was a difference between your measured and predicted values.

4. Identify a source of error in this experiment.


Name:

Electrolysis of Potassium Iodide When an aqueous solution is electrolyzed, the electrolyte or water can undergo electrolysis. In this investigation, you will build an electrolytic cell, carry out the electrolysis of an aqueous solution, and identify the products.

Problem What are the products from the electrolysis of a 1 mol/L aqueous solution of potassium iodide? Are the observed products the ones predicted using reduction potentials? Prediction Predict which product(s) are formed at the anode and which product(s) are formed at the cathode. Include a half reaction for each half cell. (4 marks) Materials • 1 mol/L KI • 2 graphite pencil leads, 2 cm long • 1 drop 1% starch solution • 1 drop 1% phenolphthalein • sheet of white paper

• 1 beaker (600 mL or 400 mL) • 1 elastic band • 25 cm clear aquarium tubing • 3 disposable pipettes • 2 wire leads (black and red) with alligator clips • 9-V battery

Procedure 1. Fold a sheet of paper lengthwise. Curl the folded paper so that it fits inside the beaker. Invert the

beaker on your lab bench. 2. Use the elastic to strap the aquarium tubing to the side of the beaker in a U shape, as shown in the diagram. 3. Fill a pipette as completely as possible with 1 mol/L KI solution. Insert the tip of the pipette firmly into one end of the aquarium tubing. Slowly inject the solution into the U-tube until the level of the solution is within 1 cm to 2 cm from the top of both ends. If air bubbles are present, try to remove them by poking them with a toothpick. You may need to repeat this step from the beginning to ensure there are no air bubbles. 4. Attach the black lead to a 2 cm piece of pencil lead. Insert the lead into one end of the U-tube. Attach the red electrical lead to the pencil lead. Insert the lead

into the other end of the U-tube. 5. Attach the leads to the 9-V battery Attach the black lead to the negative terminal and the red lead to the

positive terminal.


6. Let the reaction proceed for three minutes, while you examine the U-tube. Record your observations. Shut off the power source and remove the electrodes. Determine the product formed around the anode by adding a drop of starch solution to the end of the U-tube that contains the anode. Push the starch solution down with a toothpick if there is an air lock. Determine one of the products around the cathode by adding a drop of phenolphthalein to the appropriate end of the U-tube.

Observations Record your qualitative observations in a table. The conclusions that you make in this experiment will be based on your qualitative observations, so it is important that they are detailed and accurate. Your observations should be so detailed that someone who didn’t perform the experiment could understand what had taken place. For this experiment you should have a minimum of five qualitative observations. You should record what you saw and when you saw it. (6 marks) Conclusion Provide an answer to the problem. Use your observations to justify your conclusion. (5 marks) Analysis 1. Sketch the cell you made in this investigation. (4 marks) On your sketch, show: (a) the direction of the electron flow in the external circuit (b) the anode and the cathode (c) the positive electrode and the negative electrode (d) the movement of ions in the cell

Name _________________________


Chapter 13 Lab: Electrolysis of CuSO4(aq)

Problem What are the products of reaction and the mass of metal produced during the operation of an aqueous copper (II) sulfate electrolytic cell over a time period of 10.0 minutes?

Prediction Cathode reaction Anode reaction Net reaction Minimum voltage required for the cell to operate: Calculate the mass of copper that can be plated using a current of 1.67 amps for 10 minutes from the electrolysis of a solution of copper (II) sulfate. Materials -a carbon rod(anode) - copper wire (cathode) -2 connecting wires -blue and red litmus paper -150 mL beaker -9 V battery -copper (II) sulfate solution -electronic balance -paper towel or Kleenex Experimental Design Copper electrodes are placed in a solution of 0.500 mol/L copper (II) sulfate and a 9.00 volt battery produces a direct current of 1.67 amps to the cell. Diagnostic tests are conducted and empirical evidence is gathered. Procedure 1. Clean any tarnish off the copper wire by sanding it gently. 2. Wrap the copper wire around a pencil to make a closely spaced coil. Leave 10 cm of the wire

unwrapped. Measure and record the mass of the cathode. 3. Use the 10 cm of uncoiled wire to secure the coil on the opposite side of the beaker from the anode,

as shown in the diagram. This copper wire will serve as the cathode. 4. Pour 60 mL of the acidified CuSO4(aq) solution into the beaker. Attach the lead from the negative

terminal of the battery to the cathode. Attach the positive terminal to the anode. Do not allow your electrodes to touch.

5. Keep the battery hooked up for 10.0 min. 6. After 10 min, turn off the power. Remove the cathode and rinse it very gently with distilled water, and

then gently dry it. Measure and record the new mass of the cathode.


Evidence (construct a table to record your qualitative and quantitative evidence)

Conclusion (answer the problem) Evaluation 1. Determine the percent difference. 2. One of the products in the experiment was not confirmed by a diagnostic test. How could your design,

and hence procedure, be altered to confirm this product in a more reliable manner. 3. Identify and describe a source of uncertainty that could have contributed to your percent difference. Do not use the following: - lack of standard conditions - resistance in the wires - precision of instruments - significant digits

Documents

Eelctrochemistry_workbook.pdf