ii. Gains Oxygen and/or loses hydrogen (hydrocarbons): Example:

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1 Identify element being oxidized and element being reduced 1. Oxidation: Increasing oxidation number 2. Reduction: Decreasing oxidation number 3. Hierarchical guidelines 4. OXIDIZED i. Free elements have an oxidation # of zero (e.g, Fe, H 2, S 8) ii. The oxidation number must add up to the overall charge of the molecule or ion iii. Assign oxidation numbers to all but one of the elements using the hierarchical guidelines. Find the oxidation number of the final element mathematically. 1. Alkali metal ions (Group 1A) always have an oxidation number of +1. Alkaline earth metals (Group 2A) are assigned an oxidation number of Fluorine, in a compound, is assigned Hydrogen is usually +1 (less common oxidation number is -1) 4. Oxygen is usually assigned -2 (less common numbers include -1, peroxide) i. Charge increases: Example: ii. Gains Oxygen and/or loses hydrogen (hydrocarbons): Example: 5. REDUCED i. Charge decreases: Example ii. Gains hydrogen and/or loses oxygen (hydrocarbons) 6. OXIDIZING AGENT: 7. REDUCING AGENT: Oxidation and Reduction: Balancing using half reaction method - 1 -

2 Balance RedOx reactions by the half-reaction method 1. Assign oxidation numbers to all atoms, Identify what is being oxidized and what is being reduced. 2. Determine if the reaction is in acid or basic conditions (one extra step is required if the reaction occurs in basic conditions) 3. Split the reaction into half-reactions. a. Oxidation half reaction b. Reduction half reaction 4. Balance, in order a. All atoms except H and O b. Balance the O atoms by adding H 2O to the side of the equation that needs O c. Balance H by adding H + to the side of the equation that needs H. d. Balance the electric charge by adding electrons (e - ) to the more positive side 5. Combine the two half-reactions a. Multiply each half-reaction by a factor so that each has the same number of electrons. b. Add the two half-reactions (the electrons should cancel) c. Simplify the reaction by canceling species that occur on both side of the reaction d. Check that the reaction is balanced 6. If the reaction occurs in acidic conditions, you are finished. If it occurs in basic conditions, finish by a. Adding one OH - to both sides of the reaction for each H + (or H 3O + ) b. When H + and OH - occur on the same side, combine them to form H 2O (or H 3O + + OH - makes two waters) c. Cancel water molecules that occur on both sides. Oxidation and Reduction: Balancing using half reaction method - 2 -

3 Example 1 : In an acidified solution, Dichromate ion (Cr 2O 2-7 ) reacts with zinc metal. Products of the reaction include zinc ions (Zn 2+ ) and chromium(iii) ions (Cr 3+ ) Write the balanced ionic equation for this reaction using the half reaction method. 1. Determine what is being oxidized and what is being reduced a. Assign oxidation numbers Cr 2 O 7 2- Zn Zn 2+ Cr 3+ b. Element Reduced:, element oxidized: Oxidizing Agent: Reducing Agent 2. Reaction conditions: Acid or basic? 3. Oxidation half reaction: Reduction half reaction: 4. Balance each half reaction a. All atoms except H and O b. Balance O (add H 2O) c. Balance the H (add H + ) d. Balance the charge (e-) 5. Combine the two half reactions. a. Multiply (electrons must be equal) b. Add c. Simplify d. check 6. Adjust for basic conditions if necessary Oxidation and Reduction: Balancing using half reaction method - 3 -

4 Example 2: Write a balanced net ionic reaction: In a basic solution, Lead(II) ion (Pb 2+ ) react with hypochlorite ion (ClO-) to give lead(iv) oxide (PbO 2) and chloride ion (Cl - ). More examples: Acidic: Zn(s) + - (aq) Zn 2+ (aq) + NH 4+ (aq) Basic: H 2O 2 + ClO 2 ClO O 2 Oxidation and Reduction: Balancing using half reaction method - 4 -

5 VOLTAIC CELLS (aka galvanic cell) and Batteries I. A BATTERY 1. A package of one or more voltaic cells connected together to produce electric energy II. III. 2. Word origin: A set or series of similar units as in a number of pieces of artillery used together Voltaic Cells: Key ideas 1. Spontaneous reaction (ΔG<0) is the source of energy 2. Two half reactions are separated physically from each other 3. Connections are made two ways i. Connection to the device needing power ( connection) ii. Salt Bridge. Allows for to flow. ( connection) An atomic view of a voltaic cell: IV. ACTIVITY Batteries : How does a voltaic cell work. Objectives 1. Identify the anode and the cathode in a drawing given either The direction of the electron flow OR The half-cell reactions OR The activity series of metals 2. Describe the purpose of the salt bridge in a voltaic cell Voltaic Cell (Galvanic Cell) - 5 -

6 Why? Batteries How does a battery (voltaic cell) work? When we use portable devices like MP3 players and cell phones we need a ready source of electricity to provide a flow of electrons. Batteries are the common solution to this challenge. In a battery or voltaic cell, oxidation and reduction reactions provide electrons which power our devices. In this activity we will explore the chemistry of voltaic cells or batteries. Model 1 Voltaic Cell Bulb not lit Bulb lit Wire not connected Zn Salt Bridge K + Cl Cl Cu Wire connected Zn Salt Bridge Cu K + Anode K + Cl Zn 2+ Cu 2+ Zn 2+ Cu 2+ Cathode Time Passes Cl K + Zn 2+ Cl Zn 2+ Cl Zn 2+ K + Cu 2+ NO 3 K + Before the wire is connected A few minutes after the wire was connected 1. Consider the reaction in Model 1. Notice that there are two zinc ions (Zn 2+ ) in the beaker on the left before the wire is connected. Explain why the number of nitrate ions ( ) present in that beaker is correct. 2. Examine the system in Model 1 both before the wire is connected and after it is connected. Identify two specific pieces of evidence that a chemical reaction has occurred as time passes with the wire connected. Batteries 1-6 -

7 3. Examine the diagram in Model 1. a. Which piece of solid metal loses mass (gets smaller) as the reaction proceeds? b. Does the number of zinc ions (Zn 2+ ) in solution increase or decrease as the reaction proceeds? c. Circle the half-reaction below that represents the change in the metal identified in part a. Zn 2+ (aq) + 2e Zn(s) Zn(s) Zn 2+ (aq) + 2e Cu 2+ (aq) + 2e Cu(s) Cu(s) Cu 2+ (aq) + 2e d. Is the reaction circled in part c an oxidation or reduction reaction? 4. Examine the diagram in Model 1. a. Which piece of solid metal is gaining mass as the reaction proceeds? b. Where do those metal atoms come from? Explain. c. Circle the half-reaction below that represents the change in the metal identified in part a. Zn 2+ (aq) + 2e Zn(s) Zn(s) Zn 2+ (aq) + 2e Cu 2+ (aq) + 2e Cu(s) Cu(s) Cu 2+ (aq) + 2e d. Is the reaction circled in part c an oxidation or reduction reaction? 5. Electricity is the flow of electrons. Look back at Model 1. In which diagram can the electrons flow through the wire, the one when the bulb is not lit or the one when the bulb is lit? (Circle one.) Explain your answer. 6. Based on your answers in Questions 3 and 4, which piece of solid metal is giving up electrons, and therefore losing them into the wire? 7. On the drawing of the voltaic cell in Model 1, draw an arrow to depict the direction that the electrons are traveling through the wire. 2 POGIL Activities for High School Chemistry - 7 -

8 8. Consider the reactions occurring in Model 1. a. What type of half-reaction (oxidation or reduction) is occurring at the piece of metal labeled anode in Model 1? b. What type of half-reaction (oxidation or reduction) is occurring at the piece of metal labeled cathode in Model 1? c. Explain how the phrase an ox and a red cat can help students remember the type of halfreaction that occurs at each electrode in an electrochemical cell. 9. Explain how the direction of electron flow in a voltaic cell is consistent with what you would predict from the activity series? 10. Draw an unconnected voltaic cell similar to the one on the left side of Model 1 using iron and silver as electrodes. The solutions should include silver ions (Ag + ), iron ions (Fe 3+ ), and nitrate ions ( ). Use a salt bridge identical to that in Model Use the activity series of metals to determine which metal in the voltaic cell in Question 10 should be the anode. Explain your choice. Batteries 3-8 -

9 12. Draw an arrow on your diagram in Question 10 to indicate the direction of electron flow through the voltaic cell once the wire is connected. 13. Use Model 1 to complete the table. Initially After a Few Minutes Number of zinc ions (Zn 2+ ) dissolved in solution Number of copper ions (Cu 2+ ) dissolved in solution Number of nitrate ions ( ) dissolved in solution Number of potassium ions (K + ) dissolved in solution Number of chloride ions (Cl ) dissolved in solution 14. Even though ions and electrons move around in a voltaic cell, the cell must stay electrically neutral. a. Explain how the anode half-cell in Model 1 remains electrically neutral (no charge) even though zinc ions are being formed from neutral zinc metal. Refer to the table in Question 13 to support your answer. b. Explain how the cathode half-cell in Model 1 remains electrically neutral (no charge) even though copper ions are being removed from the solution. Refer to the table in Question 13 to support your answer. c. What is the role of the salt-bridge in a voltaic cell? 4 POGIL Activities for High School Chemistry - 9 -

10 Extension Questions 15. Work with your group to apply everything you have learned about batteries to label the voltaic cell diagram below. Use the path and direction that the electrons are traveling to help you. a. Label the following items on the diagram. Site of oxidation Site of reduction Anode Cathode b. Complete the few minutes after drawing to show what ions would be in the beakers and the salt bridge, and how the electrodes may have changed. c. Write the oxidation and reduction half-reactions for the voltaic cell. Wire not connected Fe Salt Bridge K + Cl Cl Zn Fe Salt Bridge Zn K + Cl K + Fe 2+ Zn 2+ Fe 2+ Zn 2+ Time Passes Before the wire is connected A few minutes after the wire was connected 16. Propose a reason why it is necessary to place the metal electrodes in solutions of their own ions to make a battery. Batteries

11 Standard Electrode (Reduction) Potentials Which direction will the electrons flow? Voltaic Cell Notation The oxidation half cell, the, is written on the. The reduction half cell, the, is written on the. Species oxidized I Oxidation product II Species reduced I Reduction product INTERPRET the following Zn(s) Zn 2+ (aq) Cu 2+ (aq) Cu(s) Write the reaction for the voltaic cell that is represented by the following : Al(s) Al 3+ (aq) Cu 2+ (aq) Cu(s) Sketch the voltaic cell represented by the above reaction. Label the anode and cathode. Show the directions the electrons flow (external connection) and the direction the ions flow (internal connection) Voltaic Cell (Galvanic Cell)

12 Cell Potential (EMF: Electromotive Force) MOVEMENT OF ELECTRONS THROUGH THE EXTERNAL CIRCUIT: Just as work is required to pump water from one point to another, so work is required to move electrons. Water flows from areas of high pressure to areas of low pressure. Similarly, electrons flow from high electric potential to low electric potential. Electric potential can be thought of as electric pressure. Spontaneous flow of electrons is due to the diff in potential energy between two substance) Cell potential (E cell) is measured in VOLTS (Joule/Coulomb) Summary E o cell ΔG Q & K Relationship Reaction Direction Spontaneity (as written) The standard electrode potential, E, is the electrode potential when the concentrations of solutes are 1 M, the gas pressures are 1 atm, and the temperature has a specified value (usually 25 C). The superscript degree sign ( ) signifies standard state conditions. Cell Potential (EMF)

13 Summary of electrochemistry units and constants Joule (J): Unit of work Work = -nfe cell (where n is the number of moles of electrons involved in either half-cell reaction) Coulomb (C) : Unit of charge 1 e - = X C Faraday Constant (F): The charge of one mole of electrons 1 F = Coulombs 1 mole of e Volt (V): Energy absorbed (or evolved) when a charge moves through an energy difference Watt: Unit of Power Ampere: Unit of current 1V = J C 1 W= J s 1 A = C S The emf of a particular cell is V. The cell reaction is 2Al(s) + 3Cu 2+ (aq) 2Al 3+ (aq) + 3Cu(s) Calculate the maximum electrical work (J) of this cell. Hint: To determine the value of n we need to examine the half reactions: Electrochemistry units: Work done by a cell

14 Calculating Ecell from half reactions E cell = E oxidation + E reduction Look back at the electrode potentials given in the Table Are the electrode potentials given for oxidation or reduction? How can you determine the potentials for oxidation? Example: How does the reduction potential of Li + (aq) compare to the oxidation potential of Li(s)? Write each reaction and its cell potential. Voltage is an Intensive Property This means that if you double the coefficents for the reaction the value of the standard potential. Give the standard potentials for the following reactions: Al(s) Al 3+ (aq) +3e - 2Al(s) 2Al 3+ (aq) +6e - Sn 2+ (aq) +2e- Sn(s) 3Sn 2+ (aq) +6e- 3Sn(s) Calculating E 0 cell

15 Let s find Eºcell for the following cell: Al(s) Al 3+ (aq) Cu 2+ (aq) Cu(s) A fuel cell is simply a voltaic cell that uses a continuous supply of electrode materials to provide a continuous supply of electrical energy. A fuel cell employed by NASA on spacecraft uses hydrogen and oxygen under basic conditions to produce electricity. The water produced in this way can be used for drinking. The net reaction is 2H2(g) + O2(g) 2H2O(g) Calculate the standard emf of the oxygenhydrogen fuel cell. 2H 2O(l) + 2e H 2(g) + 2OH (aq) O 2(g) + 2H 2O(l) + 4e 4OH (aq) E = 0.83 V E = 0.40 V Calculating E 0 cell

16 Predicting the Direction of a Reaction: Standard Cell Potential (Eº) to Standard Free Energy (ΔGº) E o cell ΔG Q & K Relationship Reaction Direction Spontaneity (as written) Will dichromate ion oxidize manganese(ii) ion to permanganate ion in acid solution under standard conditions? Standard potentials Cr 2O H e - Cr H 2O MnO 4 + 8H e - Mn H 2O E = 1.33 V E = 1.49 V The free energy change, DG, for a reaction equals the maximum useful work of the reaction. ΔG = wmax = nfe From the cell potential, calculate the standard free energy change for the net reaction used in the hydrogenoxygen fuel cell: 2H2(g) + O2(g) 2H2O(l) The cell potential is 1.23 V. First we need determine the number of electrons involved in the reaction. How does this compare with ΔGf of H2O(l)? (ΔGf = kj/mol) Calculating E 0 cell

17 A voltaic cell consists of one half cell with Fe dipping into an aqueous solution of 1.0 M FeCl2 and the other half cell with Ag dipping into an aqueous solution of 1.0 M AgNO3. Fe 2+ (aq) I Fe(s) Ag + (aq) I Ag(s) E = 0.41 V E = 0.80 V Determine the reaction and calculate the cell potential Continued with same reaction: Calculate ΔGrxn from ΔGf (ΔGf Ag + (aq) = 77 kj/mol, and Fe 2+ (aq)= 85 kj/mol.) Continued: Verify the relationship: ΔG =-nfe Equilibrium Constants from Cell Potentials nfe = RT ln K Rearrange, solve for E Calculating E 0 cell

18 Calculate the equilibrium constant K at 25 C for the following reaction for the standard cell potential: Pb 2+ (aq) + Fe(s) Pb(s) + Fe 2+ (aq) Find the cell potential Now find K 13 P a g e

19 Dependence of Cell Potential on Concentration: Nernst Equation ΔG = ΔG + RT ln Q o E cell Ecell RT nf lnq This reaction is sometimes written as o E cell Ecell lnq n Consider a voltaic cell, Fe(s) Fe 2+ (aq) Cu 2+ (aq) Cu(s), being run under standard conditions. a. Is ΔG positive or negative for this process? b. Change the concentrations from their standard values in such a way that Ecell is reduced. Write your answer using the shorthand notation. What is the cell potential of the following voltaic cell at 25ºC? Zn(s) Zn 2+ (0.200 M) Ag + ( M) Ag (s) 14 P a g e

20 Determination of ph (ion selective electrodes) Standard Hydrogen electrode: Glass electrode Test Solution Ag(s) Ag + (1 M) H + (test solution) H 2(1 atm) Pt What are the reactions? Anode: Cathode: Overall: Nernst Equation: Expression for Q: How is ph determined? A ph meter is constructed using hydrogen gas bubbling over an inert platinum electrode (the hydrogen electrode) at a pressure of 1.2 atm. The other electrode is aluminum metal immersed in a 0.20 M Al 3+ solution. What is the cell potential when the hydrogen electrode is immersed in a sample of acid rain with a ph of 4.0 at 25 C? Al(s) Al 3+ (0.20 M) H + (test solution) H 2(1.2 atm) Pt Corrosion Control: Cathodic Protection P a g e

21 Common Voltaic Commercial Voltaic Cell Zinc-Carbon Dry Cell (Leclanché ) Anode: Zn(s) Zn 2+ (aq) + 2e Cathode:2NH4 + (aq) + 2MnO2(s) + 2e Mn2O3(s) + H2O(l) + 2NH3(aq) Alkaline Battery By Túrelio (via Wikimedia-Commons), 2009 /, CC BY-SA 3.0 de, Zn (s) + 2OH (aq) ZnO (s) + H 2O (l) + 2e [E = V] 2MnO 2(s) + H 2O (l) + 2e Mn 2O 3(s) + 2OH (aq) [E = V] Overall reaction: Zn (s) + 2MnO 2(s) ZnO (s) + Mn 2O 3(s) [e = V] Recharable NiCD battery Corrosion Control: Cathodic Protection P a g e

22 During Discharge: E oxº= The reactions at the nickel oxide electrode are: E redº= The net reaction during discharge is Lead Storage Cell: Anode: Cathode: Pb(s) + HSO4 (aq) PbSO4(s) + H + (aq) + 2e PbO2(s) + 3H + (aq) + HSO4 (aq) + 2e PbSO4(s) + 2H2O(l) Corrosion Control: Cathodic Protection P a g e

23 Cathodic Protection: Corrosion Control: Rusting occurs when iron comes in contact with iron: Anode: Cathode: Fe(s) Fe 2+ (aq) + 2e O 2(g) + 2H 2O(l) + 4e 4OH (aq) Sacrificial Anode: A more active metal such as Zn or Mg, (More easily oxidized) is attached as protection. Examples: Compare oxidation potentials of Fe and Zn Zn (s) Zn 2+ Fe(s) Fe 2+ Applications of Cathodic Protection: Galvanized Nails Hull of Ship Water Heaters Pipelines Corrosion Control: Cathodic Protection P a g e

24 Electrolytic Cell An electrolytic cell is an electrochemical cell in which electric current an otherwise reaction. This process is called. (See Electrolytic Cells ) Corrosion Control: Cathodic Protection P a g e

25 Name Chem 163 Section: Team Number: ALE 27. Electrolytic Cells (Reference: 21.7 Silberberg 5 th edition) If an element doesn t occur naturally, where does a sample of it come from? The Model: A Battery Used as an Electron Pump If a battery is used to pump electrons from one electrode to a second electrode, it is possible to force a non-spontaneous reaction to occur. Let us at the moment consider an electrolytic cell that has inert electrodes. An inert electrode is made of some electrically-conducting substance, which is resistive to being oxidized or reduced. Graphite or an unreactive metal, like gold or platinum, may be used. In the cell to the right, the electrodes are submerged in an aqueous solution of some soluble salt, a strong electrolyte. (While the salt is represented generically as MX, nothing is implied about the charges of the cation and anion their charges do not have to be 1:1.) Key Questions 1. As the battery pumps electrons from the electrode on the left to the electrode on the right, what charge begins to accumulate on each electrode? (Circle your choices.) Electrode on left: positive or negative Electrode on right: positive or negative 2. What species flows to the electrode on the left? to the electrode on the right? Hints: Do like charges attract or repel? Do opposite charges attract or repel? (Circle your choices.) Electrode on left: anions or cations Electrode on right: anions or cations 3. In each of the following half-reactions, are electrons reactants or products? Oxidation: e - is reactant or product Reduction: e - is reactant or product 4. In the electrolytic cell, which species are likely to be oxidized? to be reduced? (Circle all choices that apply.) Species likely to collide with anode and be oxidized: anions cations water Species likely to collide with cathode and be reduced: anions cations water Page 1 of 7

26 5 a. Now consider the electrolysis of Na 2 SO 4 (aq) using inert electrodes exposed to the air. At the anode, either sulfate or water may oxidize: 2 SO 2-4 (aq) S 2 O 2-8 (aq) + 2 e - o E V 2 H 2 O(l) O 2 (g) + 4 H + (aq) + 4 e - o E V Which species would you predict is the one that is oxidized at the anode? Briefly explain why. ox ox b. When Na 2 SO 4 (aq) is electrolyzed, either sodium cations or water may be reduced: Na + (aq) + e - o Na(s) E V 2 H 2 O(l) + 2 e - H 2 (g) + 2 OH - o (aq) E V Which species would you predict is the one that is reduced at the cathode? Briefly explain why. red red c. Suppose Na 2 SO 4 (aq) is electrolyzed using inert electrodes which are exposed to the air. On the following schematic of the electrolytic cell i. draw what is being consumed and what is being produced at each of the electrodes. ii. draw arrows indicating the flow of each reactant and product. d. Why cannot pure water be electrolyzed? (i.e., Why must sodium sulfate or some similar electrolyte be dissolved in the water? Hint: Consider the flow of ions in the aqueous solution between the electrodes.) Page 2 of 7

27 6. Let us now consider the electrolysis of CuSO 4 (aq) using inert electrodes exposed to the air. The possible processes at the electrodes are: at the anode 2 SO 2-4 (aq) S 2 O 2-8 (aq) + 2 e - o E V 2 H 2 O(l) O 2 (g) + 4 H + (aq) + 4 e - o E V ox ox Cu 2+ (aq) + 2 e - Cu(s) at the cathode E 0.34 V o red 2 H 2 O(l) + 2 e - H 2 (g) + 2 OH - o (aq) E V a. Circle the half-reaction that occurs at the anode and the reaction at the cathode. b. Write the overall cell reaction and determine the absolute minimum voltage that the battery must have to force this nonspontaneous reaction to occur. (I ve said absolute minimum, because ordinarily an additional voltage known as an overvoltage (a.k.a. over potential or junction potential) must be supplied to overcome the slow reaction kinetics due to the high activation energy required for gases to form at the electrode overvoltages vary, but are about 0.4 to 0.6 V for O 2 and H 2. Overall reaction: o Minimum E = cell o o c. If overvoltage is considered, what s the approximate E? E = cell The Model: Electrolysis of Molten Salt If we want to avoid the possibility of more than one half-reaction taking place at each electrode, the best way to do that is to limit which species can interact with the anode and cathode. Suppose instead of an aqueous solution of a salt (again, generically represented as MX ) we electrolyze the molten salt [i.e., MX(l) ]. Then it is possible that both the cations of the salt would be reduced to the elemental state at the cathode and the anions of the salt would be oxidized to the elemental state at the anode. Key Questions 7. Why must a salt be in the liquid state if it is to be electrolyzed? cell red 8 a. What is an advantage of electrolyzing a molten salt over electrolyzing an aqueous solution of that salt? b. What is a procedural challenge to electrolyzing a molten salt? Page 3 of 7

28 9. Elemental magnesium does not occur freely in Nature, but Mg 2+ is the third most abundant ion in sea water. At a concentration of M, there is a virtually inexhaustible supply of magnesium in our oceans! Once magnesium chloride salt is isolated from sea water, it can be electrolyzed to yield magnesium metal. The light-weight magnesium metal is moderately strong, so it is a useful metal for construction (e.g., of airplanes). The combustion of magnesium metal produces an intense light, so it is also used in pyrotechnics. a. Write the oxidation and reduction half-reactions that occur at the anode and cathode, respectively, when MgCl 2 (l) is electrolyzed using inert electrodes. Half reaction at anode: Half reaction at cathode: b. Elemental magnesium metal, Mg, and chlorine gas, Cl 2, react explosively with each other! So when MgCl 2 (l) is electrolyzed, the products must be prevented from recombining. Use the design of the electrolytic cell to the right and the following physical data to explain how these elements may be collected when molten magnesium chloride, MgCl 2 (l), is electrolyzed using inert electrodes. T f ( C) d (g/ml) Mg MgCl Page 4 of 7

29 The Model: Electroplating Electroplating is the process of using electrical current to reduce cations of a desired material from a solution and coat a conductive object with a thin layer of the material (e.g. such as a metal). Electroplating is primarily used for depositing a layer of material to give it a desired property (e.g., abrasion and wear resistance, corrosion protection, aesthetic qualities, etc.) to a surface that otherwise lacks that property. Another application uses electroplating to build up thickness on undersized parts as seen in the electrolytic cell to the right where an external power source is used to increase the thickness of the tin cathode. Key Questions 10. The possible processes at the electrodes are: Sn 2+ (aq) + 2e - Sn(s) at the cathode o Ered V at the anode Cu(s) Cu 2+ (aq) + 2 e - E V o ox 2 H 2 O(l) + 2 e - H 2 (g) + 2 OH - o (aq) E V red 2 H 2 O(l) O 2 (g) + 4 H + (aq) + 4 e - o E V a. Circle the half-reaction that occurs at the cathode and the half-reaction at the anode if the voltage is kept low. (Water will also react if the voltage is high enough.) b. Write the overall cell reaction and determine minimum voltage that the battery must have to force this nonspontaneous reaction to occur. (Do not consider overvoltage since gases are not involved.) Overall reaction: Minimum E = o cell c. If water were to react at each electrode, what is the minimum voltage the battery must have for water to under go hydrolysis and produce hydrogen gas at the cathode and oxygen gas at the anode? (Since gases are involved remember to consider overvoltage) What is the net reaction involving water at each electrode? Overall reaction: Minimum E for the hydrolysis of water = o cell d. What is the minimum voltage (remember to consider overvoltage) the battery must have for both hydrogen gas to form (but no O 2 at the anode) and tin to plate on the cathode? ox Page 5 of 7

30 The Model: Calculating the Mass of Metal Electroplated A common method of electroplating is achieved using the object to be electroplated as the cathode (e.g. the spoon in the figure below) and use as the anode the metal to be plated (e.g. silver) both submerged in an aqueous solution containing the ions to be plated (e.g. Ag(aq) ). In the figure below, silver anode is oxidized (Ag (s) Ag + (aq) + e - ) and at the cathode Ag + is reduced (Ag + (aq) + e - Ag (s) ). Thus there is no net reaction. The overall process consists of simply moving silver atoms from one electrode to the other. Since this is not a spontaneous process, an electrical current is needed to drive the process. The mass in grams of silver electroplated on the spoon is calculated using the following relationships: The approximate mass of silver electroplated can be calculated from the electrical current in amps and the amount of time the current ran: Current: 7.50 amps at a minimum of 0.80 V Time: 30.0 minutes Useful Info. Needed to Calculate the Mass of Ag Electroplated: 1 ampere (A) = 1 Coulomb per sec: 1 A = 1 C / s Charge of 1 mol of electrons =96,485 C = 1 Faraday 1 F = 96,485 C / 1 mol e- Ag + (aq) + e - Ag (s) E o red 0.80 V o Hence, 1 mol of electrons will plate 1 mol of Ag Molar Mass of Ag: g/mol Key Question 11. If the electrolytic cell above were run at 7.50 A for 30.0 minutes, use dimensional analysis to show that the maximum mass of silver that can be electroplated is 15.1 g. 12. How long will it take (in hours) to electroplate 40.0 g of Chromium onto a small automotive accessory from an aqueous Cr( ) 3 solution using a current of 12 amps? Cr = 51.0 g/mol. Use dimensional analysis to show your work Page 6 of 7

31 Exercises 13. Zinc plating (galvanizing) is an important means of corrosion protection. Although the process is done customarily by dipping the object into molten zinc, the metal can also be electroplated from an aqueous Zn 2+ solution. How many grams of zinc can be deposited on a steel tank from a ZnSO 4 solution when a amp current flows for 2.00 days? Use dimensional analysis to show your work. 14. Electrolysis of molten NaCl is the major means of producing sodium metal. If 215 g of Na metal forms Use dimensional analysis to show your work and circle your answers. a. How many moles of electrons are required? b. How many coulombs are required? c. How many amps are required to produce this amount in 9.5 hr? 15. In the electrolysis of a molten mixture of CsBr and SrCl 2, use your knowledge of periodic trends in ionization energy and electron affinity to identify the product that forms at the negative electrode (cathode) and at the positive electrode (anode). Explain your reasoning. Substance formed at the cathode: Reasoning: Substance formed at the anode: Reasoning: Page 7 of 7

32 Voltaic Cells 1. Use the words anode and cathode to complete the sentences Oxidations occurs at the and reduction occurs and the Electrons flow from the to the Cations (from the salt bridge) flow to the cell and anions flow to the The electrode at the is positive and the electrode at the is negative If both electrodes are active metals, the mass of the increases but the mass of the decreases The concentration of metal ion at the increases and the concentration of metal ion at the decreases 2. Practice a labeling exercise: Study Guide/Practice Questions (Midterm #3) P a g e

33 3. Given a cell reaction, identify the reactions that happen at the cathode and anode. Example: The reaction for a lead storage battery is Pb(s) + PbO2(s) + 2 H + (aq) + 2 HSO4 - (aq) 2 PbSO4(s) + 2 H2O(l) What is the cathode and anode reactions? What is oxidized, what is reduced? 4. Given two reduction potentials, decide what reaction occurs at the anode and cathode, and calculate the overall potential of the cell. Also calculate ΔGº and the equilibrium constant K Ni 2+ (aq) + 2e - Ni(s) Eº red = -0.23V Fe 3+ (aq) + 3e - Fe(s) Eº red = -0.04V 5. Interpret short hand cell notation. Translate the following notations Fe(s) I Fe 2+ (aq) I I Ag + (aq) I Ag(s) Pt I H 2(g) (1atm) I H + (aq) (0.05 M) I Br 2(l) I Br - (aq)(2.5m) I Pt 6. Non- standard condition cells Calculate the value of Q for the cell: Pt I H 2(g) (1atm) I H + (aq) (0.05 M) II Br 2(l) I Br - (aq)(2.5m) I Pt Calculate the potential for the above cell at 25 ºC What is the nickel(ii) ion concentration in the voltaic cell if the cell potential is 0.34 V at 25 ºC Zn(s) I Zn 2+ (aq)(1.00m) I I Ni2 + (aq) (? M)I Ag(s) Electrolytic Cell 7. Use the words anode and cathode to complete the sentences Oxidations occurs at the and reduction occurs and the Electrons flow from the to the Cations (from the salt bridge) flow to the cell and anions flow to the The electrode at the is positive and the electrode at the is negative If both electrodes are active metals, the mass of the increases but the mass of the decreases Study Guide/Practice Questions (Midterm #3) P a g e

34 8. Electrolysis of water Decide what is likely to happen during electrolysis of aqueous solutions. You must consider that water might undergo oxidation or reduction. Stoichiometry Of Electrolysis What is the balanced half reaction for the reduction of water? What is the balanced half reaction for the oxidation of water? What are the anode and cathode reactions for the electrolysis of NaCl(l)? What are the anode and cathode reactions for the electrolysis of NaCl(aq)? 9. Important relationships to know What is a Volt? What is an Ampere? What is a Faraday? 10. A constant electric current deposits 365 mg of copper in 216 min from an aqueous copper(ii) nitrate solution. What is the current? Spontaneous vs Non-spontaneous 11. Know the meaning of the sign and or magnitude of E, ΔG, and K. Study Guide/Practice Questions (Midterm #3) P a g e