CHM 170 1nd EditionFinal Exam Study Guide Lectures: 21 - 25Lecture 21 (April 21)This section for exam four is about electrochemistry and radioactivity and nuclear chemistry. There are some general rules for assigning an oxidation number. General Rules:1. For an atom in its elemental form (Na, O2, Cl2, etc.): O.N. = 02. For a monoatomic ion: O.N. = ion charge3. The sum of O.N. values for the atoms in a compound equals zero. The sum of O.N. values for the atoms in a polyatomic ion equals the ion’s charge.Rules for specific atoms or periodic table groups:1. for Group 1A (1): O.N. = +1 in all compounds2. for Group 2A (2): O.N. = +2 in all compounds3. for hydrogen: O.N. = +1 in combination with nonmetals or -1 with some metals4. for fluorine: O.N. = -1 in combination with metals and boron5. for oxygen: O.N. = -1 in peroxides. O.N.= -2 in all other compounds (except with F)6. For Group 7A (17): O.N. = -1 in combination with metals, nonmetals (except O), and other halogens lower in the groupPolyatomic don’t change for example, OH-.H Cl O2+1 -2-2= +1 = +3 = -4Since the oxidation should equal to zero and you know the rules for H and O then Cl is just the remainder to make the whole chemical equation to equal zero. If the equation was HClO2 – then you would have to find the numbers that equals to -1 and not 0. Redox Reactions:Redox reaction are those involving the oxidation and reduction of species.OIL – Oxidation Is Loss of electrons.RIG – Reduction Is Gain of electrons.Oxidation and reduction must occur together.They cannot exist alone.Key points about redox reactions:o Oxidation (electron loss) always accompanies reduction (electron gain).o The oxidizing agent is reduced, and the reducing agent is oxidized.o The number of electrons gained by the oxidizing agent always equals the number lost by the reducing agent.o Oxidized and reduce can only be a reactant.Examples:Identify the Oxidizing and Reducing Agents in Each of the Following3 H2S + 2 NO3– + 2 H+ ® 3 S + 2 NO + 4 H2O+1 -2 +5 -2 +1 0 +2 -2 +1 -2 Lecture 22 (April 23rd)Balancing Redox Reactions Example:Fe2+ + Cr2O72- → Fe3+ + Cr3+1. Separate the equation into two half-reactions.Oxidation: Fe2+ →Fe3+Reduction: Cr2O72- → Cr3+2. Balance the atoms other than O and H in each half-reactionCr2O72- →2Cr3+3. For reactions in acid, add H2O to balance O atoms and H+ to balance H atoms.Cr2O72- →2Cr3+ + 7H2O14H+ + Cr2O72- → 2Cr3+ + 7H2O4. Add electrons to one side of each half-reaction to balance the charges on the half-reaction.Fe2+ →Fe3+ + 1e-6e- + 14H+ + Cr2O72- →2Cr3+ + 7H2O5. If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.6Fe2+→ 6Fe3+ + 6e-6e- + 14H+ + Cr2O72- →2Cr3+ + 7H2O6. Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel.14H+ + Cr2O72- + 6Fe2+→6Fe3+ + 2Cr3+ + 7H2O7. For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation.Redox Reactions & Current:o redox reactions involve the transfer of electrons from one substance to anothero therefore, redox reactions have the potential to generate an electric currento In order to use that current, we need to separate the place where oxidation is occurring from the place that reduction is occurring.A voltaic cell using inactive electrodes:Oxidation half-reaction: Reduction half-reaction:2I-(aq) → I2(s) + 2e-MnO4-(aq) + 8H+(aq) + 5e- →Mn2+(aq) + 4H2O(l)Overall (cell) reaction:2MnO4-(aq) + 16H+(aq) + 10I-(aq) → 2Mn2+(aq) + 5I2(s) + 8H2O(l)Electrochemical CellsAn electrochemical cell is a system consisting of electrodes that dip into an electrolyte in which a chemical reaction either uses or generates an electric current.o Electrodes: are usually metal strips/wires connected by an electrically conducting wire.o Salt Bridge: is a U-shaped tube that contains a gel permeated with a solution of an inert electrolyte.o Anode: is the electrode where oxidation takes place.o Cathode: is the electrode where reduction takes place. A voltaic, or galvanic, cell is an electrochemical cell in which a spontaneous reaction generates an electric current.o Energy is released from spontaneous redox reactionAn electrolytic cell is an electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction.o Energy is absorbed to drive a nonspontaneous redox reactionLecture 23 (April 28th)Notation for Voltaic Cells• It is convenient to have a shorthand way of designating particular voltaic cells.The cell consisting of the zinc-zinc ion half-cell and the copper-copper ion half-cell, is writtenZn( s )|Zn2+(aq )||Cu2 +( aq)|Cu(s)Anode salt bridge Cathode- The anode (oxidation half-cell) is written on the left. The cathode (reduction half-cell) is written on the right.- The two electrodes are connected by a salt bridge, denoted by two vertical bars.- The cell terminals are at the extreme ends in the cell notation.- A single vertical bar indicates a phase boundary, such as between a solid terminal and the electrode solution.Example:The answer is E.Another problem to consider:Give the overall cell reaction for the voltaic cellCd( s)|Cd2+(1 . 0 M )||H+(aq )|H2(1. 0 atm )|PtThe half-cell reactions are…Cd( s)→Cd2 +( aq)+2 e−2 H+(aq )+2e−→ H2(g )Overall: Cd( s )+2 H+(aq )→Cd2+(aq )+ H2( g)Electromotive Force:- You measure this quantity in volts.- The volt, V, is the SI unit of potential difference equivalent to 1 joule of energy per coulomb of charge.Formula: 1 volt=1 JC• The Faraday constant, F, is the magnitude of charge on one mole of electrons; it equals 96,500 coulombs (9.65 x 104 C).Formula: work( J )=−F (coulombs )×volts( J/ coulomb )Standard Cell emf’s and Standard Electrode PotentialsA cell emf is a measure of the driving force of the cell reaction.The reaction at the anode has a definite oxidation potential, while the reaction at the cathode has a definite reduction potential. Thus, the overall cell emf is a combination of these two potentials…Formula: Ecell = Ecatode - EanodeA reduction potential is a measure of the tendency to gain electrons in the reduction half-reaction. The oxidation potential for an oxidation half-reaction is the negative of the reduction potential for the reverse reaction.Calculating Cell emf’s from Standard Potentials•