17.1 Redox Reactions. Oxidation Numbers. Assigning Oxidation Numbers. Redox Reactions. Ch. 17: Electrochemistry 12/14/2017. Creative Commons License

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Ch. 17: Electrochemistry Electric vehicles contain batteries that can be recharged, thereby using electric energy to bring about a chemical change and vice versa.

1 Ch. 17: Electrochemistry Electric vehicles contain batteries that can be recharged, thereby using electric energy to bring about a chemical change and vice versa. (credit: modification of work by Robert Couse-Baker) Electrodeposition of metals: letter&utm_campaign=cen Chemistry: OpenStax 1 Creative Commons License Images and tables in this file have been used from the following sources: OpenStax: Creative Commons Attribution License 4.0. ChemWiki (CC BY-NC-SA 3.0): Unless otherwise noted, the StatWiki is licensed under a Creative Commons Attribution- Noncommercial-Share Alike 3.0 United States License. Permissions beyond the scope of this license may be available at copyright@ucdavis.edu. Principles of General Chemistry (CC BY-NC-SA 3.0): Redox Reactions Electrochemistry is the study of batteries (Galvanic Cells or voltaic cells): using a chemical reaction to generate electricity. Reactions in batteries are redox (oxidation-reduction) reactions in which one element gains electrons and another loses electrons. These two processes MUST happen together. Assign oxidation numbers to determine what is oxidized and what is reduced; oxidizing agent and reducing agent. Review Section 4.2 for assigning oxidation numbers; determining what is oxidized/reduced/oxidizing agent/reducing agent; activity series. Electrodeposition of metals: wsletter&utm_medium=newsletter&utm_campaign=cen 3 Oxidation Numbers 1) Elements in their natural states are 0. 2) Elements in binary ionic compounds are the same as their charges. (e.g., MgCl 2 ) 3) H: usually +1, except with alkali metals (LiH, NaH, etc.) Examples: H 2 O, H 2 (g), H 2 SO 4 4) O: usually -2, except in peroxides (H 2 O 2, K 2 O 2 ) Examples: H 2 O, KMnO 4, Na 2 O 5) The non-oxygen element in a polyatomic ion has to be determined from the other oxidation numbers. Examples: NO 3-, MnO 4-, Cr 2 O Oxidation: loss of electrons Reduction: gain of electrons LEO the lion goes GER Redox Reactions The substance oxidized is also called the reducing agent (it caused the other substance to be reduced). The substance reduced is also called the oxidizing agent (it caused the other substance to be oxidized). Zn(s) + CuCl 2 (aq) Cu(s) + ZnCl 2 (aq) Assigning Oxidation Numbers Assign oxidation numbers to each element in the equation below. Then identify the substance oxidized, reduced, the oxidizing agent, and the reducing agent. Cu(NO 3 ) 2 (aq) + Mg(s) Cu(s) + Mg(NO 3 ) 2 (aq) Elements get oxidized or reduced. 5 Compounds are oxidizing or reducing agents. 6 1

Balancing Redox Reactions Simple redox reactions can be balanced using half-reactions to balance atoms and charges.

Combine half reactions into one equation; atoms and charges have to balance. 7 Redox Reactions Figure 17.

exchange of Cu 2+ for Ag + ions in solution.

(credit: modification of work by Mark Ott) 8 Figure 17.4: If we separate the halfreactions into two separate containers, we can connect the metals via wire in a galvanic cell.

2 Balancing Redox Reactions Simple redox reactions can be balanced using half-reactions to balance atoms and charges. Zn(s) + Al 3+ (aq) Al(s) + Zn 2+ (aq) Separate equation into the oxidation half reaction and the reduction half reaction, including electrons. Then balance the equations so electrons cancel. Combine half reactions into one equation; atoms and charges have to balance. 7 Redox Reactions Figure 17.3: Cu (s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag(s) When a clean piece of copper metal is placed into a clear solution of silver nitrate (a), an oxidation-reduction reaction occurs that results in the exchange of Cu 2+ for Ag + ions in solution. As the reaction proceeds (b), the solution turns blue (c) because of the copper ions present, and silver metal is deposited on the copper strip as the silver ions are removed from solution. (credit: modification of work by Mark Ott) 8 Figure 17.4: If we separate the halfreactions into two separate containers, we can connect the metals via wire in a galvanic cell. Voltaic Cell: movie Electron motion in a cell 17.2 Galvanic Cells 9 Galvanic Cells - Components : Oxidation reaction (both vowels); mass of electrode decreases; solution charge increases; negative charge Metal is placed in its own solution (e.g., Cu(s) in Cu(NO 3 ) 2 (aq)) : Reduction reaction (both consonants); mass of electrode increases; solution charge decreases; positive charge Metal is placed in its own solution (e.g., Ag(s) in AgNO 3 (aq)) Wire: Electrons are transferred across a wire (through a voltmeter) from the anode (-) to the cathode (+) The cell potential, E (or Electromotive Force, EMF), is measured in Volts (V) Salt bridge - soluble salt solution used to maintain neutral charges of ions in each solution; cations travel toward cathode, anions toward anode 10 Galvanic Cells - Components Figure 17.5: A nonreactive, or inert, platinum wire allows electrons from the left beaker to move into the right beaker. Overall reaction is: Mg(s) + 2H + (aq) Mg 2+ (aq) + H 2 (g). Inert electrodes allow electrons to transfer but don t take part in the reaction. Non-metal Electrodes Zn/H2 cell: movie 12 2

Cu(s) Cu 2+ (1 M) Ag + (1 M) Ag(s) anode is phase boundary (separates solid from ions) is salt bridge electrons flow from left to right The Mg/H 2 Galvanic cell is represented by the notation: Mg(s)

3 Standard Reduction Potentials The standard potential of a Galvanic cell is measured by a voltmeter and is called the cell potential (E o cell ), measured in Volts (V).

3 Cu(s) Cu 2+ (1 M) Ag + (1 M) Ag(s) anode is phase boundary (separates solid from ions) is salt bridge electrons flow from left to right The Mg/H 2 Galvanic cell is represented by the notation: Mg(s) Mg 2+ (aq) H + (aq), H 2 (g) Pt(s). Write the cell notation for the reaction: 3 Zn(s) + 2 Al 3+ (aq) 2 Al(s) + 3 Zn 2+ (aq) Zn(s) Zn 2+ (aq) Al 3+ (aq) Al(s) Cell Notation cathode Standard Reduction Potentials The standard potential of a Galvanic cell is measured by a voltmeter and is called the cell potential (E o cell ), measured in Volts (V). The cell potential of a battery results from the difference in the electrical potentials for each electrode. The hydrogen electrode has been assigned a potential of 0 V, and is referred to as the Standard Hydrogen Electrode (SHE). 2H + (aq) + 2e H 2 (g) E o red = 0.0 V E = electromotive force (EMF) is the electrical potential that pushes e - away from anode (-) and pulls them toward cathode (+). Also called cell potential or cell voltage. 14 Standard Reduction Potentials In the Galvanic cell below, hydrogen is the anode (oxidized) and copper is the cathode (reduced). : H 2 (g) 2 H + (aq) + 2 e - E o ox = 0 V : Cu 2+ (aq) + 2e - Cu(s) E o red = V Figure 17.7 Standard Reduction Potentials In the Cu/Ag Galvanic cell, we can calculate the standard cell potential (E o cell) using cell potentials of each half reaction. : Cu(s) Cu 2+ (aq) + 2e - E o ox = V : Ag + (aq) + e - Ag(s) E o red = 0.80 V E o cell = E o red + E o ox = 0.80 V V = 0.46 V Notice that the copper half reaction is flipped from the previous example. When we reverse an equation, we change the sign of the potential (electron flow is reversed). How do we determine which half-reaction will be oxidized and which will be reduced? It depends on a species ability to reduce another substance. Standard Reduction Potentials Standard Reduction Potentials at 25 o C (See Table 17.2 partial table). Full list is in Appendix L (Standard Electrode (Half-Cell) Potentials. Increasing strength as oxidizing agents Half-Reaction Increasing strength as reducing agents E o red (V) F 2 (g) + 2e 2F (aq) Br 2 (l) + 2e 2Br (aq) Ag + (aq) + e Ag(s) Cu 2+ (aq) + 2e Cu(s) H + (aq) + 2e H 2 (g) 0.00 Pb 2+ (aq) + 2e Pb(s) Cd 2+ (aq) + 2e Cd(s) Zn 2+ (aq) + 2e Zn(s) Li + (aq) + e Li(s)

4 Activity Series Cell Potentials F 2 (g) + 2e - 2F - (aq) E o red = V This is the most positive E o red value. Makes sense: F 2 is the most electronegative element and easily gains an electron. F 2 (g) is easiest to reduce, a good oxidizing agent. Li + (aq) + e - Li(s) E o red = V This is the most negative E o red. Li + (aq) is hardest to reduce. Li(s) Li + (aq) + e - E o ox = V Li(s) is easiest to oxidize (easily loses an electron), a good reducing agent. 20 Based on their E o red values, determine the specific element or ion (include physical states) that is the: best oxidizing agent best reducing agent Au e - Au(s) Br 2 (l) + 2e - 2Br - (aq) Pb e - Pb(s) Ni e - Ni(s) Cell Potentials E o red = V E o red = V E o red = V E o red = V Standard Cell Potentials When given two half reactions, the more positive E o red will be the reduction half-reaction (no change). The more negative (or less positive) E o red will be reversed to give the oxidation half-reaction (E o ox = - E o red). Positive E o cell means the reaction is product-favored. The reaction will go forward until the battery dies. die when salt bridge, anode, or cathode runs out of material. Therefore, we want two half reactions that yield the most positive E o cell value. E o cell = E o ox + E o red For each pair of electrodes below, determine what will be the anode and what will be the cathode. Calculate E o cell for each combination. 1) Cu(s) and Cr(s) 2) Cu(s) and Mn(s) 3) Cu(s) and Au(s) Half-Reaction Au 3+ (aq) + 3e - Au(s) E red (V) Cu 2+ (aq) + 2e - Cu(s) Cu/Cr: Cr(s) anode, Cu(s) cathode; E o cell = V Cu/Mn: Mn(s) anode, Cu(s) cathode; E o cell = V Cu/Au: Cu(s) anode, Au(s) cathode; E o cell = V Example 17.4 Standard Cell Potentials Cr 3+ (aq) + 3e - Cr(s) Mn 2+ (aq) + 2e - Mn(s) The Nernst Equation E o cell is related to the thermodynamic quantity G o (both determine whether a reaction is reactant-favored or product-favored). Chemical energy in a Galvanic cell can do work, which is the product of the charge transferred (Coulombs discovered by Millikan and his Oil Drop experiment) and the potential difference (Volts). work = Volts x Coulombs = J F is Faraday s constant 24 4

5 The Nernst Equation G o = - nfe o cell n is the number of moles electrons F is Faraday s constant (96,500 C/mol e ) G = G o + RT ln Q (from Chapter 16) substitute nfe for G -nfe = -nfe o + RT ln Q (divide both sides by nf) E = E o (RT/nF) ln Q Nernst equation Non-Standard Cell Potential E = E o (RT/nF) ln Q Nernst equation When Q > 1, E < E ; When Q < 1, E > E Variations on this equation: E = E o cell - (2.303 RT) / nf log Q E = E o cell - ( V / n) log Q (at 25 o C) also on equation sheet At equilibrium, E = 0 and Q = K The equation becomes: E o cell = (RT/nF) ln K Variation on this equation: E o cell = ( V / n) log K (at 25 o C) Summary of Relationships Positive E o cell value G o is negative K is large Reaction is product-favored Negative E o cell value G o is positive K is small Reaction is reactant-favored G o, K from E o cell Calculate G o and K for a product-favored reaction at 25 o C using a Mg(s)-Al(s) Galvanic cell. Half-Reaction E red (V) Al 3+ (aq) + 3e Al(s) Mg 2+ (aq) + 2e Mg(s) Example Spontaneity at Nonstandard Conditions Calculate the cell potential, E, for the Mg-Al galvanic cell when [Mg 2+ ] = 0.60 M and [Al 3+ ] = M at 50 o C. Half-Reaction E red (V) Al 3+ (aq) + 3e Al(s) Mg 2+ (aq) + 2e Mg(s) Calculations Practice Calculate the cell potential, E, at 298 K when [Fe 2+ ] = 0.62 M, [Fe 3+ ] = 2.1 M and [Cd 2+ ] = 0.15 M? Is it spontaneous? Fe 3+ (aq) + e - Fe 2+ (aq) Cd 2+ (aq) + 2e - Cd(s) E o red = V E o red = V Balanced equation: Cd(s) + 2Fe 3+ (aq) 2Fe 2+ (aq) + Cd 2+ (aq) E o cell = V V = V Q = [Fe 2+ ] 2 [Cd 2+ ]/[Fe 3+ ] 2 = ( *0.15)/(2.1 2 ) = E = E o cell - (RT/nF) ln Q = V (spontaneous) Example

Calculate the cell potential at 25 o C when [Pb(NO 3 ) 2 ] = 0.88 M and [Ag(NO 3 )] = 0.14 M Ag + (aq) + e Ag(s) Pb 2+ (aq) + 2e - Pb(s) E o red = + 0.7996 V E o red = - 0.

9 E = 0.877 V Calculations Practice 31 17.6 Corrosion The term corrosion generally refers to the deterioration of a metal by an electrochemical process (e.g., rusting of iron).

The patina is formed by corrosion of the copper skin of the statue, which forms a thin layer of an insoluble compound that contains copper(ii), sulfate, and hydroxide ions.

(b) Exposure to the elements has resulted in the formation of the blue-green patina seen today. Image from Principles of General Chemistry http://2012books.lardbucket.

7: Once the paint is scratched on a painted iron surface, corrosion occurs and rust begins to form.

6 Calculate the cell potential at 25 o C when [Pb(NO 3 ) 2 ] = 0.88 M and [Ag(NO 3 )] = 0.14 M Ag + (aq) + e Ag(s) Pb 2+ (aq) + 2e - Pb(s) E o red = V E o red = V Write the short-hand notation of this reaction. Side note: gas pressures can also be used to calculate Q. E o cell = V V = V; n = 2 mol e - Q = (0.88 M) / (0.14 M) 2 = 44.9 E = V Calculations Practice Corrosion The term corrosion generally refers to the deterioration of a metal by an electrochemical process (e.g., rusting of iron). A view from the top of the Statue of Liberty, showing the green patina coating the statue. The patina is formed by corrosion of the copper skin of the statue, which forms a thin layer of an insoluble compound that contains copper(ii), sulfate, and hydroxide ions. (a) The Statue of Liberty is covered with a copper skin, and was originally brown, as shown in this painting. (b) Exposure to the elements has resulted in the formation of the blue-green patina seen today. Image from Principles of General Chemistry CC-BY-NC-SA license Corrosion of iron Figure 17.7: Once the paint is scratched on a painted iron surface, corrosion occurs and rust begins to form. The speed of the spontaneous reaction is increased in the presence of electrolytes, such as the sodium chloride used on roads to melt ice and snow or in salt water. Cathodic Protection Cathodic protection (CP) is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. A simple method of protection connects the metal to be protected to a more easily corroded sacrificial metal to act as the anode. The sacrificial metal then corrodes instead of the protected metal. Common applications are: steel water or fuel pipelines (e.g., water heaters), ship and boat hulls, metal reinforcement bars in concrete buildings and structures, and galvanized steel (zinc is sacrificial metal to protect steel metal from rusting). Source: Electrolysis Electrolysis The use of electric energy to drive a non-spontaneous chemical reaction (running a battery in reverse) is called electrolysis and uses an electrolytic cell. Electrolysis of molten sodium chloride; E o cell = 4 V Figure Water doesn t naturally decompose into hydrogen and oxygen. Electrolysis of water: E o cell = V; G o = kj/mol : 2H 2 O(l) O 2 (g) + 4H + (aq) + 4e : 4H + (aq) + 4e 2H 2 (g) 2H 2 O(l) O 2 (g) + 2H 2 (g) Figure 17.20: Water decomposes into oxygen and hydrogen gas during electrolysis. Sulfuric acid was added to increase the concentration of hydrogen ions and the total number of ions in solution, 35 2Cl (l) Cl 2 (g) + 2e 2Na + (l) + 2e 2Na(l) 2Na + (l) + 2Cl (l) 2Na(l) + Cl 2 (g) 36 6

Electrolysis Faraday developed the quantitative treatment of electrolysis.

2 A is passed through the cell for 1.50 hr? Electrolysis If a constant current of 0.912 A is passed through an electrolytic cell containing molten MgCl 2 and produces 7.

5 and Fuel Cells Pb(s) + 2HSO 4 (aq) PbSO 4 (s) + H + (aq) + 2e PbO 2 (s) + 3H + (aq) + HSO 4 (aq) + 2e PbSO 4 (s) + 2H 2 O(l) Pb(s) + PbO 2 (s) + 2H 2 SO 4 (aq) 2PbSO 4 (s) + 2H 2 O(l)

5 V It is important to remember that the voltage delivered by a battery is the same regardless of the size of a battery.

As the zinc container oxidizes, its contents eventually leak out, so this type of battery should not be left in any electrical device for extended periods.

7 Electrolysis Faraday developed the quantitative treatment of electrolysis. 2Cl (l) Cl 2 (g) + 2e Ca 2+ (l) + 2e Ca(l) Ca 2+ (l) + 2Cl (l) Ca(l) + Cl 2 (g) Charge (C) = Current (A) x time (s) What mass of chlorine gas can be produced when 45.2 A is passed through the cell for 1.50 hr? Electrolysis If a constant current of A is passed through an electrolytic cell containing molten MgCl 2 and produces 7.44 g of Mg, how long does the current have to be passed? Mg 2+ (l) + 2e Mg(l) 37 Example Figure 17.4: Lead Storage (Car) 6 cells 2 V per cell 12 V total Rechargeable 17.5 and Fuel Cells Pb(s) + 2HSO 4 (aq) PbSO 4 (s) + H + (aq) + 2e PbO 2 (s) + 3H + (aq) + HSO 4 (aq) + 2e PbSO 4 (s) + 2H 2 O(l) Pb(s) + PbO 2 (s) + 2H 2 SO 4 (aq) 2PbSO 4 (s) + 2H 2 O(l) Figure 17.10: Dry Cell (or Laclanche) : Zinc container (anode), Graphite rod (cathode), 1.5 V It is important to remember that the voltage delivered by a battery is the same regardless of the size of a battery. For this reason, D, C, A, AA, and AAA batteries all have the same voltage rating. However, larger batteries can deliver more moles of electrons. As the zinc container oxidizes, its contents eventually leak out, so this type of battery should not be left in any electrical device for extended periods. Zn(s) Zn 2+ (aq) + 2e 2 NH 4+ (aq) + 2MnO 2 (s) + 2e Mn 2 O 3 (s) + 2NH 3 (aq) + H 2 O(l) Zn(s) + 2NH 4+ (aq) + 2MnO 2 (s) Zn 2+ (aq)+ Mn 2 O 3 (s) + 2NH 3 (aq) + H 2 O(l) 40 9 Volt 6 x 1.5 V Dry Cell batteries connected in series Lithium ion (Rechargeable) batteries: charge flows between the electrodes as the lithium ions move between the anode and cathode Cell potential: 3.7 V 41 Li(s) Li + + e Li + (aq) + CoO 2 + e LiCoO 2 (s) Li + (s) + CoO 2 LiCoO 2 (s) 42 7

Tesla Motors 60 kwh battery rated to deliver 230 miles (208 by EPA) 85 kwh battery rated to deliver 320 miles (265 by EPA)

How it works 43 2H 2 (g) + 2O 2- (aq) 2H 2 O(l) + 4e O 2 (aq) + 4e 2O 2- (aq) 2H 2 (g) + O 2 (g) 2H 2 O(l) 44 Figure 17.

43 V Alkaline batteries were designed as direct replacements for zinc-carbon (dry cell) batteries in the 1950 s.

NiCd : Nickel-plated cathode and cadmium-plated anode.

8 Tesla Motors 60 kwh battery rated to deliver 230 miles (208 by EPA) 85 kwh battery rated to deliver 320 miles (265 by EPA) Contains 7,104 lithium-ion battery cells in 16 modules wired in series. The battery pack uses Ni-Co-Al cathodes. Fuel Cells; E o cell = 1.23 V In this hydrogen fuel-cell schematic, oxygen from the air reacts with hydrogen, producing water and electricity. How it works 43 2H 2 (g) + 2O 2- (aq) 2H 2 O(l) + 4e O 2 (aq) + 4e 2O 2- (aq) 2H 2 (g) + O 2 (g) 2H 2 O(l) 44 Figure 17.11: Alkaline : Zinc container (anode), Graphite rod (cathode), 1.43 V Alkaline batteries were designed as direct replacements for zinc-carbon (dry cell) batteries in the 1950 s. They can deliver about 3 5x the energy of a zinc-carbon dry cell battery of similar size. NiCd : Nickel-plated cathode and cadmium-plated anode. NiCd batteries use a jelly-roll design that significantly increases the amount of current the battery can deliver as compared to a similarsized alkaline battery. Zn(s) + 2OH - (aq) ZnO(s) + H 2 O(l) + 2e 2MnO 2 (s) +H 2 O(l) + 2e Mn 2 O 3 (s) + 2OH - (aq) Cd(s) + 2OH - (aq) Cd(OH) 2 (s) + 2e NiO 2 (s) + 2H 2 O(l) + 2e Ni(OH0 2 (s) + 2OH - (aq) Zn(s) + 2MnO 2 (s) ZnO(s)+ Mn 2 O 3 (s) 45 Cd(s) + NiO 2 (s)+ 2H 2 O(l) Cd(OH) 2 (s)+ Ni(OH) 2 (s) 46 8

Oxidation-Reduction Review. Electrochemistry. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Sample Problem.

Oxidation-Reduction Review. Electrochemistry. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Sample Problem. 1 Electrochemistry Oxidation-Reduction Review Topics Covered Oxidation-reduction reactions Balancing oxidationreduction equations Voltaic cells Cell EMF Spontaneity of redox reactions Batteries Electrolysis