Presentation on theme: "Unit 11- Redox and Electrochemistry"— Presentation transcript:
1 Unit 11- Redox and Electrochemistry AnodeCathodeElectrochemical cellElectrodeElectrolysisElectrolyteElectrolytic cellHalf-reactionOxidationOxidation numberRedoxReductionSalt bridgeVoltaic cell
2 What’s the point ? REDOX reactions are important in … 30/09/99What’s the point ?REDOX reactions are important in …C3H8O + CrO3 + H2SO4 Cr2(SO4)3 + C3H6O + H2OPurifying metals (e.g. Al, Na, Li)Producing gases (e.g. Cl2, O2, H2)Electroplating metalsElectrical production (batteries, fuel cells)Protecting metals from corrosionBalancing complex chemical equationsSensors and machines (e.g. pH meter)
3 What is redox? Oxidation- loss of electrons by an atom or ion Reduction- gain of electrons by an atom or ion**since one can’t occur without the otherCombine terms to RedoxMnemonic: LEO the lion says GERLose Electrons OxidationGain Electrons Reduction
4 Oxidation numbers On periodic table Determines what is oxidized and reduced in a reactionIf they change it’s a redox reactionWhat type of reaction is this (besides redox)???
5 Assigning Oxidation numbers Identify the formulaIf element is free (uncombined) its ox # is 0Monotomic ions- ox # is same as ion chargeMetals in Groups 1,2 and 3 have ox #’s of +1, +2 and +3 respectivelyFluorine is always -1 in a compoundHydrogen is always +1 unless it’s combined with a metal then it’s -1Oxygen is usually -2, except when combined with a more electronegative element then it’s +2*sum of oxidation #’s in a compound must be 0*sum of oxidation #’s in a polyatomic ion must equal its charge
7 Redox reactionsOnce you determine oxidation numbers you can see what element was oxidized and what was reducedOxidizing agent- substance that was reduced (gained electrons)Reducing agent substance that was oxidized (lost electrons)Oxidizing agent (element reduced) because it gains electrons from oxidized element so it allowed for oxidation to take place because it was a place for the electrons to go
8 Half-reactions Show oxidation or reduction of redox rx Ex: Shows conservation of mass and chargeCharge does not have to be 0
9 Balancing redox rx’sAssign oxidation numbers to determine what is oxidized and what is reduced.Write the oxidation and reduction half-reactions.Balance each half-reaction.Balance charge by adding electrons.Multiply the half-reactions by integers so that the electrons gained and lost are the same
10 Example: Cu + AgNO3 Cu(NO3)2 + Ag Add the half-reactions, subtracting things that appear on both sides.Make sure the equation is balanced according to mass.Make sure the equation is balanced according to charge.Example:Cu + AgNO3 Cu(NO3)2 + Ag
11 Practical use for redox reactions Electrochemical cellsInvolves a chemical reaction and flow of electrons2 types:Voltaic- spontaneousElectrolytic- requires electric current (nonspontaneous)Each have 2 electrodes- site of oxidation and reductionOxidation occurs at the anodeReduction occurs at the cathodeAn Ox Red CatAnode- oxidation, reduction-cathode
12 Voltaic cellsOnce even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop
13 Voltaic cellsTherefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. (completes the circuit)Cations move toward the cathode.Anions move toward the anode.
14 Voltaic CellsIn the cell, then, electrons leave the anode and flow through the wire to the cathode.As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.
15 Voltaic CellsAs the electrons reach the cathode, cations in the cathode solution are attracted to the now negative cathode.The electrons are taken by the cation, and the neutral metal is deposited on the cathode.
16 Activity series helps identify anode and cathode Metal higher on chart- oxized (anode)Metal lower on chart- site of reduction (cathode)
17 Determining electric potential Voltmeter is usedVoltage is compared to the reduction of H which is 0 voltsThe more “+” the reading; reduction is more likelyE is positive- happens naturallyE is negaitve- not spontaneousStandard electrode potentials-Eo= ability for an electrode to gain electrons
18 Reduction potentials for many electrodes has already been measured
19 Cell potentialsAt standard conditions can be determined using this equation:The strongest oxidizers have the most positive reduction potentials.The strongest reducers have the most negative reduction potentials.Ecell= Ered (cathode) − Ered (anode)
20 Cell Potentials Ered = −0.76 V Ered = +0.34 V For the oxidation in this cell,For the reduction,Ered = −0.76 VEred = V
21 Cell Potentials Ecell = Ered (cathode) − (anode) = V − (−0.76 V)= V
22 Examples of Voltaic Cells: Dry Cells Dry cells use two electrodes and a “paste” as an electrolyte.Some pastes are acidic and others are alkaline.Carbon is generally used as the cathode and zinc as the anode.
23 Lead-Acid BatteriesLead-Acid batteries usually contain six cells.(2 V each)The battery contains lead plates, lead oxide plates, dividers, and a sulfuric acid electrolyte.The lead plate is the anode and the lead oxide plate is the cathode.Each cell is connected to form one cathode and one anode on the top or side of the battery.
24 Fuel CellsFuel cells bring in the oxidizing and reducing agents as gasesGraphite is typically the anode and cathode for the reaction which produced electricity.Fuel cells are clean and efficient.
25 Corrosion Corrosion is defined as the disintegration of metals. Corrosion is typically caused by oxygen (O2).A familiar example of corrosion is iron rusting.Corrosion is a result of a redox reaction involving a metal.Iron oxide (rust)
27 Corrosion PreventionTypical corrosion protection involves plating the iron with another metal.The production of steel (iron and carbon) reduces the rate of corrosion of the iron.Aluminum, zinc, titanium are some metals which corrode slowly, or have different properties used to protect iron.
29 Electrolytic Cells Electricity is used to force a chemical reaction ElectrolysisUsed to obtain metals from molten saltsStarting/keeping a car runningPlating metalsWhen starting a car- spont rx occurs in battery to provide electricity to start the carTo keep the car running-alternator (non-spont rx) recharges the batteryAlso making aluminumElectrolysis of water into h and o gas
30 Electroplating Item to be plated is cathode Metal that will plate is anodePut in solution containing ions- electrolyte
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