Chapter Eighteen Electrochemistry Today Turn in: Nothing Our Plan: Test Results Notes Redox Equations
Worksheet #1 Homework (Write in Planner): Redox Equation WS due Friday 6% Curve A 2 B 5 C 4 D 1 F Average 72.72% High Score 90% (x2) OxidationReduction: The Transfer of Electrons
Silver metal is formed, and the solution turns blue from copper(II) ions formed. Electrons from copper metal are transferred to silver ions. Half-Reactions (Review)
In any oxidationreduction reaction, there are two half-reactions: Oxidation: a species loses electrons to another species. (LEO) Reduction: a species gains electrons from another species. (GER) Both oxidation and reduction must occur simultaneously. A species that loses electrons must lose them to something else (something that gains them). A species that gains electrons must gain them from something else (something that loses them). Oxidation Numbers An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion. Examples: in NaCl, the oxidation number of Na is +1, that of Cl is 1 (the actual charge). In CO2 (a molecular compound, no ions) the oxidation number of oxygen is 2, because oxygen as an ion would be expected to have a 2 charge. What is the oxidation number on carbon in CO2? Rules for Assigning Oxidation Numbers
For the atoms in a neutral speciesan isolated atom, a molecule, or a formula unitthe sum of all the oxidation numbers is 0. This includes elements in their standard state (Cu (s), Cl2 (g), etc.) For the atoms in an ion, the sum of the oxidation numbers is equal to the charge on the ion. In compounds, the group 1A metals all have an oxidation number of +1 and the group 2A metals all have an oxidation number of +2. In compounds, the oxidation number of fluorine is 1. In compounds, hydrogen has an oxidation number of +1. In most compounds, oxygen has an oxidation number of 2. In binary compounds with metals, group 7A elements have an oxidation number of 1, group 6A elements have an oxidation number of 2, and group 5A elements have an oxidation number of 3. Assign oxidation numbers to the elements in each compound: NO3-1 SO2
Lets Try It Assign oxidation numbers to the elements in each compound: NO3-1 SO2 Fe2O3 Cu (s) Identifying OxidationReduction Reactions
In a redox reaction, the oxidation number of a species changes during the reaction. Oxidation occurs when the oxidation number increases (species loses electrons). LEO Reduction occurs when the oxidation number decreases (species gains electrons). GER Another mnemonic is OIL RIG! If any species is oxidized or reduced in a reaction, that reaction is a redox reaction. Examples of redox reactions: displacement of an element by another element; combustion; incorporation of an element into a compound, etc. Review Fe2O3 + S --> Fe + SO2
In the reactions below, write oxidation numbers for each substance and identify which substance is oxidized and which is reduced? N2O H > H2O NH3 K KNO >N K2O Fe2O S--> Fe+ SO2 The Half-Reaction Method of Balancing Redox Equations
Separate a redox equation into two half-equations, one for oxidation and one for reduction. Balance the number of atoms of each element in each half-equation. Usually we balance O and H atoms last. Balance each half-reaction for charge by adding electrons to the left in the reduction half-equation and to the right in the oxidation half-equation. Adjust the coefficients in the half-equations so that the same number of electrons appears in each half-equation. Add together the two adjusted half-equations to obtain an overall redox equation. Simplify the overall redox equation as necessary. Redox Reactions in Acidic and in Basic Solution
Redox reactions in acidic solution and in basic solution may be very different from one another. If acidic solution is specified, we must add H2O and/or H+ as needed when we balance the number of atoms. If basic solution is specified, the final equation may have OH and/or water molecules in it. A simple way to balance an equation in basic solution: Balance the equation as though it were in acidic solution. Add as many OH ions to each side as there are H+ ions in the equation. Combine the H+ and OH ions to give water molecules on one side, and simplify the equation as necessary. MnO4-1 + S2O3-2 --> Mn+2 + SO4-2
Example 18.1 Balance the equation in acidic solution: MnO4-1 + S2O3-2 --> Mn+2 + SO4-2 P4 (s) + H+1 (aq) + NO3-1 (aq) --> H2PO4-1 (aq) + NO (g)
Example 18.1 B Write a balanced equation for the oxidation of phosphorus by nitric acid, which is described by P4 (s) + H+1 (aq) + NO3-1 (aq) --> H2PO4-1 (aq) + NO (g) Br2 (l) --> Br-1 (aq) + BrO3-1 (aq)
Example 18.2 Balance the following in basic solution: Br2 (l) --> Br-1 (aq) + BrO3-1 (aq) CrO4-2 (aq) + AsH3 (g) --> Cr(OH)3 (s) + As (s)
Example 30a Balance the following in basic solution: CrO4-2 (aq) + AsH3 (g) --> Cr(OH)3 (s) + As (s) HCl + K2Cr2O7 --> KCl + CrCl3 + H2O + Cl2
WS Example 1 HCl + K2Cr2O7 --> KCl + CrCl3 + H2O + Cl2 FeCl2 + KMnO4 + HCl --> FeCl3 + KCl + MnCl2 + H2O
WS Example 2 FeCl2 + KMnO4 + HCl --> FeCl3 + KCl + MnCl2 + H2O S-2 + MnO4-1 --> S + MnO2 (basic solution)
WS Example 3 S-2 + MnO4-1 --> S + MnO2 (basic solution) CuS + NO3-1 --> Cu+2 + S + NO (acidic)
WS Example 4 CuS + NO3-1 --> Cu+2 + S + NO (acidic) Stop! Do the Redox Equations Worksheet. Today Turn in: Get out Redox WS Our Plan: Questions on Redox WS Quiz
Notes Galvanic Cells WS Homework (Write in Planner): Worksheet Due Wednesday A Qualitative Description of Voltaic Cells
A voltaic cell uses a spontaneous redox reaction to produce electricity. A half-cell consists of an electrode (strip of metal or other conductor) immersed in a solution of ions. This Zn2+ becomes a Zn atom. Both oxidation and reduction occur at the electrode surface, and equilibrium is reached. This Zn atom leaves the surface to become a Zn2+ ion. Important Electrochemical Terms
An electrochemical cell consists of two half-cells with the appropriate connections between electrodes and solutions. Two half-cells may be joined by a salt bridge (U shaped tube suspended in gel) that permits migration of ions, without completely mixing the solutions. The anode is the electrode at which oxidation occurs. The cathode is the electrode at which reduction occurs. Helpful Mnemonic AUTO (anode = oxidation) CAR (cathode = reduction)
OR: To remember the charge: Ca+ions are attracted to the Ca+hode (the t is a plus sign) To remember which reaction occurs at which terminal: An Ox and Red Cat - Anode Oxidation, Reduction Cathode Important Electrochemical Terms
In a voltaic cell, a spontaneous redox reaction occurs and current (electricity) is generated. Cell potential (Ecell) is the potential difference in volts between anode and cathode. Ecell is the driving force that moves electrons and ions. A ZincCopper Voltaic Cell
Positive and negative ions move through the salt bridge to equalize the charge. the electrons produced move through the wire to the Cu(s) electrode, where they are accepted by Cu2+ ions to form more Cu(s). Zn(s) is oxidized to Zn2+ ions, and Reaction:Zn(s) + Cu2+(aq)--> Cu(s)+ Zn2+(aq) Cell Diagrams A cell diagram is shorthand for an electrochemical cell. The anode is placed on the left side of the diagram. The cathode is placed on the right side. A single vertical line ( | ) represents a boundary betweenphases, such as between an electrode and a solution. A double vertical line ( || ) represents a salt bridge or porous barrier separating two half-cells. Reaction: Zn(s) + Cu2+(aq) --> Cu(s) + Zn2+(aq) Another look at Cell Diagrams Example 18.3 Describe the half-reactions and the overall reaction that occur in the voltaic cell represented by the cell diagram: Pt(s) | Fe2+(aq), Fe3+(aq) || Cl(aq) | Cl2(g) | Pt(s) Standard Electrode Potentials
Since an electrode represents only a half-reaction, it is not possible to measure the absolute potential of an electrode. The standard hydrogen electrode (SHE) provides a reference for measurement of other electrode potentials. The SHE is arbitrarily assigned a potential of V. Standard Electrode Potentials
The standard electrode potential, E, is based on the tendency for reduction to occur at an electrode. (Sometimes it is referred to as standard reduction potential) All standard reduction potentials are for 25 C and 1 M solutions. Efor the standard hydrogen electrode is arbitrarily assigned a value of V. All other values of E are determined relative to the standard hydrogen electrode. Standard Electrode Potentials
The standard cell potential (Ecell) is the difference between E of the cathode and E of the anode. Ecell= E(cathode)E(anode) Remember AUTO CAR! Measuring the Standard Potential of the Cu2+/Cu Electrode
The voltmeter reading and the direction of electron flow tell us that Cu2+ is more easily reduced than H+, by volts. Standard hydrogen electrode Cu2+ + 2e --> Cu E = V Measuring the Standard Potential of the Zn2+/Zn Electrode
The voltmeter reading and the direction of electron flow tell us that Zn2+ is harder to reduce than H+, by volts. Standard hydrogen electrode Zn2+ + 2e --> Zn E = V F2 is the strongest oxidizing agent F is the weakest reducing agent
Li is the strongest reducing agent Important Note The formula Ecell= E(cathode)E(anode) allows us to simply use the numbers in the table, but if you were asked for the oxidation value of a reaction, you would have to reverse the sign since the values are all reduction potentials. Important Points about Electrode and Cell Potentials
Standard electrode potentials and cell voltages are intensive properties; they do not depend on the total amounts of the species present. A flashlight battery (D-cell) and a penlight battery (AA cell) produce the same potential1.5 volts. E does depend on the particular species in the reaction (or half-reaction). As we shall learn later, cell and electrode potentials can depend on concentration of the species present. Example 18.4 Determine E for the reduction half-reaction
Ce4+(aq) + e --> Ce3+(aq), given that the cell voltage for the voltaic cell Co(s) | Co2+(1 M) || Ce4+(1 M), Ce3+(1 M) | Pt(s) is Ecell = V. Example 18.5 Balance the following oxidationreduction equation, and determine Ecell for the reaction. O2(g) + H+(aq) + I(aq)-->H2O(l) + I2(s) Electrode Potentials, Spontaneous Change, and Equilibrium
An electrochemical cell does work. welec = nFEcell n = number of electrons in the balanced equation F = 96,485 coulombs per mole. The amount of electrical work is also equal to DG: DG = nFEcell Under standard conditions: DG = nFEcell Criteria for Spontaneous Change in Redox Reactions
If Ecell is positive, the forward reaction is spontaneous. If Ecell is negative, the forward reaction is nonspontaneous (the reverse reaction is spontaneous). If Ecell = 0, the system is at equilibrium. When a cell reaction is reversed, Ecell and DG change signs. Example 18.6 Will copper metal displace silver ion from aqueous solution? That is, does the reaction Cu(s) + 2 Ag+(1 M)-->Cu2+(1 M) + 2 Ag(s) occur spontaneously from left to right? The Activity Series Revisited
In the activity series of metals (Section 4.4), any metal in the series will displace a metal below it from a solution of that metals ions. Theoretical basis: The activity series lists metals in order of their standard potentials. Displacement of a metal from a solution of its ions by a metal higher in the series corresponds to a positive value of Ecell and a spontaneous reaction. Visual Example Equilibrium Constants in Redox Reactions
Whereas potential and free energy are related, and free energy and equilibrium are related, equilibrium and potential must be related to one another. DG = nFEcell and DG = RT ln Keq thereforeRT ln Keq = nFEocell RT ln Keq RT Ecell == ln Keq nF nF R and F are constant, therefore at 298 K: V Ecell = ln Keq n Example 18.8 Calculate the values of G and Keq at 25 C for the reaction Cu(s) + 2 Ag+(1 M)-->Cu2+(1 M) + 2 Ag(s) Thermodynamics, Equilibrium, and Electrochemistry: A Summary
From any one of the three quantities Keq, G, Ecell, we can determine the others. Stop! Begin working on Worksheet #2! Today Turn in: Nothing Our Plan: Quick Review
Notes Concentration & Electrochemical Cells Finish Electrochemical Cells WS Homework (Write in Planner): Worksheet Due Wednesday Effect of Concentrations on Cell Voltage
A nonstandard cell differs in potential from a standard cell (1 M concentrations, 1 atm partial pressures). Effect of Concentrations on Cell Voltage
From the previous relationships we can show that: RT Ecell = Ecell ln Q nF At 25 C, and converting to common logarithms: V Ecell = Ecelllog Q n This Nernst equation relates a cell voltage for nonstandard conditions, Ecell, to a standard cell voltage, Ecell, and to the concentrations of reactants and products expressed through the reaction quotient, Q. We can use the Nernst equation to find cell potential from concentrations, or we can measure Ecell and determine the concentration of a species in the cell. Example 18.9 Calculate the expected voltmeter reading for the voltaic cell pictured in Figure Another Nernst Equation Example
Use the Nernst equation to determine Ecell at 25C for the following voltaic cells. Zn(s) | Zn2+(2.0 M) || Cu2+(0.050 M) | Cu (s) Zn(s) | Zn2+(0.050 M) || Cu2+(2.0 M) | Cu (s) Batteries: Using Chemical Reactions to Make Electricity
We often call any device that stores chemical energy for later release as electricity a battery. Technically, a D, C, or AA battery is actually a single electrochemical cell. A battery consists of several cells joined together to produce higher current or higher voltage. A 9-volt transistor battery, an automobile battery, and a rechargable battery pack are all true batteries. The Dry Cell A primary cell employs an irreversible chemical reaction.
When the reactants inside the cell are largely used up, the cell is dead. The LeClanch cell or dry cell (right) is the ordinary type of flashlight battery. Alkaline cells cost more than the LeClanch cell but they have a longer shelf life and longer service life. The LeadAcid Storage Battery
A leadacid storage battery used in an automobile uses secondary cells; they are rechargeable. By connecting the cell to an external electric energy source, the discharge reaction is reversed. Cell reaction:Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4(aq) --> 2 PbSO4(s) + 2 H2O(l) Charging reaction: 2 PbSO4(s) + 2 H2O(l) --> Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4(aq) Other Secondary Cells The nickelcadmium (NiCd) cell uses a cadmium anode and a cathode containing Ni(OH)2. A NiCd cell can be recharged hundreds of times. It produces 1.2 V (a Leclanch cell produces 1.5 V). Nickelmetal hydride cells (NiMH) use a metal alloy anode that contains hydrogen. In use, the anode releases the hydrogen, forming water. Like the NiCd cell, a NiMH cell produces 1.2 V. Lithium-ion cells use a lithiumcobalt oxide or lithiummanganese oxide material as the anode. The electrolyte is an organic solvent containing a dissolved lithium salt. Many modern laptop computers and cellular phones use lithium-ion cells. Fuel Cells In a fuel cell, the cell reaction is equivalent to a combustion reaction. The reactants are supplied externally; the cell does not go dead as long as the oxidizing and reducing agents are provided. Fuel cells are generally operated under nonstandard conditions and at temperatures considerably higher than 25 C. H2O2 fuel cells are seeing use in some automobiles. Corrosion: Metal Loss Through Voltaic Cells
In moist air, iron is oxidized to Fe2+, particularly at scratches, nicks, or dents. These areas are referred to as anodic areas. Other regions of the iron serve as cathodic areas, where the electrons from the anodic areas reduce O2 to OH. Iron(II) ions migrate from the anodic areas to the cathodic areas where they combine with the hydroxide ions. The iron(II) is then further oxidized to iron(III) by atmospheric oxygen. Common rust is Fe2O3 x H2O. Corrosion of an Iron Piling
One way to minimize rusting is to provide a different anode reaction. Protecting Iron from Corrosion
The simplest defense against corrosion of iron is to coat it (with paint or metal) to exclude oxygen from the surface. An entirely different approach is to protect iron with a more active metal. Galvanized iron has been coated with zinc. The zinc provides an alternative anode reaction; the zinc corrodes, protecting the iron. Cathodic Protection In cathodic protection, an iron object to be protected is connected to a chunk of an active metal. The iron serves as the reduction electrode and remains metallic. The active metal is oxidized. Water heaters often employ a magnesium anode for cathodic protection. Stop! Finish Worksheet #2 Electrochemical Cells Worksheet Today Turn in: Get out Electrochemical Cells WS Our Plan:
Questions on WS Electrochemical Cell Review Notes Electrolysis WS Homework (Write in Planner): Worksheet Due Friday (LAST HOMEWORK) Electrolytic Cells A voltaic cell corresponds to a spontaneous cell reaction. An electrolytic cell corresponds to a nonspontaneous cell reaction. The reaction is called electrolysis. An electrolytic cell is the opposite of a galvanic cell. The external source of electricity acts like an electron pump. It pulls electrons away from the anode, where oxidation takes place, and pushes them toward the cathode, where reduction takes place. The polarities of the electrodes are reversed from those in the voltaic cell, because now the external source controls the flow of electrons. Electrolysis of Molten Sodium Chloride
2 NaCl(l) --> 2 Na(l)+Cl2(g) The nonspontaneous reaction is driven by external potential. Molten NaCl (around 1000 C) Predicting Electrolysis Reactions
In an electrolytic cell, all combinations of cathode and anode half-reactions give negative values of Ecell. The reaction most likely to occur is the one with the least negative value of Ecell (requires the lowest applied voltage from the external electricity source).HOWEVER In many half-reactions, particularly those involving gases, various interactions at electrode surfaces make the required voltage for electrolysis higher than the voltage calculated from E data. Overvoltage is the excess voltage above the voltage calculated from E values that is required in electrolysis. Quantitative Electrolysis
The unit of electric charge is the coulomb (C), and the charge carried by one electron is x 1019 C. The conversion factor well use is 96,485 C/1 mole electrons Electric current, expressed in amperes (A), is the rate of flow of electric charge (C/s). Quantitative Electrolysis
To calculate the quantitative outcome of an electrolysis reaction: Determine the amount of charge (C)the product of current and time. Convert the amount of charge to moles of electrons. Use a half-equation to relate moles of electrons to moles of a reactant or a product. Convert from moles of reactant or product to the final quantity desired. Example 18.13 We can use electrolysis to determine the gold content of a sample. The sample is dissolved, and all the gold is converted to Au3+(aq), which is then reduced back to Au(s) on an electrode of known mass. The reduction half-reaction is Au3+(aq) + 3e Au(s).What mass of gold will be deposited at the cathode in 1.00 hour by a current of 1.50 A? Try It Out! 18.13 A:For how many minutes must the electrolysis of a solution of CuSO4 (aq) be carried out, at a current of 2.25 A, to deposit 1.00 g of Cu (s) at the cathode? Example 18.14 An Estimation Example
Without doing detailed calculations, determine which of the following solutions will yield the greatest mass of metal at a platinum cathode during electrolysis by a 1.50-A electric current for 30.2 min: CuSO4(aq), AgNO3(aq), or AuCl3(aq). Producing Chemicals by Electrolysis
Electrolysis plays an important role in the manufacture and purification of many substances, including chlorine, copper, silver, magnesium, aluminum, lead, zinc, sodium, fluorine, titanium, sodium hydroxide, hydrogen Electrolysis of NaCl(aq) is used to produce H2, Cl2, and NaOH, all of which have important industrial uses. Electroplating Electrolysis can be used to coat one metal onto another, a process called electroplating. Usually, the object to be electroplated, such as a spoon, is cast of an inexpensive metal. It is then coated with a thin layer of a more attractive, corrosion-resistant, and expensive metal, such as silver or gold. Complete the Electrolysis Worksheet
Stop Complete the Electrolysis Worksheet Today Turn in: Get out Electrolysis WS Our Plan: Questions on WS
Electrochemical Cells Online Lab Homework (Write in Planner): Take Home Test Due Monday OR prepare for the AP Exam! A Little Redox Humor Today Turn in: Take Home Test Our Plan:
Begin Qualitative Analysis Lab (read the pre-lab and procedure FIRST) Homework (Write in Planner): Work on Lab Report Lab Preparation Read the Packet, including the entire procedure.
Answer the Pre-Lab Questions Begin working on your lab report I recommend recording data in the handout that I gave you and then transferring it to your final copy. You could begin working on the title, purpose, pre-lab, and procedure, though. Reminders Safety Goggles, gloves, apron at ALL TIMES!
Clean all glassware and lab supplies before and after use. Label things well (include your initials)! Seal containers to store until next class period.Place them on the back bookshelf. All solids go in the trash and liquids down the drain EXCEPT barium (anion step 2 &3).Put them in the waste container in the hood! Run tests until you get clear and reproducible results (you can repeat steps if you need to). Report is due 5/14 at the beginning of class! Today Homework (Write in Planner): Turn in: Nothing Our Plan:
Qualitative Analysis Lab Homework (Write in Planner): Report Due on 8/14