Download - Unit 11- Redox and Electrochemistry
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Unit 11- Redox and Electrochemistry
• Anode• Cathode• Electrochemical cell• Electrode• Electrolysis• Electrolyte• Electrolytic cell• Half-reaction• Oxidation• Oxidation number
• Redox• Reduction• Salt bridge• Voltaic cell
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What’s the point ?
• Electrical production (batteries, fuel cells)
REDOX reactions are important in …
• Purifying metals (e.g. Al, Na, Li)
• Producing gases (e.g. Cl2, O2, H2)
• Electroplating metals
• Protecting metals from corrosion• Balancing complex chemical equations• Sensors and machines (e.g. pH meter)
C3H8O + CrO3 + H2SO4 Cr2(SO4)3 + C3H6O + H2O
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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 other
– Combine terms to Redox– Mnemonic: LEO the lion says GER
• Lose Electrons Oxidation• Gain Electrons Reduction
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Oxidation numbers• On periodic table• Determines what is oxidized and reduced in a
reaction• If they change it’s a redox reaction
What type of reaction is this (besides redox)???
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Assigning Oxidation numbers• Identify the formula• If element is free (uncombined) its ox # is 0• Monotomic ions- ox # is same as ion charge• Metals in Groups 1,2 and 3 have ox #’s of +1, +2 and +3
respectively• Fluorine is always -1 in a compound• Hydrogen is always +1 unless it’s combined with a metal then
it’s -1• Oxygen 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
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Try these:
• HNO3
• CO2
• K2PtCl6
• PCl5
• H2SO4
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Redox reactions• Once you determine oxidation numbers you
can see what element was oxidized and what was reduced
• Oxidizing agent- substance that was reduced (gained electrons)
• Reducing agent- substance that was oxidized (lost electrons)
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Half-reactions
• Show oxidation or reduction of redox rx• Ex:
• Shows conservation of mass and charge – Charge does not have to be 0
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Balancing redox rx’s• Assign 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
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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.
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Practical use for redox reactions
• Electrochemical cells– Involves a chemical reaction and flow of electrons– 2 types:
• Voltaic- spontaneous• Electrolytic- requires electric current (nonspontaneous)• Each have 2 electrodes- site of oxidation and reduction
– Oxidation occurs at the anode– Reduction occurs at the cathode– An Ox Red Cat– Anode- oxidation, reduction-cathode
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Voltaic cells
• Once 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
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Voltaic cells• Therefore, 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.
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Voltaic Cells• In 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.
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Voltaic Cells• As 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.
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• Activity series helps identify anode and cathode– Metal higher on
chart- oxized (anode)– Metal lower on
chart- site of reduction (cathode)
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Determining electric potential• Voltmeter is used• Voltage is compared to the reduction of H
which is 0 volts• The more “+” the reading; reduction is more
likely
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• Reduction potentials for many electrodes has already been measured
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Cell potentials
• At 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)
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Cell Potentials• For the oxidation in this cell,
• For the reduction,
Ered = −0.76 V
Ered = +0.34 V
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Cell Potentials
Ecell = Ered (cathode) − Ered (anode)
= +0.34 V − (−0.76 V)= +1.10 V
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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.
Examples of Voltaic Cells:
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Lead-Acid Batteries
• Lead-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.
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Fuel Cells
• Fuel cells bring in the oxidizing and reducing agents as gases
• Graphite is typically the anode and cathode for the reaction which produced electricity.
• Fuel cells are clean and efficient.
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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)
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Corrosion con’t…
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Corrosion Prevention
• Typical 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.
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…Corrosion Prevention
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Electrolytic Cells• Electricity is used to force a chemical reaction
– Electrolysis• Used to obtain metals from molten salts
• Starting/keeping a car running• Plating metals
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Electroplating
• Item to be plated is cathode
• Metal that will plate is anode
• Put in solution containing ions- electrolyte
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Electroplating con’t
• Benefits– Resists corrosion– Improves appearance– Cheaper
• Drawbacks– Plating isn’t always even– Can wear off– Solutions are toxic