Electrochemistry: Voltaic Cells

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Marybeth Hart
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1 Name: Band: Date: An Overview and Review Electrochemistry: Voltaic Cells Electrochemistry is the field of chemistry that focuses on reactions involving electrical energy. All electrochemical reactions are oxidation-reduction (redox) reactions, which should make sense: redox reactions involve moving electrons, and moving electrons are an electric current! You ve already studied oxidation-reduction reactions, and you should be able to assign oxidation numbers, write and balance half-reactions, identify oxidizing and reducing agents, and balance redox reactions. Here are a few reminders about balancing redox reactions from your favorite textbook (Kotz): Hydrogen balance can be achieved only with H + /H 2 O (in acid) or OH - /H 2 O (in base). Never add H or H 2 to balance hydrogen. Use H 2 O or OH - as appropriate to balance oxygen. Never add O atoms, O 2- ions, or O 2 for O balance. Never include H + (aq) and OH - (aq) in the same equation. A solution can be acidic or basic, never both. The number of electrons in a half-reaction reflects the change in oxidation state of the element being oxidized or reduced. Electrons are always a component of half-reactions but should never appear in the overall equation. Include charges in the formulas for ions. Omitting the charge, or writing the charge incorrectly, is one of the most common errors seen on student papers. The best way to become competent in balancing redox equations is to practice, practice, practice. Voltaic/Galvanic Cells One type of electrochemical process involves using a chemical change to produce an electric current. Voltaic (also known as galvanic) cells are devices that do just that: they use product-favored redox reactions to move electrons through a circuit. To make a voltaic cell at home or in the lab, you need to choose two complementary halfreactions (one oxidation and one reduction) that are product-favored. Right now, you ll just have to trust me when I say that the oxidation of copper metal and the reduction of silver ions are both product-favored. 1. What type of reaction is product-favored: a spontaneous reaction or nonspontaneous reaction? 2. If the two half-reactions mentioned above are both product favored: a. Predict the sign (positive or negative) of the change in Gibbs free energy for these half-reactions. b. Predict the value of K for these half-reactions (greater than one or less than one). 3. Write the half-reaction for the oxidation of copper metal involving two electrons.

4. Write the half-reaction for the reduction of silver ions. Once you ve chosen your two half-reactions, you need to find two containers: one for each half-reaction.

2 4. Write the half-reaction for the reduction of silver ions. Once you ve chosen your two half-reactions, you need to find two containers: one for each half-reaction. The half-reactions need to be separated so that when copper metal transfers its electrons to silver ions, the electrons have to go over the river and through the woods before reaching their destination. In other words, you want the electrons to travel through a wire to generate that electric current! Check out the set-up below: Notice the contents of each container: In the container on the left, you have copper metal (the oxidation reactant) and a solution of copper(ii) ions (the oxidation product). In the container on the right, you have silver ions (the reduction reactant) and silver metal (the reduction product). The solid metals are called the electrodes. More specifically, the anode is the electrode at which oxidation occurs, and the cathode is the electrode at which reduction occurs. To help you remember this, consider this memory device: an ox; red cat. Inert (unreactive) electrodes are often used when a gas is involved in a half-reaction OR when ion-to-ion oxidation/reduction occurs (such as from Fe 3+ to Fe 2+ instead of Fe). Examples of inert electrodes include platinum, Pt, or graphite. 5. Why is the copper metal the anode in this voltaic cell? 6. Why is the silver metal the cathode in this voltaic cell? 7. Hydrogen gas is sometimes generated in redox reactions as hydrogen ions are reduced. Identify a possible electrode for this half-reaction.

3 Another important feature of the voltaic cell above is the salt bridge. A salt bridge is basically salty Jello; gelatin is added to a solution of an electrolyte/ionic salt that will be unreactive with the components of the half-reactions. The ions within the salt bridge help prevent either container from becoming too negative or too positive. For example, as copper is oxidized, the copper atoms of the anode go into solution as copper(ii) ions, creating a solution that has a lot of cations not good. To neutralize the situation (haha!), nitrate ions from the salt bridge enter into the container on the left. Similarly, as silver ions gain electrons and attach themselves to the cathode as neutral metal atoms, the solution loses cations and becomes more negative; as a result, sodium ions from the salt bridge travel into the container on the right. For our purposes, ions in the solution do NOT flow back into the salt bridge; ignore any parts of the diagram that suggest this. 8. Which electrode the anode or the cathode will gain mass during the redox reaction in a voltaic cell? Why? 9. Which electrode the anode or the cathode will lose mass during the redox reaction in a voltaic cell? Why? 10. Why are the ionic salts KNO 3 and NaNO 3 always good choices for salt bridges? 11. Anions in the salt bridge always travel to the, and cations from the salt bridge always travel to the. You can assign the anode a negative (-) charge and the cathode a positive (+) charge. This is because electrons are generated at the anode during oxidation and move to the cathode, where there is an electron deficit and therefore a positive charge. Notice the direction of electron flow as a result of these charge differences in the electrodes. Electrons always flow from the anode to the cathode. To help you remember this, think: electrons flow in alphabetical order, from A to C. You can also think fat cat : From the Anode To the CAThode 12. Based on the net reaction, how many silver ions are reduced for everyone one atom of copper that is oxidized? 13. What half-reaction takes place at the negative electrode: oxidation or reduction? 14. What half-reaction takes place at the positive electrode: oxidation or reduction?

4 Practice: Voltaic Cells 15. Describe how to set up a voltaic cell to generate an electric current using the oxidation of iron metal to iron(ii) ions and the reduction of copper(ii) ions to copper metal. Draw a diagram similar to the worked example above. 16. Standard cell notation (aka line notation) is written as follows: anode/anode solution // cathode solution/cathode The / refers to a phase boundary, and the // indicates a salt bridge. Notice how the notation is in alphabetical order! For the example voltaic cell, you d write: Cu/Cu 2+ (1.0 M) // Ag + (1.0 M)/Ag The solution concentrations are always 1.0 M if the cell is at standard conditions (more on this later). Try writing standard cell notation for the cell in question The hydrogen electrode is particularly important in electrochemistry because, as you ll see later, it s used as a point of reference for assigning important electrical values to electrodes. The voltaic cell below uses a hydrogen electrode as the anode. Based on the diagram, describe the voltaic cell in as much detail as possible. Write standard cell notation for this cell, too.

5 Cell Potential Cell potential is a measure of electromotive force, which is just the pull of the electrons as they travel from anode to cathode. We use the symbol Emf, E cell or ε cell and units of volts, V. You ll notice in the diagrams above that all of the cells have a voltmeter attached to the wire connecting the two half-cells, and the copper/silver cell has a lower Emf than the iron/hydrogen cell (+0.46 V vs V). This implies that the electrons generated in the copper/silver cell do not have as strong of a pull as they travel through the wire. Not only does every redox reaction have a cell potential, but half-reactions have cell potentials, too. To find the potential of a half-reaction, the electromotive force of the half-reaction is compared to the electromotive force of a standard hydrogen electrode; by assigning the hydrogen electrode a value of zero volts, you can calculate the potential of the half-reaction. A list of half-reactions and their potentials are in your Reference Tables. Find that list right now, and use it to answer the questions below. 18. What does the naught symbol after E mean: a. In terms of temperature? b. In terms of pressure (for gases)? c. In terms of molarity (for solutions)? 19. What half-reaction has a potential of 0.00 V? 20. What half-reaction has a potential of 1.50 V? 21. What half-reaction has a potential of V? 22. What type of reaction is represented by every equation on the sheet: oxidation or reduction? How can you tell? Standard Reduction Potentials Now that you re familiar with the general structure of that reference sheet, let s figure out what all those numbers mean. First, substances that have the most positive reduction potentials are most easily reduced. Secondly, substances that have the least positive (most negative) reduction potentials are most easily oxidized. 23. What types of elements metals or nonmetals are most easily reduced? 24. What types of elements metals or nonmetals are most easily oxidized?

6 This table of reduction potentials can also be used to determine the strength of various oxidizing and reducing agents. Substances that are high up on the table (easily reduced) make the best oxidizing agents, and substances that are low on the table (easily oxidized) make the best reducing agents. Finally, let s connect this information to those voltaic cells: if you are given two half-reactions for use in a voltaic cell, you can use this table to figure out which half-reaction takes place at the anode (oxidation) and which half-reaction takes place at the cathode (reduction). 25. A student wants to construct a battery using sodium metal and an aqueous solution of sodium ions as well as zinc metal and an aqueous solution of zinc ions. a. Write the reduction half-reaction that will take place in this battery. b. Write the oxidation half-reaction that will take place in this battery. Hint: you ll need to flip the equation on the table in order to show oxidation! c. Write the balanced redox reaction for this battery. d. What is the oxidizing agent in this redox reaction? e. What is the reducing agent in this redox reaction? Calculating Standard Cell Potential A popular topic on the AP exam is calculating standard cell potentials, E cell. Since you don t have a voltmeter, you ll have to do some math. Just follow these steps: --Decide which substance is oxidized and which substance is reduced using the table of reduction potentials. --Write both equations EXACTLY AS THEY APPEAR ON THE TABLE. --Reverse the equation that will be oxidized and change the sign of the voltage. This E value is now E ox. --Balance the two half-reactions using coefficients/electrons but do NOT multiply E values at any stage of this process. --Add the two half-reactions and the voltages together. Remember: E cell = E ox + E red 26. Consider a galvanic cell based on the reaction: Al 3+ (aq) + Mg(s) Al(s) + Mg 2+ (aq) Write a balanced cell reaction and calculate the E for the cell. 27. Calculate the cell voltage for the following reaction: Fe 3+ (aq) + Cu(s) Cu 2+ (aq) + Fe 2+ (aq) Write a balanced cell reaction and calculate the E for the cell.

AP CHEMISTRY NOTES 12-1 ELECTROCHEMISTRY: ELECTROCHEMICAL CELLS

AP CHEMISTRY NOTES 12-1 ELECTROCHEMISTRY: ELECTROCHEMICAL CELLS Review: OXIDATION-REDUCTION REACTIONS the changes that occur when electrons are transferred between reactants (also known as a redox reaction)