Presentation on theme: "Electrochemistry Electrochemical reactions are oxidation-reduction reactions. The two parts of the reaction are physically separated. The oxidation reaction."— Presentation transcript:
1ElectrochemistryElectrochemical reactions are oxidation-reduction reactions.The two parts of the reaction are physically separated.The oxidation reaction occurs in one cell.The reduction reaction occurs in the other cell.There are two kinds electrochemical cells.Electrochemical cells containing nonspontaneous chemical reactions are called electrolytic cells.Electrochemical cells containing spontaneous chemical reactions are called voltaic or galvanic cells.
2Electrical Conduction Metals conduct electric currents well in a process called metallic conduction.In metallic conduction there is electron flow with no atomic motion.In ionic or electrolytic conduction ionic motion transports the electrons.Positively charged ions, cations, move toward the negative electrode, cathode.Negatively charged ions, anions, move toward the positive electrode, anode.
3ElectrodesThe following convention for electrodes is correct for either electrolytic or voltaic cells:The cathode is the electrode at which reduction occurs.The cathode is negative in electrolytic cells and positive in voltaic cells.The anode is the electrode at which oxidation occurs.The anode is positive in electrolytic cells and negative in voltaic cells.
4The Half-Reaction Method Half reaction method rules:Write the unbalanced reaction.Break the reaction into 2 half reactions:One oxidation half-reaction andOne reduction half-reactionEach reaction must have complete formulas for molecules and ions.Mass balance each half reaction by adding appropriate stoichiometric coefficients. To balance H and O we can add:H+ or H2O in acidic solutions.OH- or H2O in basic solutions.Charge balance the half reactions by adding appropriate numbers of electrons.Electrons will be products in the oxidation half-reaction.Electrons will be reactants in the reduction half-reaction.Multiply each half reaction by a number to make the number of electrons in the oxidation half-reaction equal to the number of electrons reduction half-reaction.Add the two half reactions.Eliminate any common terms and reduce coefficients to smallest whole numbers.3837393939
5The Half-Reaction Method Tin (II) ions are oxidized to tin (IV) by bromine. Use the half reaction method to write and balance the net ionic equation.4039414141
6The Half-Reaction Method Dichromate ions oxidize iron (II) ions to iron (III) ions and are reduced to chromium (III) ions in acidic solution. Write and balance the net ionic equation for the reaction.4443454545
7The Half-Reaction Method In basic solution hydrogen peroxide oxidizes chromite ions, Cr(OH)4-, to chromate ions, CrO42-. The hydrogen peroxide is reduced to hydroxide ions. Write and balance the net ionic equation for this reaction.4847505049
8The Half-Reaction Method When chlorine is bubbled into basic solution, it forms hypochlorite ions and chloride ions. Write and balance the net ionic equation.This is a disproportionation redox reaction. The same species, in this case Cl2, is both reduced and oxidized.5049525251
9Stoichiometry of Redox Reactions Just as we have done stoichiometry with acid-base reactions, it can also be done with redox reactions.What volume of M KMnO4 is required to oxidize 35.0 mL of M HCl? The balanced reaction is:5251545453
10Stoichiometry of Redox Reactions A volume of 40.0 mL of iron (II) sulfate is oxidized to iron (III) by 20.0 mL of M potassium dichromate solution. What is the concentration of the iron (II) sulfate solution? The balanced equation is:5453565655
11Voltaic or Galvanic Cells Electrochemical cells in which a spontaneous chemical reaction produces electrical energy.Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference.Examples of voltaic cells include:
12The Construction of Simple Voltaic Cells Voltaic cells consist of two half-cells which contain the oxidized and reduced forms of an element (or other chemical species) in contact with each other.A simple half-cell consists of:A piece of metal immersed in a solution of its ions.A wire to connect the two half-cells.And a salt bridge to complete the circuit, maintain neutrality, and prevent solution mixing.
13The Zinc-Copper Cell Cell components for the Zn-Cu cell are: A metallic Cu strip immersed in 1.0 M copper (II) sulfate.A metallic Zn strip immersed in 1.0 M zinc (II) sulfate.A wire and a salt bridge to complete circuitThe cell’s initial voltage is 1.10 volts
14The Zinc-Copper CellIn all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).
15The Zinc-Copper CellThere is a commonly used short hand notation for voltaic cells.The Zn-Cu cell provides a good example.
16The Copper - Silver Cell Cell components:A Cu strip immersed in 1.0 M copper (II) sulfate.A Ag strip immersed in 1.0 M silver (I) nitrate.A wire and a salt bridge to complete the circuit.The initial cell voltage is 0.46 volts.
17The Copper - Silver Cell Compare the Zn-Cu cell to the Cu-Ag cellThe Cu electrode is the cathode in the Zn-Cu cell.The Cu electrode is the anode in the Cu-Ag cell.Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.
18The Copper - Silver Cell These experimental facts demonstrate that Cu2+ is a stronger oxidizing agent than Zn2+.In other words Cu2+ oxidizes metallic Zn to Zn2+.Similarly, Ag+ is is a stronger oxidizing agent than Cu2+.Because Ag+ oxidizes metallic Cu to Cu 2+.If we arrange these species in order of increasing strengths, we see that:
19Easiest to reduce isStrongest Oxidizing AgentAg +Cu 2+Zn 2+(find order using left hand side of table)Least likely to be reduced isStrongest Reducing AgentZnCuAg(find order using right hand side of table)
20Standard Electrode Potential To measure relative electrode potentials, we must establish an arbitrary standard.That standard is the Standard Hydrogen Electrode (SHE).The SHE is assigned an arbitrary voltage of … V
21The Zinc-SHE Cell For this cell the components are: A Zn strip immersed in 1.0 M zinc (II) sulfate.The other electrode is the Standard Hydrogen Electrode.A wire and a salt bridge to complete the circuit.The initial cell voltage is volts.
22The Zinc-SHE Cell The cathode is the Standard Hydrogen Electrode. In other words Zn reduces H+ to H2.The anode is Zn metal.Zn metal is oxidized to Zn2+ ions.
23The Copper-SHE Cell The cell components are: A Cu strip immersed in 1.0 M copper (II) sulfate.The other electrode is a Standard Hydrogen Electrode.A wire and a salt bridge to complete the circuit.The initial cell voltage is volts.
24The Copper-SHE Cell In this cell the SHE is the anode The Cu2+ ions oxidize H2 to H+.The Cu is the cathode.The Cu2+ ions are reduced to Cu metal.
25Primary Voltaic CellsAs a voltaic cell discharges, its chemicals are consumed.Once the chemicals are consumed, further chemical action is impossible.The electrodes and electrolytes cannot be regenerated by reversing current flow through cell.These cells are not rechargable.
26The Dry CellOne example of a dry cell is flashlight and radio batteries.The cell’s container is made of zinc which acts as an electrode.A graphite rod is in the center of the cell which acts as the other electrode.The space between the electrodes is filled with a mixture of:ammonium chloride, NH4Clmanganese (IV) oxide, MnO2zinc chloride, ZnCl2and a porous inactive solid.
27The Dry Cell The anode reaction is: The graphite rod is the positive electrode (cathode).Ammonium ions from the NH4Cl are reduced at the cathode.
28The Dry CellThe other components in the cell are included to remove the byproducts of the reaction.MnO2 prevents H2 from collecting on graphite rod.At the anode, NH3 combines with Zn2+ to form a soluble complex and removing the Zn2+ ions from the reaction.
29The Dry CellAlkaline dry cells are similar to ordinary dry cells except that KOH, an alkaline substance, is added to the mixture.Half reactions for an alkaline cell are:
30Secondary Voltaic Cells Secondary cells are reversible, rechargeable.The electrodes in a secondary cell can be regenerated by the addition of electricity.These cells can be switched from voltaic to electrolytic cells.One example of a secondary voltaic cell is the lead storage or car battery.The Lead Storage BatteryIn the lead storage battery the electrodes are two sets of lead alloy grids (plates).Holes in one of the grids are filled with lead (IV) oxide, PbO2.The other holes are filled with spongy lead.The electrolyte is dilute sulfuric acid.
31The Lead Storage Battery Diagram of the lead storage battery.
32The Lead Storage Battery As the battery discharges, spongy lead is oxidized to lead ions and the plate becomes negatively charged.The Pb2+ ions that are formed combine with SO42- from sulfuric acid to form solid lead sulfate on the Pb electrode.
33The Lead Storage Battery The net reaction at the anode during discharge is:As the cell discharges, the cathode reaction is:The cell reaction for a discharging lead storage battery is:As the cell discharges, the cathode reaction is:
34The Lead Storage Battery What happens at each electrode during recharging?At the lead (IV) oxide, PbO2, electrode, lead ions are oxidized to lead (IV) oxide.The concentration of the H2SO4 decreases as the cell discharges.Recharging the cell regenerates the H2SO4.
35The Nickel-Cadmium (Nicad) Cell Nicad batteries are the rechargeable cells used in calculators, cameras, watches, etc.As the battery discharges, the half-reactions are:
36The Hydrogen-Oxygen Fuel Cell Fuel cells are batteries that must have their reactants continuously supplied in the presence of appropriate catalysts.A hydrogen-oxygen fuel cell is used in the space shuttleThe fuel cell is what exploded in Apollo 13.Hydrogen is oxidized at the anode.Oxygen is reduced at the cathode.
37The Hydrogen-Oxygen Fuel Cell Notice that the overall reaction is the combination of hydrogen and oxygen to form water.The cell provides a drinking water supply for the astronauts as well as the electricity for the lights, computers, etc. on board.Fuel cells are very efficient.Energy conversion rates of 60-70% are common!
38Uses of Standard Electrode Potentials Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials.Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials.Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.For example, the half-reaction for the standard potassium electrode is:
39Uses of Standard Electrode Potentials Compare the potassium half-reaction to fluorine’s half-reaction:
40Uses of Standard Electrode Potentials Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously.Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn2+ ions oxidize metallic Ag to Ag+ ions?Steps for obtaining the equation for the spontaneous reaction.Choose the appropriate half-reactions from a table of standard reduction potentials.Write the equation for the half-reaction with the more positive E0 value first, along with its E0 value.Write the equation for the other half-reaction as an oxidation with its oxidation potential, i.e. reverse the tabulated reduction half-reaction and change the sign of the tabulated E0.Balance the electron transfer.Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous.
42Electrode Potentials for Other Half-Reactions Will permanganate ions, MnO4-, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution?Follow the steps outlined in the previous slides.Note that E0 values are not multiplied by any stoichiometric relationships in this procedure.
43Electrode Potentials for Other Half-Reactions Will nitric acid, HNO3, oxidize arsenous acid, H3AsO3, in acidic solution? The reduction product of HNO3 is NO in this reaction.
44Effect of Concentrations (or Partial Pressures) on Electrode Potentials The Nernst EquationStandard electrode potentials, those compiled in appendices, are determined at thermodynamic standard conditions.Reminder of standard conditions.
45The Nernst EquationThe value of the cell potentials change if conditions are nonstandard.The Nernst equation describes the electrode potentials at nonstandard conditions.The Nernst equation is:Substitution of the values of the constants into the Nernst equation at 25o C gives:For this half-reaction:The corresponding Nernst equation is:
46The Nernst Equation Substituting E0 into the above expression gives: If [Cu2+] and [Cu+] are both 1.0 M, i.e. at standard conditions, then E = E0 because the concentration term equals zero.
47The Nernst EquationCalculate the potential for the Cu2+/ Cu+ electrode at 250C when the concentration of Cu+ ions is three times that of Cu2+ ions.
48The Nernst EquationCalculate the potential for the Cu2+/Cu+ electrode at 250C when the Cu+ ion concentration is 1/3 of the Cu2+ ion concentration.
49The Nernst EquationCalculate the electrode potential for a hydrogen electrode in which the [H+] is 1.0 x 10-3 M and the H2 pressure is 0.50 atmosphere.
50The Nernst EquationThe Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes.Calculate the initial potential of a cell that consists of an Fe3+/Fe2+ electrode in which [Fe3+]=1.0 x 10-2 M and [Fe2+]=0.1 M connected to a Sn4+/Sn2+ electrode in which [Sn4+]=1.0 M and [Sn2+]=0.10 M . A wire and salt bridge complete the circuit.
51The Nernst EquationCalculate the E0 cell by the usual procedure.
52Counting Electrons: Coulometry and Faraday’s Law of Electrolysis A coulomb is the amount of charge that passes a given point when a current of one ampere (A) flows for one second.1 amp = 1 coulomb/second
53Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Faraday’s Law states that during electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent.This corresponds to the passage of one mole of electrons through the electrolytic cell.
54Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes.
55Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in the previous example.