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Electrochemistry Oxidation – Reduction (Redox) Reactions Involves the transfer of electrons from one species to another Both oxidation and reduction must occur Burning Magnesium Demo The balance of the overall reaction seen above can be written as follows. 2 Mg (s) + O2 (g) -------> 2 MgO 1) What is happening to magnesium in this reaction? Answer: Magnesium is losing electrons to become a positive ion . This is the process of oxidation. The reaction we can write for this process is called the oxidation half - reaction. Mg0 (s) --------> Mg 2+ + 2 e - ( electrons ) Since two magnesium atoms are reacting a total of 4 electrons are lost by the magnesium 2) What is happening to the Oxygen in this reaction? Answer: Oxygen is gaining electrons to become a negative ion. This is the process of reduction. The reaction we write for this process is called the reduction half- reaction. O0 (g) ----------> O 2- + 2e Since oxygen is a diatomic molecule two oxygen atoms are undergoing reduction and a total of 4 electrons are gained Note: The total number of electrons lost in oxidation = the total number gained in reduction. This is true of all redox reactions Oxidation – Reduction (Redox) Reactions Oxidation number: Charge on a single atom or ion or the apparent charge on atoms in multiatom combinations Oxidation: the lose of electrons, oxidation number increases Reduction: the gain of electrons, oxidation number decreases Oxidizing Agent (oxidant): substance oxidizing another (being reduced) Reducing Agent (reductant): substance reducing another (being oxidized) Oxidation – Reduction (Redox) Reactions Oxidation state of a free element is zero Oxidation state of group I, II, III metals: +1, +2, +3 Oxidation state of F is always –1 Oxidation state of H is +1 except in metallic hydrides where it is –1 Oxidation state of O is –2 except with F where it is +2, in peroxides where it is –1 or in superoxides where it is –1/2 Oxidation state of monatomic ion = ion’s charge Sum of oxidation states of all atoms in a polyatomic ion = ion’s charge Sum of oxidation states of all atoms or ions in the formula of a compound = zero Example In each of the following equations, is hydrogen peroxide acting as a reducing or oxidizing agent? 2MnO4-(aq) + 5H2O2(aq) + 6H+(aq) 2Mn2+(aq) + 5O2(g) + 8H2O(l) PbS(s) + 4H2O2(aq) PbSO4(aq) + 4 H2O(l) Balancing Redox Equations Half Reaction Method (Acidic Solution) Write the oxidation half reaction ( electrons are products) Write the reduction half reaction ( electrons are reactants) Balance elements other than H and O in each reaction Balance O by adding H2O Balance H by adding H+ Balance charges by adding electrons Add the 2 half reactions and simplify where possible Check Balancing Redox Equations Half Reaction Method (Basis Solution) Follow steps 1 thru 5 as if it were in acidic solution Balance H+ by adding OH- Example: Complete and balance the following equations by the method if half reactions. Cr2O72-(aq) + Cl-(aq) Cr3+ (aq) + Cl2(g) (acidic) CN-(aq) + MnO4-(aq) CNO-(aq) + MnO2(s) (basic) Voltaic (galvanic) Cells Energy released in a chemical reaction can be used to do electrical work Device used to perform electrical work by the transfer of electrons through an external pathway Electrodes: 2 solid metals connected by the external circuit Cathode (+): electrode where reduction occurs Anode (-): electrode where oxidation occurs Salt Bridge: Maintains neutrality in a galvanic cell Contains an electrolytic solution whose ions will not react with other ions in the cell Electron Flow: From anode to cathode Voltaic (galvanic) Cells Standard cell notation: anode/anode solution // cathode solution/ cathode Zn/Zn2+ (1.0 M) //Cu2+ (1.0 M)/Cu Voltmeter: Measures the cell potential usually in volts Voltaic Cell Animation Voltaic cell animation http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/voltaicCellEMF.html Example: The following redox reaction is spontaneous Cr2O72-(aq) + 14H+(aq) + 6I-(aq) 2Cr3+(aq) + 3I2(s) + 7H2O(l) A solution containing K2Cr2O7 and H2SO4 is poured into one beaker and a solution of KI is poured into another beaker. A salt bridge is used to join the beakers. A metallic conductor that will not react with either solution is suspended in each solution and the two conductors are connected with wires through a voltmeter. Indicate the reaction occurring at the anode, the reaction at the cathode, the direction of electron and the ion migrations, and the signs of the electrodes Cell EMF (Electromotive Force) The potential difference between the two electrodes of a voltaic cell which provides the driving force that pushes the electrons through the external circuit EMF of a cell (Ecell) is also called cell potential Measured in volts, so called cell voltage Must be positive value if spontaneous reaction Standard Reduction Potential (Eored) The tendency of a given half cell to occur as a reduction The most positive value on the reduction potential table will be the reduction half reaction Standard Cell Potential (Eocell) The standard cell potential is the measured cell potential when the ion concentrations in the half cells are 1 M, any gases are at a pressure of 1 atm and the temperature is 25oC. Ecell = Ered + Eox Positive values of E mean spontaneous Example: A voltaic cell is based on the following two standard half reactions: Cd2+(aq) + 2 e- Cd(s) Sn2+(aq) + 2 e- Sn(s) By using the data in appendix E, determine the half reactions that occur at the cathode and the anode and the standard cell potential Practice: A voltaic cell is to be designed at standard state conditions using the following two half reactions. Mn2+(aq) + 2e- Mn(s) Eored = -1.18 V Cr3+(aq) + 3e- Cr(s) Eored = -0.74 V Write the chemical equation describing the spontaneous cell reaction and calculate Eocell. Draw a figure representing the voltaic cell. Label all components. Practice Determine which one of the three metals Zn, Fe, or Na is the most active metal given the following data: Na+(aq) + e- Na(s) Eored = -2.71 V Zn2+(aq) + 2e- Zn(s) Eored = -0.76 V Fe2+(aq) + 2e- Fe(s) Eored = - 0.44 V Oxidizing Agents: Reduced in a redox reaction Large positive E values on reduction potential table MnO4-(aq) Mn2+(aq) (in acidic) MnO4-(aq) MnO2 (s) (in basic) Cr2O72-(aq) Cr3+(aq) (in acidic) CrO42-(aq) Cr(OH)3(s) (in basic) All strong nonmetals are excellent oxidizing agents O2(g) O2- (combustion if fast, corrosion if slow) Cl2(g) Cl- (free halogens reduce to halide anions) Reducing Agents: Oxidized in a redox reaction More negative E values on reduction potential table HSO3-(aq) SO42(in acidic) SO32-(aq) SO42(in basic) S2O82-(aq) SO42(in acidic with Cl2) S2O82-(aq) S4O62(in acidic with I2) Fe2+(aq) or Sn2+(aq) Fe3+(aq) or Sn4+(aq) Also, all strong metals are excellent reducing agents Na(s) Na+ Ba(s) Ba2+ (free metals oxidize to metal cations) Spontaneity of Redox Reactions Ecell = Ered + Eox Positive value of Ecell mean spontaneous ∆G = - nFE n = # electrons transferred, E is emf (voltage) F (Faraday’s constant) = 96,500 C/mol = 96,500 J/V-mol Nernst Equation: Relates cell emf to concentration E = Eo – (RT/nF)lnQ E =Eo – (2.303 RT/ nF)log Q E =Eo – (0.0592 / n)log Q Q is reaction quotient F is Faraday’s constant N is # electrons transferred Cell EMF and Chemical Equilibrium EMF of a voltaic cell drops as it discharges As reactants convert to products the value of Q increases so the value of E decreases until E =0 If E = 0 then ∆G = 0 and the system is at equilibrium and Q = Keq Log Keq= nEo / 0.0592 Batteries A battery is a portable, self-contained electrochemical power source of one or more voltaic cells. Lead – Acid Battery Automotive battery uses 6 voltaic cells in series Cathode is PbO2 and anode is Pb Both electrodes are immersed in sulfuric acid Alkaline Battery Anode is powdered Zn in contact with KOH Cathode is graphite and MnO2 Nickel-cadmium, Nickel-metal halide and lithium-ion batteries Light weight rechargeable Cd is anode and NiO(OH)(s) is cathode Corrosion Reactions: Spontaneous redox reactions that lead to the corrosion of metals Lead Battery Recycling Lead-Acid Battery Lead-Acid Battery Alkaline Battery Nickel-Cadmium Battery Cathodic Protection Prevents corrosion of iron Make the metal the cathode in an electrochemical cell Sacrificial anode is the metal that will oxidize Choose an anode metal with a reduction potential that is less than that of the metal to be protected. These metals are more easily oxidized Electrolysis It is possible to use electricity to cause nonspontaneous redox reactions to occur These are called electrolysis reactions and they take place in electrolytic cells Electrolysis of molten salts is an important process for the production of active metals such as sodium. w = -nFE 1 F = 96500 C/mol eCouloumbs = amperes X seconds watt (W) is a unit of electrical power 1 W = 1 J/s If given current and time find quantity of charge (coulombs) find moles of electrons (faradays) find moles of substance oxidized or reduced find grams of substance Electrolytic cell with salt bridge Electrolysis http://www.cressex.bucks.sch.uk//department s/Sc/revision%20movies/electrolysis.htm electrolysis of copper (II) chloride Examples: What mass of zinc metal is produced at the cathode in an electrolysis cell when a constant current of 10.00 amp is passed for 1 hr? Calculate the mass of Mg produced in an electrolytic cell if 2.67 x 102 kwh of electricity is passed through a solution of MgCl2 at 4.20 volts.