name:& && & &&&&&unit&13:&redox&€¦ ·...
TRANSCRIPT
Name: _______________________ Unit 13: Redox
1
REGENTS CHEMISTRY
Name: _______________________ Unit 13: Redox
2
• An oxidation-‐reduction (redox) reaction involves the transfer of electrons (e-‐). (3.2d) • Reduction is the gain of electrons. (3.2e) • A half-‐reaction can be written to represent reduction. (3.2f) • Oxidation is the loss of electrons. (3.2g) • A half-‐reaction can be written to represent oxidation. (3.2h) • In a redox reaction the number of electrons lost is equal to the number of electrons gained. (3.3b) • Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate
that oxidation and reduction have occurred. (3.2i) • An electrochemical cell can be either voltaic or electrolytic. In an electrochemical cell, oxidation occurs
at the anode and reduction at the cathode. (3.2j) • A voltaic cell spontaneously converts chemical energy to electrical energy. (3.2k) • An electrolytic cell requires electrical energy to produce chemical change. This process is known as
electrolysis. (3.2l)
Name: _______________________ Unit 13: Redox
3
For each word, provide a short but specific definition from YOUR OWN BRAIN! No boring textbook definitions. Write something to help you remember the word. Explain the word as if you were explaining it to an elementary school student. Give an example if you can. Don’t use the words given in your definition!
Reduction: _________________________________________________________________________________
Oxidation: _________________________________________________________________________________
Spectator Ion: ______________________________________________________________________________
Half Reaction: ______________________________________________________________________________
Redox Reaction: ____________________________________________________________________________
Electrochemical Cell (Voltaic): _________________________________________________________________
Anode: ____________________________________________________________________________________
Cathode: __________________________________________________________________________________
Salt Bridge: ________________________________________________________________________________
Electrolytic Cell: ____________________________________________________________________________
Name: _______________________ Unit 13: Redox
4
Section 1 Oxidation Numbers
Oxidation Number Rules: 1. All pure non-‐bonded elements have an oxidation number of ______. Elements are not charged. 2. All ions have the _________ charge and oxidation number. 3. All compounds must have a sum of oxidation numbers equal to ______. Compounds are not charged. 4. All group 1 elements have a ______ oxidation number as seen on the periodic table (except H).
Similarly, group two must have ______. Aluminum is ______. 5. If a halogen (F, Cl, Br, I) is at the end of the molecule it is ______. Otherwise you have options for
oxidation numbers if it is in the middle of the compound. 6. If Oxygen is the anion, it will have a charge of ______UNLESS it is with F or an alkali metal in a 1:1 (or
2:2) ratio. 7. H is ______ in the front and ______ in the back. 8. The sum of oxidation numbers for an ion must equal the ion’s charge.
Oxidation numbers are very important in this chapter “Redox Reactions.” Without the complete understanding of how to assign these numbers, we cannot move ahead with this chapter. They are much like ionic charges, except the every element will be assigned a number. The most important rules that cannot be broken are:
• Free elements are zero. • Group 1 is +1 • Group 2 is +2 • Fluorine is -‐1
Assign oxidation numbers to each element in the following (use the Periodic Table to help you)
(a) NaCl Na___ Cl___
(b) H2S H___ S___
(c) H2O H___ O ___
(d) CO2 C ___ O___
(e) H2SO4 H ___ S___ O___
(f) FeCO3 Fe___ C___ O___
(g) AgI Ag___ I___
(h) H2 H___
(i) PbCl2 Pb___ Cl___
(j) BaCO3 Ba___ C ___ O___
(k) Fe2O3 Fe___ O___
(l) I2 I____
(m) BeO Be____ O____
(n) CaF2 Ca____ F ____
(o) FeCl3 Fe____ Cl____
(p) PF5 P____ F____
(q) H3PO4 H____ P____ O___
(r) KCl K ____ Cl____
(s) K2O K ____ O____
(t) O3 O ____
(u) LiH Li ____ H____
(v) HBr H ____ Br____
(w) Li+ Li____
(x) PO43-‐ P____ O___
Name: _______________________
5
Section 2 2 RedOx reactions, spectators, and agents
Key Ideas: Circle the best fit in the parenthesis or fill in the blanks. • Elements oxidize because they (lose/gain) electrons and their charges (increase /decrease). • Elements reduce because they (lose/gain) electrons and their charges (increase /decrease). • Elements that do not change charges are known as ____________________________. • Redox reactions occur as long as reduction and oxidation both occur. • Table ______ shows a list of the most active metals and nonmetals. Active metals tend to ___
electrons and therefore (oxidize/reduce). Metals at the top of this table (oxide/reduce) the best. • Active nonmetals tend to _______ electrons and therefore (oxidize/reduce). Nonmetals at the top of
this table (oxide/reduce) the best.
Redox reactions are usually synthesis reactions, decomposition, combustion or single replacement reactions.
Double replacement and neutralization reactions are NOT redox reactions. Usually they are easy to spot
because if an element goes from being “free” (with an oxidation number of 0) to being in a compound (with a
new oxidation number) it shows there was an exchange of electrons. In the following examples, identify what
type of reaction they are and then state if they are redox reactions.
1. N2 + O2 à 2NO _________________________ ____________
2. Cl2 + 2NaBr à NaCl + Br2 _________________________ ____________
3. 2NaOH + HCl à H2O + NaCl _________________________ ____________
Are these redox?
What type of reaction (S, D, C, SR, or DR) is NEVER redox? ______________________________
Name: _______________________
6
Questions: 1. For the following reactions, identify the redox reactions. For the redox reactions ONLY: identify the
elements as reducing, oxidizing, or spectators.
a. Cl2 + 2KBr à 2KCl + Br2 Is it redox? _____ Spectator: _____ Reducing: _____ Oxidizing: _____
b. Cu + AgNO3 à CuNO3 + Ag Is it redox? _____ Spectator: _____ Reducing: _____ Oxidizing: _____
c. Zn + 2HCl à H2 + ZnCl2 Is it redox? _____ Spectator: _____ Reducing: _____ Oxidizing: _____
d. CaCO3 + HCl à H2O + CO2 + CaCl2 Is it redox? _____ Spectator: _____ Reducing: _____ Oxidizing: _____
2. Identify redox reactions. For redox reactions ONLY: Use table J to determine if the reaction is spontaneous.
a. Cu + 2HClà CuCl2 + H2 Which element oxidizes? _____ Which element reduces? _____ Should it according to table J? _____
b. Mg + 2HClà MgCl2 + H2 Which element oxidizes? _____ Which element reduces? _____
Should it according to table J? _____
Name: _______________________
7
Section 3 Redox Half Reactions
A half reaction shows either the oxidation or reduction portion of a redox equation including if the electrons are gained or lost. A reduction half reaction shows an atom or ion gaining one of more e-‐: Fe3+ + 3e-‐ à Fe Notice that the e-‐ is one the left. An oxidation half reaction shows an atom or ion losing one or more e-‐: Mg à Mg2+ + 2 e-‐ Notice that the e-‐ is one the right. In a half reaction, only one element is shown and the charges must be conserved. To write a half reaction, first assign all the oxidation numbers to all the elements. Second, cross out any elements that are spectators (they do not change oxidation number). Then, write a half reaction showing a change in oxidation state and label which element is being oxidized and which is being reduced. Last, add in the number of electrons needed to conserve the charge.
Directions: For the following examples, determine if they represent oxidation or reduction.
1. e-‐ + Cr+3 à Cr+2
2. Mn+7 à Cr+5 + 2e-‐
3. Mg+2 + 2e-‐ à Mg
4. K à e-‐ + K+1
5. e-‐ + Cu+2 à Cu+1
6. Al+3 à Al + 3e-‐
7. Li+ + e-‐ à Li
8. S-‐2 à 2e-‐ + S
Directions: For the following examples, add in the e-‐ to balance the charge.
1. Cr+5 à Cr+2
2. Mn+4 à Cr+7
3. Ca+2 à Ca
4. Rb à Rb+1
5. Cu+1 à Cu+2
6. B+3 à B
7. N-‐3 à N+4
8. Te-‐2 à Te
Directions: For the following examples, balance the elements first, then add in the e-‐ to balance the charge.
1. Cl-‐ à Cl2
2. Br2 à Br-‐1
3. I2 à I-‐
4. H2 à H+1
5. F-‐1 à F2
6. N+3 à N2
7. S8 à S-‐2
8. O-‐2 à O2
Name: _______________________
8
Directions: For each reaction, decide if it is redox. If it is redox, write the half reactions below. If they are not redox, leave the answers blank.
1. Cr3+ + Fe2+ à Cr2+ + Fe3+
2. F2 + O2-‐ à F1-‐ + O2
3. Sn + N5+ à Sn4+ + N4+
4. NaCl à Na+ + Cl-‐
5. Cu2O à Cu + O2
6. Cl2 + KBr à KCl + Br2
7. CH4 + O2 à CO2 + H2O
8. H3PO4 + Ca(OH) 2 à Ca3(PO4) 2 + H2O
Reduction Oxidation
1. _____________________________________ ___________________________________
2. _____________________________________ ___________________________________
3. _____________________________________ ___________________________________
4. _____________________________________ ___________________________________
5. _____________________________________ ___________________________________
6. _____________________________________ ___________________________________
7. _____________________________________ ___________________________________
8. _____________________________________ ___________________________________
Name: _______________________
9
Fill in the table below. Reducing half Reaction Oxidizing half reaction How
many e-‐ are
trans-‐ferred?
1 H2 + O2 à H2O
2 K + B2O3 à K2O + B
3 C + S8 à CS2
4 N2 + O2 à N2O5
5 Na + O2 à Na2O
6 Cs + N2 à Cs3N
7 N2 + H2 à NH3
8 Li + AlCl3 à LiCl + Al
9 CH4 + O2à CO2 + H2O
10 Li + H2O à LiOH + H2
Name: _______________________
10
Use Table J to help you with the following questions.
1. Write the oxidation and reduction half reaction for: Ca + Cu2+ à Ca2+ + Cu
2. According to Table J, the element higher on the list will oxidize. Which element is oxidizing here and does that mean that this reaction is spontaneous?
3. Write the oxidation and reduction half reaction for Mg + Ca2+ à Mg2+ + Ca
4. Which element is oxidizing in number 3 and is the reaction spontaneous?
5. Which of the following ions is most easily oxidized? a. F-‐ b. Cl-‐ c. Br-‐ d. I-‐
6. Which element is more easily reduced?
a. Cu b. Mg c. Al d. Zn
7. Which element will reduce Mg2+ to Mg? a. Fe b. Ba c. Pb d. Ag
8. Which ion will oxidize Fe?
a. Zn2+ b. Ca2+ c. Mg2+ d. Cu2+
9. Which metal will react spontaneously with Ag+ but not Zn2+? a. Cu b. Au c. Al d. Mg
10. Which reaction will take place spontaneously?
a. Mg + Ca2+ à Mg2+ + Ca b. Ba + 2Na+ à Ba2+ + 2Na c. Cl2 + 2F-‐à 2Cl-‐ +F2 d. I2 + 2Br-‐ à 2I-‐ + Br2
Name: _______________________
11
1. Which reaction is an example of an oxidation-‐ reduction reaction?
(1) AgNO3 + KI → AgI + KNO3 (2) Cu + 2AgNO3 → Cu(NO3)2 + 2Ag (3) 2KOH + H2SO4 → K2SO4 + 2H2O (4) Ba(OH)2 + 2HCl → BaCl2 + 2H2O
2. In an oxidation-‐reduction reaction, reduction is defined as the (1) loss of protons (3) loss of electrons (2) gain of protons (4) gain of electrons 3. When a lithium atom forms a Li+ ion, the lithium atom
(1) gains a proton (3) loses a proton (2) gains an electron (4) loses an electron
4. Which type of reaction occurs when nonmetal atoms become negative nonmetal ions?
(1) oxidation (3) substitution (2) reduction (4) condensation
5. When a neutral atom undergoes oxidation, the atom’s oxidation state
(1) decreases as it gains electrons (2) decreases as it loses electrons (3) increases as it gains electrons (4) increases as it loses electrons
6. In a redox reaction, there is a conservation of
(1) mass, only (2) both mass and charge (3) neither mass nor charge
7. In any redox reaction, the substance that undergoes reduction will (1) lose e-‐ & have a decrease in oxidation number (2) lose e-‐ & have an increase in oxidation number (3) gain e-‐ & have a decrease in oxidation number (4) gain e-‐ & have an increase in oxidation number
8. What occurs during the reaction below?
(1) The manganese is reduced and its oxidation
number changes from +4 to +2. (2) The manganese is oxidized and its oxidation
number changes from +4 to +2. (3) The manganese is reduced and its oxidation
number changes from +2 to +4. (4) The manganese is oxidized and its oxidation
number changes from +2 to +4.
9. Given the balanced equation:
What is the total number of moles of electrons lost by 2 moles of Al(s)?
(1) 1 mole (3) 3 moles (2) 6 moles (4) 9 moles 10.Given the balanced equation: Mg(s) + Ni2+(aq) à Mg2+(aq) + Ni(s) What is the total number of moles of electrons lost by 2 moles of Mg(s)? (1) 1.0 mol (3) 3.0 mol (2) 2.0 mol (4) 4.0 mol 11. Given the equation representing a reaction: Mg(s) + Ni2+(aq) à Mg2+(aq) + Ni(s) What is the total number of moles of e-‐ lost by Mg when 2.0 moles of e-‐ are gained by Ni2+(aq)? (1) 1.0 mol (3) 3.0 mol (2) 2.0 mol (4) 4.0 mol 12. Given the reaction:
Which species undergoes oxidation?
(1) Mg(s) (3) Cl–(aq) (2) H+(aq) (4) H2(g)
Name: _______________________
12
13. Given the redox reaction:
As the reaction takes place, there is a transfer of
(1) electrons from Al to Cr3+ (2) electrons from Cr3+ to Al (3) protons from Al to Cr3+ (4) protons from Cr3+ to Al
14. Given the redox reaction: 2 Fe3+ + 3 Zn à 2 Fe + 3 Zn2+
As the reaction takes place, there is a transfer of electrons
(1) from Fe3+ to Zn (2) from Zn to Fe3+ (3) from Zn2+ to Fe (4) from Fe to Zn2+
__________________________________________________________________________________________ 15. Circle the electrons in the half-‐reactions below and identify as oxidation or reduction. 16. Complete the half-‐reactions below by ADDING in electrons to the correct side in order to equalize charge (show conservation of charge).
(a) Fe2+ à Fe3+ (b) K à K+ (c) Sn4+ à Sn2+ (d) Cr6+ à Cr3+ (e) O2 à 2O–
(f) Mn3+ à Mn4+ (g) Cr2+ à Cr3+ (h) Cl7+ à Cl1+ (i) 3Cl2 à 6Cl– (j) 4H+ à 2H2
17. For each of the following equations, write the reduction half-‐reaction and the oxidation half-‐reaction.
(a) Cl2 + 2 KBr à 2 KCl + Br2 Reduction:
Oxidation:
(b) Cu + 2 Ag+ à 2 Ag + Cu2+ Reduction: Oxidation:
(c) 2 Mg + O2 à 2 MgO Reduction:
Oxidation:
(d) 2 F2 + 2 H2O à 4 HF + O2 Reduction:
Oxidation:
(a) Br2 + 2 e– � 2 Br– (b) Na � Na+ + e–
(c) Ca2+ + 2e– � Ca
(d) Cl2 + 2 e– � 2 Cl– (e) Na+ + e– � Na (f) S2– � S + 2e–
(g) Cu2+ + 2 e– � Cu (h) Fe � Fe2+ + 2 e– (i) Mn7+ + 3 e– � Mn4+
Name: _______________________
13
18. Which half-‐reaction correctly represents oxidation?
(1) Fe(s) → Fe2+(aq) + 2e– (2) Fe2+(aq) → Fe(s) + 2e– (3) Fe(s) + 2e– → Fe2+(aq) (4) Fe2+(aq) + 2e– → Fe(s)
19. Which equation shows a conservation of both mass and charge?
(1) Cl2 + Br– → Cl– + Br2 (2) Cu + 2Ag+ → Cu2+ + Ag+ (3) Zn + Cr3+ → Zn2+ + Cr (4) Ni + Pb2+ → Ni2+ + Pb
20. Given the balanced ionic equation:
Which equation represents the oxidation half-‐reaction?
(1) Zn(s) + 2e– → Zn2+(aq) (2) Zn(s) → Zn2+(aq) + 2e– (3) Cu2+(aq) → Cu(s) + 2e– (4) Cu2+(aq) + 2e– → Cu(s)
21. Which half-‐reaction equation represents
the reduction of a potassium ion? (1) K+ + e– à K (3) K+ à K + e– (2) K + e– à K+ (4) K à K+ + e–
22. Given the equation:
The reduction half-‐reaction is
(1) Al → Al3+ + 3e– (2) Cu2+ + 2e– → Cu (3) Al + 3e– → Al3+ (4) Cu2+ → Cu + 2e–
23. Base your answers to the questions below on the following redox reaction, which occurs in a battery.
____ Zn + ____ Cr3+ à ____ Zn2+ + ____ Cr
(a) Write the half-‐reaction for the reduction that occurs.
(b) Write the half-‐reaction for the oxidation that occurs.
(c) Balance the equation below using the smallest whole-‐number coefficients.
____ Zn + ____ Cr3+ à ____ Zn2+ + ____ Cr (d) Which species loses electrons and which species gains electrons?
(e) State what happens to the number of protons in a Zn atom when it changes to Zn2+ as the redox
reaction occurs.
Name: _______________________
14
24. The outer structure of the Statue of Liberty is made of copper metal. The framework is made of iron. Over time, a thin green layer (patina) forms on the copper surface.
(a) When copper oxidized to form this patina layer, the copper atoms became copper(II) ions (Cu2+). Write a balanced half-‐reaction for this oxidation of copper.
Cu → Cu2+ + 2e– (b) Where the iron framework came in contact with the copper surface, a reaction occurred in which iron
was oxidized. Using information from Reference Table J, explain why the iron was oxidized. Iron is a more active metal. 25. Litharge, PbO, is an ore that can be roasted (heated) in the presence of carbon monoxide, CO, to produce elemental lead. The reaction that takes place during this roasting process is represented by the balanced equation below.
PbO(s) + CO(g) à Pb(ℓ) + CO2(g) (a) Write the balanced equation for the reduction half-‐reaction that occurs during this roasting process. (b) Determine the oxidation number of carbon in carbon monoxide. _____ 26. The catalytic converter in an automobile changes harmful gases produced during fuel combustion to less harmful exhaust gases. In the catalytic converter, nitrogen dioxide reacts with carbon monoxide to produce nitrogen and carbon dioxide. In addition, some carbon monoxide reacts with oxygen, producing carbon dioxide in the converter. These reactions are represented by the balanced equations below.
Determine the oxidation number of carbon in each carbon compound in reaction 2. Your response must include both the sign and value of each oxidation number. +2 for carbon in CO and +4 for carbon in CO2. 27. A flashlight can be powered by a rechargeable nickel-‐cadmium battery. The unbalanced equation below represents the reaction that occurs as the battery produces electricity. When a nickel-‐cadmium battery is recharged, the reverse reaction occurs.
(a) Balance the equation above using the smallest whole-‐number coefficients. 1,1,2,1,1 (b) Determine the change in oxidation number for Cd. from 0 to +2 (c) Explain why Cd would be above Ni if placed on Table J. Cd is more reactive than Ni.
Name: _______________________
15
Section 4 Electrochemical/Voltaic Cells
Key Ideas: • In a RedOx reaction, ___________ are transferred from the oxidizing elements which loses to the
reducing element which gains. • LEO stands for __________________________ and GER stands for ____________________________. • In a voltaic cell, or battery, _______________ occurs at the anode, which is _________ charged. • In a voltaic cell, or battery, _______________ occurs at the cathode, which is _________ charged. • Electrons flow from the _________________ to the ______________ of the battery. • The __________ __________ permits the flow of ions in the voltaic cell in order to maintain the
charges of the anode and cathode.
Directions: In each of the following, determine which element oxidized easier on table J. Then label the anode, cathode, direction of e-‐ flow, and the half reactions.
1. 2.
3. 4.
Name: _______________________
16
Additional Questions: 5. On diagram 1, which way will anions travel through the salt bridge? ____________ 6. On diagram 2, towards which electrode will cations travel through the salt bridge? ____________ 7. On diagram 3, how many e-‐ are exchanged per mole of Mg? ____________ 8. On diagram 4, how many e-‐ are transferred between Ag and Ni? ____________ 9. On all diagrams, at which electrode does oxidation occur? ____________ 10. On all diagrams, at which electrode does reduction occur? ____________ 11. On all diagrams, from which electrode will electrons travel? ____________ 12. What is the purpose of the salt bridge? ____________________________________________________ 13. Describe the change in energy that occurs in voltaic cells in terms of electric and chemical energies:
14. Why does it make sense that electrons flow from the anode to the cathode in terms of charge? Discuss the charges of electrons, the anode, and the cathode in your answer.
15. Using the diagram below, answer the following questions:
a. Which element is present at the anode? _____ b. Which element is present at the cathode? _____ c. Draw an arrow to show the direction of electron flow in the battery, through the wire. d. Where are spectator ions such as Cl-‐ and Na+ located (They are not on the diagram labeled-‐you
have to think about it)? What do they do? e. How many total electrons between one atom of each metal? _____ f. Write the half reaction for magnesium including the electrons.
g. Write the half reaction for copper including the electrons.
Name: _______________________
17
Electrolytic Cells
Key Ideas:
• Electrolytic cells are not spontaneous so you must apply power.
• Electrolytic cells are different from voltaic cells (or batteries) because the electrodes charges are
reversed (however, electrons still travel from ____________ to ______________.
•
1. Which element is oxidizing in the diagram to the left? _____
2. According to Table J, should that element oxidize? ______
3. Is this reaction spontaneous? _____
4. What type of cell is it? ___________________
5. Label the anode and the cathode with charges on the diagram to the left.
6. Write the overall reaction for this cell.
____________________________________________________
7. Water is being decomposed using a battery in the diagram to the right. Write the equation for the decomposition of water.
_______________________________________
8. Which element is being oxidized? ____ How many e-‐ are lost? ___
9. O would be below H on table J, so should that element oxidize? __
10. Is this reaction spontaneous? _____
11. What type of cell is it? ___________________
12. Label where oxidation and reduction on the diagram.
13. Why is more H2 gas being formed that O2 gas?
6. Write the overall reaction for this cell.
Name: _______________________
18
14. Only one element is being used in the diagram to the left. The silver plate ionizes and the ions attach to the spoon. Show a half reaction for silver oxidizing.
____________________________________
15. Label the anode and cathode with charges on the diagram.
16. Explain the direction of e-‐ flow through the wire.
__________________________________________________________
17. Is this reaction spontaneous? _____ How can you tell?
_________________________________________________________
18. Write the half reaction for the anode in the diagram to the left:
_________________________________________________
19. Write the half reaction for the cathode (use Fe+2):
__________________________________________________
20. How many e-‐ are transferred per mole? _____
21. What is the charge of the Zn electrode? _____
22. Why is a power source needed in this cell? __________________
_________________________________________________
23. Show a half reaction for silver reducing.
__________________________________________________
24. Label the anode and cathode on the diagram to the right.
25. What will happen to the mass of the key?
__________________________________________________
26. What will happen to the mass of the silver metal?
__________________________________________________
27. Show the direction of e-‐ flow through the wire on the diagram to the right.
28. State the difference between voltaic and electrolytic cells in terms of electrical and chemical energy.
Name: _______________________
19
Questions:
1. What is the charge of a voltaic cell’s anode? ______
2. What is the charge of an electrolytic cell’s anode? ______
3. What is the charge of a voltaic cell’s cathode? ______
4. What is the charge of an electrolytic cell’s cathode? ______
5. Regardless of the cell, electrons always travel from which electrode? __________
6. Regardless of the cell, oxidation always occurs at the ______________.
7. Regardless of the cell, reduction always occurs at the ______________.
29. Why will the mass of the key increase?
__________________________________________________
30. Label the anode, cathode, and the direction of e-‐flow through the wire.
31. State the difference between voltaic and electrolytic cells in terms of spontaneity.
__________________________________________________
32. State the difference between voltaic and electrolytic cells in terms of energy being released or absorbed.
__________________________________________________
Name: _______________________
20
8. Use the diagram below to answer the following questions:
a. The cathode has hydrogen gas generating from the water. Write the half reaction that shows the hydrogen atom’s oxidation number in water changing to normal H2 gas. Is this reaction oxidation or reduction?
b. The anode has oxygen gas generating from the water. Write the half reaction that shows the oxygen atom’s oxidation number in water changing to normal O2 gas. Is this reaction oxidation or reduction?
c. Why is a battery put in this picture? Does water simply become gaseous H2 and O2 spontaneously?
1. In a voltaic cell, chemical energy is converted to (1) electrical energy, spontaneously (2) electrical energy, nonspontaneously (3) nuclear energy, spontaneously (4) nuclear energy, nonspontaneously
2. A voltaic cell spontaneously converts
(1) electrical energy to chemical energy (2) chemical energy to electrical energy (3) electrical energy to nuclear energy
The picture to the left show what happens when water is exposed to an electric current. The water is literally ZAPPED with electric and decomposes into H2 and O2.
2H2O à 2H2 + O2
Name: _______________________
21
(4) nuclear energy to electrical energy 3. A voltaic cell differs from an electrolytic cell in that in a voltaic cell
(1) energy is produced when the reaction occurs (2) energy is required for the reaction to occur (3) both oxidation and reduction occur (4) neither oxidation nor reduction occurs
4. Which half-‐reaction can occur at the anode in a voltaic cell?
(1) Ni2+ + 2e-‐ → Ni (2) Sn + 2e-‐ → Sn2+ (3) Zn → Zn2+ + 2e-‐ (4) Fe3+ → Fe2+ + e-‐
5. Which process requires an external power source?
(1) neutralization (3) fermentation (2) synthesis (4) electrolysis
6. Which energy transformation occurs when an electrolytic cell is in operation?
(1) chemical energy → electrical energy (2) electrical energy → chemical energy (3) light energy → heat energy (4) light energy → chemical energy
7. What is the purpose of the salt bridge in a voltaic cell?
(1) It blocks the flow of electrons. (2) It blocks the flow of positive and negative ions. (3) It is a path for the flow of electrons. (4) It is a path for the flow of positive and negative ions.
8. Which statement is true for any electrochemical cell?
(1) Oxidation occurs at the anode, only. (2) Reduction occurs at the anode, only. (3) Oxidation occurs at both the anode and the cathode. (4) Reduction occurs at both the anode and the cathode.
9. Given the balanced equation representing a reaction occurring in an electrolytic cell: 2NaCl(l) à 2Na(l) + Cl2(g) Where is Na(l) produced in the cell? (1) at the anode, where oxidation occurs (2) at the anode, where reduction occurs (3) at the cathode, where oxidation occurs (4) at the cathode, where reduction occurs
Name: _______________________
22
Answer questions 10 and 11 using the diagram below, which represents an electrochemical cell. Use the diagram of a key being plated with copper to answer questions 12 through 15.
16. 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. Na: +1 Al: +3 F: –1
(b) Write the balanced half-‐reaction equation for the reduction of Al3+ to Al. Al3+ + 3e– àAl
(c) Explain, in terms of ions, why molten cryolite conducts electricity. There are freely moving ions in the molten c
(d) 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.
Electrolysis uses electrical
10. When the switch is closed, in which half-‐cell does oxidation occur? 11. What occurs when the switch is closed? (1) Zn is reduced. (2) Cu is oxidized. (3) Electrons flow from Cu to Zn. (4) Electrons flow from Zn to Cu.
12. What is the name of the process shown in the diagram? 13. What is the purpose of the battery in this electrolytic cell? 14. Which electrode, A or B, attracts positive copper ions? 15. Given the reduction reaction for this cell: Cu2+(aq) + 2e– → Cu(s) This reduction occurs at
(1) A, which is the anode (3) B, which is the anode (2) A, which is the cathode (4) B, which is the cathode
Name: _______________________
23
energy. Voltaic cells produce electrical energy. 17. Base your answers to the following questions on the diagram of the voltaic cell below.
(a) Identify the anode and the cathode. Anode = Pb Cathode = Ag (b) Write the oxidation and reduction half-‐reactions for this voltaic cell. Red: 2Ag+ + 2 e-‐ à 2Ag 2 e-‐ (c) What is the total number of moles of electrons needed to completely reduce 6 moles of Ag+(aq) ions?
(d) Describe the direction of electron flow between the electrodes. Electrons flow from the Pb electrode to the Ag electrode
(e) State the purpose of the salt bridge in this cell. Maintains a balance of charge; allows ions to migrate
(f) State the electrode to which positive ions migrate when the switch is closed. Ag electrode (the cathode) – to balanced the electrons that are arriving
(g) As this voltaic cell operates, the mass of the Ag(s) electrode increases. Explain, in terms of silver ions and silver atoms, why this increase in mass occurs.
At the Ag electrode, silver ions in the water are being reduced to solid silver atoms. The solid silver deposits on the Ag electrode, which increases its mass. 18. Base your answers to the following questions on the diagram below, which represents a voltaic cell at 298 K and 1 atm.
a) In which half-‐cell will oxidation occur
when switch S is closed? b) Write the balanced half-‐reaction equation that will occur in half-‐cell 1 when switch S is closed. Pb2+ + 2e– → c) Describe the direction of electron flow between the electrodes when switch S is closed.
anode to cathode
Name: _______________________
24
19. Base your answers to the following questions on the information below. 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) State the direction of the flow of electrons between the electrodes in this cell. From Mg (anode) to Fe (cathode) (b) Explain, in terms of reactivity, why magnesium is preferred over zinc to protect underground iron pipes. Your response must include both magnesium and zinc. ore readily than Zn.
20. Base your answers to the following questions 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? 12
(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. The salt bridge allows ions to flow between the half-‐cells. Maintains a balance of charge
21. Base your answers to the following questions on the diagram and balanced equation below, which
represent the electrolysis of molten NaCl.
(a) When the switch is closed, which electrode will attract the sodium ions? cathode, one on the right
(b) What is the purpose of the battery in this electrolytic cell? energy source
(c) Write the balanced half-‐reaction for the reduction that occurs in this electrolytic cell.
Name: _______________________
25
Na+ + e– à Na OR 2 22. State two similarities and two differences between voltaic and electrolytic cells. Electrons flow anode to cathode Diff: voltaic: produces energy, spongy Electrolytic: requires energy, nonspontaneous, 1 cell, electrical energy à chemical energy 23. The apparatus shown in the diagram consists of two inert platinum electrodes immersed in water. A small amount of an electrolyte, H2SO4, must be added to the water for the reaction to take place. The electrodes are
connected to a source that supplies electricity.
(a) What type of electrochemical cell is shown? electrolytic or electrolysis
(b) What particles are provided by the electrolyte that allow an electric current to flow? Ions, charged particles, H3O+, or
SO42–
24. 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.
(a) Identify the information in the diagram that indicates this system is an electrolytic cell. A battery is part of the cell and is providing energy that causes the reaction.
Electricity is used to operate the cell. (b) Determine the change in oxidation number of oxygen during the reaction in test tube 1. -‐2 to 0 (c) Explain, in terms of the products formed in test tube 2, why litmus turns blue in test tube 2. Litmus turns blue when hydroxide ions are produce
Name: _______________________
26
25. Which reaction occurs spontaneously? (1) Cl2(g) + 2NaBr(aq) à Br2(l) + 2NaCl(aq) (2) Cl2(g) + 2NaF(aq) à F2(g) + 2NaCl(aq) (3) I2(s) + 2NaBr(aq) à Br2(l) + 2NaI(aq) (4) I2(s) + 2NaF(aq) à F2(g) + 2NaI(aq) 26. Which metal reacts spontaneously with a solution containing zinc ions? (1) magnesium (3) copper (2) nickel (4) silver
27. Which metal with react with Zn2+
spontaneously, but will not react with Mg2+? (1) Al (3) Ni (2) Cu (4) Ba 28. Which of the following metals has the least tendency to undergo oxidation? (1) Ag (3) Zn (2) Pb (4) Li
29. Because tap water is slightly acidic, water pipes made of iron corrode over time, as shown by the balanced ionic equation: 2Fe(s) + 6H+(aq) à 2Fe3+(aq) + 3H2(g) Explain, in terms of chemical reactivity, why copper pipes are less likely to corrode than iron pipes.
30. Identify one metal that does not react spontaneously with HCl(aq).