1

Chemistry Meets

Electricity Electrochemistry

2

Review: Assigning Oxidation Numbers

To identify whether atoms loses or gains electrons, chemists use a model of oxidation numbers, which can help them identify differences in an atom of an

element in different compounds. By following the set of rules described in your homework packet, you can assign an oxidation number to each element in a molecule or in a formula unit. Let’s Practice: Example: Determine the oxidation number of each element in the following substances.

Example #1

___ ___ ___ KMnO4

Example #2

__ __ SO42-

Assigning Oxidation Numbers Complete the table below and assign oxidation numbers to all of the elements in each of the

compounds or ions below.

Mg3N2

Cl2

KClO3

LiH

MnO2

Al(NO3)3

ClO2-

Na

H2CO3

P4

IO3-

Cu2SO4

Assigning Oxidation States in a Chemical Equation

Mg(s) + S(s) ! MgS(s)

3

Introduction: Redox Reaction What happens when electrons are transferred in a chemical reaction?

Why? Silver tarnishes when it comes in contact with sulfur compounds in the air. Copper gets coated in beautiful green patina as it ages. Metals rust or corrode in the presence of air and water. Minerals (ionic compounds) found in ore can be decomposed with the use of electricity to produce pure metals and nonmetals. All of these reactions are

examples of oxidation and reduction, otherwise known as redox reactions. In this activity, you will explore what is happening at the atomic level in redox reactions.

Model 1-Redox Reactions Redox Reactions

4Fe(s) + 3O2(g) ! 2Fe2O3(s)

2I-(aq) + S2O82-(aq) ! I2(s) + 2SO42-(aq)

Zn(s) + Cu2+(aq) ! Zn2+(aq) + Cu(s)

Non-redox Reactions

2AgNO3(aq) + CaCl2(aq) ! Ca(NO3)2(aq) + AgCl(s)

NaOH(aq) + HCl(aq) ! NaCl(aq) + H2O(l)

4

1. In the space above each reaction in Model 1, write the oxidation number for every element. An example is shown below: O O 3+ 2-

4Fe(s) + 3O2(g) ! 2Fe2O3(s) 2. Identify any elements that changed oxidation number in the reactions in Model 1. Connect the starting and ending oxidation numbers with a line. An example is shown below:

0 0 3+ 2- 4Fe(s) + 3O2(g) ! 2Fe2O3(s)

3. Based on the oxidation number analysis you just performed for the reactions in Model 1, what features differentiates redox reaction from non-redox reactions? __________________________________________________________________________________________________________________________________________________________ 4. Identify the following reactions as either redox or non-redox using oxidation numbers as evidence. Identify the type of reaction.

Chemical Reaction Redox or Nonredox Type of Reaction

PbBr2(aq) + 2NaI(aq) ! PbI2(s) + 2NaBr

2H2O(l) ! 2H2(g) + O2(g)

CH4(g) + 2O2(g) ! 2H2O(l) + CO2(g)

LiOH(aq) + HBr(aq) ! LiCl (aq) + H2O(l)

2 NaBr + Cl2 2 NaCl + Br2

Conclusion: What type of reactions are generally redox reactions? __________________________________________________________________________________________ What type of reactions are generally non-redox reactions?

5

Practice Regents Questions:

1. What is the oxidation number of chromium in the chromate ion, CrO42–?

(a) +6 (b) +2 (c) +3 (d) +8

2. In which compound does chlorine have the highest oxidation number? (a) NaClO (b) NaClO2 (c) NaClO3 (d) NaClO4

3. What is the oxidation number of sulfur in Na2S2O3?

(a) -1 (b) +2 (c) +6 (d) +4

4. Given the balanced equation representing a reaction: Fe2O3 + 2Al--> Al2O3 + 2Fe During this reaction, the oxidation number of Fe changes from (a) +2 to 0 as electrons are transferred (b) +2 to 0 as protons are transferred (c) +3 to 0 as electrons are transferred (d) +3 to 0 as protons are transferred 5. Which equation represents an oxidation- reduction reaction? (a) CH4 + 2O2 ==> CO2 + 2H2O (b) H2SO4 + Ca(OH)2 ==> CaSO4 + 2H2O (c) MgCrO4 + BaCl2 ==> MgCl2 + BaCrO4 (d) Zn(NO3)2 + Na2CO3 � ==> 2NaNO3 + ZnCO3 6. Which balanced equation represents a redox reaction? (a) AgNO3 + NaCl ==>AgCl + NaNO3 (b) BaCl2 + K2CO3 �==> BaCO3 + 2KCl (c) CuO + CO==> Cu + CO2 (d) HCl + KOH �==> KCl + H2O 7. Which reaction is an example of an oxidation-reduction reaction? (a) AgNO3 + KI -->AgI + KNO3 (b) Cu + 2 AgNO3 -->Cu(NO3)2 + 2 Ag (c) 2 KOH + H2SO4 -->K2SO4 + 2 H2O (d) Ba(OH)2 + 2 HCl -->BaCl2 + 2 H2O Do Now: Using your glossary, define the following terms:

Oxidation-reduction

(redox) reaction

Oxidation

Reduction

6

Read This! The process of oxidation and reduction can be can be thought of as a transfer of electrons from one atom to another. Thus, one atom gives up electrons and the other atom gains them. As a result of this process, the oxidation numbers of both atoms change. All redox reactions can be divided up into two reactions:

• An oxidation half-reaction • A reduction half- reaction

This allows for better understanding of the electron transfer process. Model 2- Half Reactions

Redox Reaction 0 0 2+ 2-

Mg(s) + S(s) ! MgS(s)

Oxidation Half Reaction Mg0 ! Mg2+ + 2e- Reduction Half Reaction S0(g) + 2e- ! S2-

Questions: 1. What does the “e-“ symbol represent in the oxidation and reduction half-reactions shown in Model 2?

____________________________________

2. Look at the oxidation half-reaction in Model 2. a. Are electrons lost or gained by the atom during the process of oxidation? ________________________________________________________________________ b. Does the oxidation number of the atom involved in the process of oxidation increase or decrease? ______________________ The oxidation number of the atom __________________ as a result of oxidation. 3. Look at the reduction half-reaction in Model 2. a. Are electrons lost or gained by the atom during the process of reduction? __________________________________________________________________________

b. Does the oxidation number of the atom involved in the process of reduction increase or decrease? ________________________ The oxidation number of the atom __________________ as a result of reduction. c. Consider the word “reduction” as it is used in the English language. In reduction half-reaction, what is “reduced”? Use the example in Model 2 to verify your answer.

__________________________________________________________________________

7

“LEO the Lion Goes GER” LEO=Loss of Electrons is Oxidation

GER= Gain of Electrons is Reduction

Another Redox Reaction (Assign an oxidation # to each element in the equation)

2Na(s) + Cl2(g) ! 2NaCl

Oxidation Half Reaction

__________________________

Reduction Half Reaction

__________________________

Reducing Agent Always found on the reactant side of the equation.

Identify the particle is oxidized and acting as the reducing agent.

_______________

Oxidizing Agent Always found on the reactant side of the equation.

Identify the particle that is reduced and acting as the oxidizing agent.

_______________

Summary 2Na(s) + Cl2(g) ! 2NaCl

Oxidation Reduction " Will always have an increase in

oxidation state (reactant! product). " Involves loss of electrons (LEO) " Substance oxidized is also the

reducing agent. " The number of e-‘s lost is written on

the product side of the half reaction. Example:

2Na0 ! 2Na+ + 2e-

" Will always have a decrease in oxidation state (reactant! product)

" Involves gaining of electrons (GER) " Substance reduced is also the oxidizing

agent. " The number of e-‘s gained is written on

the reactant side of the half reaction. Example:

Cl2 + 2e- ! 2Cl-

2Na(s) + Cl2(g) ! 2NaCl

8

For the reactions below, identify the substance oxidized, the substance reduced, the oxidizing agent, the reducing agent, and write out the oxidation and reduction half reactions.

Example: Oxidized Reduced

Mg + Br2 ! MgBr2 Reducing Oxidizing Agent Agent

Oxidation half reaction: Mg0 ! Mg+2 + 2e- Reduction half reaction: 2e- + Br2o ! 2Br-

Fe + Zn+2 ! Fe+2 + Zn

Oxidation half reaction: _________________________

Reduction half reaction: _________________________

3Fe+2 + 2Al ! 3Fe + 2Al+3

Oxidation half reaction: _________________________

Reduction half reaction: _________________________

Cu + 2AgNO3 ! Cu(NO3)2 + 2Ag Oxidation half reaction: _________________________

Reduction half reaction: _________________________ Identify the spectator ion: _______________ Spectator ions are not involved in redox reactions.

9

BALANCING HALF REACTIONS

BALANCE THE FOLLOWING HALF REACTIONS WITH REGARD TO MASS (Atoms) AND CHARGE.

1. Fill in any coefficients when appropriate (balancing the # of atoms). 2. Write in the appropriate number of electrons to the correct side of the equation. 3. Label the reaction as oxidation or reduction.

HALF REACTION OXIDATION OR REDUCTION

Ga -------> Ga+3

O3 -------> O-2

2Cr+3 -------> Cr+2

S+6 -------> S-2

F-1 -------> F2

As+3 ------> As+5

10

Use the redox reaction below to answer the following questions.

___Sn0 + ___Ag+ -> ___Sn2+ + ___Ag0

_____1. Identify the particle is oxidized and acting as the reducing agent. _____2. Identify the particle that is reduced and acting as the oxidizing agent. 3. Oxidation half reaction: _______________________________ Reduction half reaction: ______________________________ ____4. Is the number of electrons lost = number of electrons gained? _____________ 5. Balancing using the redox method. 6. Add the appropriate coefficients to the original redox equation.

Use the redox reaction below to answer the following questions.

___Cr0 + ___Pb2+ -> ___Cr3+ + ___Pb0

_____1. Identify the particle is oxidized and acting as the reducing agent. _____2. Identify the particle that is reduced and acting as the oxidizing agent. 3. Oxidation half reaction: __________________________________ Reduction half reaction: _________________________________ ____4. Is the number of electrons lost = number of electrons gained? __________ 5. How would you balance this redox reaction in terms of electrons? 6. Add the appropriate coefficients to the original redox equation.

Summary: 1. The e- on each side must be made equal; if they are not equal, they must be multiplied

by appropriate integers (the lowest common multiple) to be made the same.

2. The half-equations are added together, canceling out the electrons to form one balanced equation.

3. Add the appropriate coefficients to the original redox equation.

11

BALANCING REDOX EQUATIONS Balance the equations below using the half-reaction method.

____Zn(s)0 + ____Cu2+ ! ____Zn2+ + ____Cu(s)0

____Mg(s)0 + ____Fe3+ ! ____Mg2+ + ____Fe(s)0

___Zn(s)0 + ___Ag+ ! ___Zn2+ + ___Ag(s)0

____Ca(s) + ____HNO3(aq) ! _____Ca(NO3)2(aq) + _____H2(s)

____AlBr3(s) + ____Fe(s) ! ____FeBr2(s) + ____Al(s)

12

Practice Regents Questions 1. Given the redox reaction: 2I-(aq) + Br2(l) -> 2Br-(aq) + I2(s) What occurs during this reaction?

(A) The I- ion is oxidized, and its oxidation number increases. (B) The I- ion is oxidized, and its oxidation number decreases. (C) The I- ion is reduced, and its oxidation number increases. (D) The I- ion is reduced, and its oxidation number decreases.

2. In which type of chemical reaction are electrons transferred? (A) organic addition (B) oxidation-reduction (C) double replacement (D) acid-base neutralization

3. In an oxidation-reduction reaction, the total number of electrons lost is (A) equal to the total number of electrons gained (B) equal to the total number of protons gained (C) less than the total number of electrons gained (D) less than the total number of protons gained

4. Given the balanced equation representing a reaction: 4Al(s)+ 3O2(g) → 2Al2O3(s) As the aluminum loses 12 moles of electrons, the oxygen (A) gains 4 moles of electrons (C) gains 12 moles of electrons

(B) loses 4 moles of electrons (D) loses 12 moles of electrons

5. Given the balanced ionic equation: 3Pb2+(aq) + 2Cr(s) → 3Pb(s) + 2Cr3+(aq)

What is the number of moles of electrons gained by 3.0 moles of lead ions? (A) 5.0 mol (B) 3.0 mol (C) 2.0 mol (D) 6.0 mol

6. Which half-reaction correctly represents reduction?

(A) Cr3+ + 3e- -> Cr(s) (B) Cr3++ -> Cr(s) + 3e- (C) Cr(s) -> Cr3+ + 3e- (D) Cr(s) + 3e- -> Cr3+

7. Which quantities are conserved in all oxidation-reduction reactions?

(A) charge, only (B) mass only (C) both charge and mass (D)neither charge and mass

8. Given the reaction: __Mg + __Cr3+ -> __Mg2+ + __Cr When the equation is correctly balanced using smallest whole numbers, the sum of the

coefficients will be (A) 10 (B) 7 (C) 5 (D) 4

9. When a substance is oxidized, it (A) loses protons (B) gains protons (C) acts as an oxidizing agent (D) acts as a reducing agent

13

Chemistry Meets Electricity The science of electrochemistry deals with the connections between chemistry and electricity. It is an important subject because it is involved in many of the things you use everyday. Electrochemical devices change electrical energy into chemical energy and vice versa. A simple flashlight is an everyday example of something that converts chemical energy into electrical energy, which is then converted into light energy.

Electrochemical Cells

Zn0 + Cu2+! Zn2+ + Cu0 + energy

energy + Cu0 + Mg2+ ! Cu2+ + Mg0

" Redox reaction producing energy " Chemical energy is converted to

electrical energy " The more active metal (Table J)

undergoes oxidation. " Spontaneous Redox Reaction

" Redox reaction requiring energy " Electrical energy is converted to

chemical energy. " The less active metal (Table J)

undergoes oxidation. " Nonspontaneous Redox Rxn.

Galvanic Cell (Voltaic Cell)

Electrolytic Cell

14

What Makes Up A Galvanic Cell (Voltaic Cell)? (Spontaneous Redox Reaction)

1. Using Table J, which is more easily oxidized, zinc (Zn) or Copper (Cu)?________ 2. Determine the following half-cells reactions: Oxidation half reaction: _______________________________________ Reduction half reaction: _______________________________________ 3. Write out the balanced redox equation. __________________________________________________ 4. Which piece of solid metal loses mass (gets smaller) as the reaction proceeds? ____________ 5. Which piece of solid metal is gaining mass (gets bigger) as the reaction proceeds? _________ 6. As the voltaic cell (battery) produces electricity, the anode will:

a. lose electrons b. gains electrons c. not react 7. A salt bridge is needed in a battery (voltaic cell) so that: a. the salt has somewhere to go b. the circuit is completed and electrons can travel through aqueous solutions c. salt ions can move into the solutions to balance the overall charge in each solution and keep it neutral. d. it has no special function and is not needed.

Half Cell 1 Half Cell 2

15

Using the Galvanic cell above, answer the following questions:

______1. At which electrode does oxidation occur? (Hint: Table J) ______2. Which electrode is the anode?

______3. What is the charge of the anode?

4. As the redox reaction proceeds, what happens to the mass of the anode? _________________ ______5. At which electrode does reduction occur? ______6. Which electrode is the cathode? ______7. What is the charge of the cathode?

8. As the redox reaction proceeds, what happens to the mass of the cathode? _______________ 9. What is the purpose of the salt bridge? ______________________________________________________________________________________________________________________________________________ 10. Using arrows( !) indicate the flow of electrons.

Electrons will flow from the________________ to _________________ Electrons will flow from the________________ to _________________

11. What can be added to this cell to measure the amount of electrical energy produced?

__________________________

Practice Regents Questions 1. Given the balanced equation representing a reaction: 2Al(s) + 3Cu2+(aq) → 2Al3+(aq) + 3Cu(s) Which particles are transferred in this reaction?

(A) electrons (B) positrons (C) neutrons (D) protons 2. In an operating voltaic cell, reduction occurs (A) at the anode (B) in the salt bridge (C) at the cathode (D) in the wire

16

3. A small digital clock can be powered by a battery made from two potatoes and some household materials. The “potato clock” battery consists of two cells connected in a way to produce enough electricity to allow the clock to operate. In each cell, zinc atoms react to form zinc ions. Hydrogen ions from phosphoric acid in the potatoes react to form hydrogen gas. The labeled diagram and balanced ionic equation below show the reaction, the materials, and connections necessary to make a “potato clock” battery.

a. State the direction of electron flow in wire A as the two cells operate.

_____________________________________________ b. Write a balanced half-reaction equation for the oxidation that occurs in the “potato clock” battery.

___________________________________________________ c. Explain why phosphoric acid is needed for the battery to operate. ________________________________________________________________________________

4. Base your answers to questions a –c on the diagram of a voltaic cell and the balanced ionic equation below.

a. What is the total number of moles of electrons needed to completely reduce

6.0 moles of Ni2+(aq) ions? ____________

b. Identify one metal from Reference Table J that is more easily oxidized than Mg(s). ____________

c. Explain the function of the salt bridge in the voltaic cell. _________________________________________________________

17

What Makes Up An Electrolytic Cell?

(Non-Spontaneous Redox Reaction)

Extraction of Metals by Molten Salts Electrolysis

The decomposition of sodium chloride is a non-spontaneous redox reaction. Use the QR code to watch a video describing the

extraction of pure metals from molten salts. 1. Explain why NaCl(s) cannot conduct electricity? ________________________________________________________________________________________ 2. How can sodium chloride become an electrolyte? ________________________________________________________________________________________ 3. Write out the balanced equation for the decomposition of sodium chloride into its elements.

________________________________________________________________________________ 4. Complete the following statements: a. Sodium ions (Na+) will be attracted to the _______________ electrode known as the _________________ resulting in an _______________________ reaction to produce ________________ metal.

Write out the balanced half reaction:

______________________________________________________

b. Chloride ions (Cl-) will be attracted to the ______________electrode known as the ___________________ resulting in a ____________________ reaction to produce ______________ gas.

Write out the balanced half reaction:

______________________________________________________

18

4. The electrodes are usually made up of what type of material? _______________________________________________________________________

5. Complete the following for the electrolysis of molten lead (II) bromide.

Write out the decomposition reaction of lead (II) bromide

_____________________________________________________ Anode

Half Reaction

_________________________________________ Cathode

Half Reaction

_________________________________________

6. In terms of the electrodes, how is an electrolytic cell similar to a voltaic cell? ____________________________________________________________________ 7. In terms of the electrodes, how is an electrolytic cell different from a voltaic cell?

______________________________________________________________________________________ 8. Electrons flow from the _______________ to the _____________________

Electroplating a Key

Using your glossary, define electroplating. _________________________________________________________________________________________________________________________________________________________________________________________________________

19

Electrolysis of H2O (Demonstration)

Write out the balanced equation for the decomposition of water

_____________________________________

Practice Regents Questions 1. The diagram below represents an operating electrolytic cell used to plate silver onto a nickel key. As the cell operates, oxidation occurs at the silver electrode and the mass of the silver electrode decreases.

a. State the purpose of the power source in the cell. ___________________________________________________________________________________

b. Explain, in terms of Ag atoms and Ag+(aq) ions, why the mass of the silver

electrode decreases as the cell operates.

1. a. Explain why AgNO3 is a better choice than AgCl use in this electrolytic process.

___________________________________ ___________________________________ b. Explain the purpose of the battery in this cell. ___________________________________ ___________________________________

20

2. The diagram below shows a system in which water is being decomposed into oxygen gas and hydrogen gas. Litmus is used as an indicator in the water. The litmus turns red in test tube 1 and blue in test tube 2. The oxidation and reduction occurring in the test tubes are represented by the balanced equations below.

Test tube 1: 2H2O(l) → O2(g) + 4H+(aq) + 4e−

Test tube 2: 4H2O(l) + 4e− → 2H2(g) + 4OH− (aq)

a. Determine the change in oxidation number of hydrogen during the reaction in test tube 2. ______________________

b. Explain, in terms of the products formed in test tube 2, why litmus turns blue in test tube 2. _________________________________________________________

3. Aluminum is one of the most abundant metals in Earth’s crust. The aluminum compound found in bauxite ore is Al2O3. Over one hundred years ago, it was difficult and expensive to isolate aluminum from bauxite ore. In 1886, a brother and sister team, Charles and Julia Hall, found that molten (melted) cryolite, Na3AlF6, would dissolve bauxite ore. Electrolysis of the resulting mixture caused the aluminum ions in the Al2O3 to be reduced to molten aluminum metal. This less expensive process is known as the Hall process. a. Write the oxidation state for each of the elements in cryolite.

b. Write the balanced half-reaction equation for the reduction of Al3+ to Al.

c. Explain, in terms of electrical energy, how the operation of a voltaic cell

differs from the operation of an electrolytic cell used in the Hall process. Include both the voltaic cell and the electrolytic cell in your answer.