chapter 18 electrochemistry. redox reaction elements change oxidation number e.g., single...
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Chapter 18Electrochemistry
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Redox Reaction• Elements change oxidation number
e.g., single displacement, and combustion, some synthesis and decomposition
• Oxidation--oxidation number increases
• Reduction--oxidation number decreases Both must occur in a reaction--two half reactions
• oxidizing agent is reactant molecule that causes oxidation contains element reduced
• reducing agent is reactant molecule that causes reduction contains the element oxidized
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Rules for Assigning Oxidation States1. Free elements have an oxidation state = 0
2. Monatomic ions have an oxidation state equal to their charge.
3. The sum of the oxidation states of all the atoms in a compound is 0.
4. The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.
5. The oxidation number of fluorine is always -1 in compounds with other elements.
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Rules for Assigning Oxidation States
6. Chlorine, bromine and iodine always have oxidation numbers of -1 except when bonded to O or F.
7. The oxidation number of oxygen is almost always -2; the oxidation number of hydrogen is almost always +1.
Exceptions:
--When oxygen is in the form of a peroxide (O22-), the oxidation number is -1.
--When hydrogen forms a binary compound with a metal, the oxidation number is -1 and the compound is called a hydride.
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Oxidation and Reduction
• Oxidation occurs when an atom’s oxidation state increases during a reaction
• Reduction occurs when an atom’s oxidation state decreases during a reaction
CH4 + 2 O2 → CO2 + 2 H2O-4 +1 0 +4 –2 +1 -2
oxidationreduction
Reducing agent Oxidizing agent
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Identify the element that is oxidized and the element that is reduced in each of the following reactions. What is the oxidizing and the reducing agent in each reaction?
3 H2S + 2 NO3– + 2 H+ S + 2 NO + 4 H2O
MnO2 + 4 HBr MnBr2 + Br2 + 2 H2O
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Common Oxidizing AgentsOxidizing Agent Product when Reduced
O2 O-2
H2O2 H2O
F2, Cl2, Br2, I2 F-1, Cl-1, Br-1, I-1
ClO3-1 (BrO3
-1, IO3-1) Cl-1, (Br-1, I-1)
H2SO4 (conc) SO2 or S or H2S
SO3-2 S2O3
-2, or S or H2S
HNO3 (conc) or NO3-1 NO2, or NO, or N2O, or N2, or NH3
MnO4-1 (base) MnO2
MnO4-1 (acid) Mn+2
CrO4-2 (base) Cr(OH)3
Cr2O7-2 (acid) Cr+3
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Common Reducing AgentsReducing Agent Product when Oxidized
H2 H+1
H2O2 O2
I-1 I2
NH3, N2H4 N2
S-2, H2S S
SO3-2 SO4
-2
NO2-1 NO3
-1
C (as coke or charcoal) CO or CO2
Fe+2 (acid) Fe+3
Cr+2 Cr+3
Sn+2 Sn+4
metals metal ions
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Balancing Redox Reactions
1. Assign oxidation numbers--determine element oxidized and element reduced
2. Separate the reaction into oxidation and reduction half- reactions.3. Balance half-reactions by mass
a. First balance elements other than H and Ob. Balance O using H2O c. Balance H using H+
4. Balance each half-reaction by charge by adding electrons to the reactants side of the reduction and the product side of the oxidation.
5. Multiply half-reactions by integers to make # electrons the same in both half-reactions
° Add half-reactions and cancel the electrons to produce a balanced equation.
° For reactions that occur in acidic solutions, skip to step 9.° For reactions that occur in basic solutions, add the same # of OH- as
H+ to both sides of the equation. 9. Check that reaction is balanced for mass and charge.
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Practice - Balance the Equation H2O2 + KI + H2SO4 K2SO4 + I2 + H2O
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Practice - Balance the Equation H2O2 + KI + H2SO4 K2SO4 + I2 + H2O+1 -1 +1 -1 +1 +6 -2 +1 +6 -2 0 +1 -2
oxidationreduction
ox: 2 I-1 I2 + 2e-1
red: H2O2 + 2e-1 + 2 H+ 2 H2Otot 2 I-1 + H2O2 + 2 H+ I2 + 2 H2O
1 H2O2 + 2 KI + H2SO4 K2SO4 + 1 I2 + 2 H2O
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Electric Current Flowing Directly Between Atoms
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Redox Reactions & Current
• Redox reactions involve the transfer of electrons from one substance to another.
• Therefore, redox reactions have the potential to generate an electric current.
• In order to harness the energy produced by moving electrons, we need to separate the half reactions.
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Voltaic (or Galvanic) Cell
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Electrochemical Cells• Electrochemistry -- the study of redox reactions
that produce or require an electric current.
• The conversion between chemical energy and electrical energy is carried out in an electrochemical cell
• Spontaneous redox reactions take place in a voltaic cell (galvanic cell).
• Non-spontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy.
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Electrodes• Anode
electrode where oxidation occurs
• Cathodeelectrode where reduction occurs
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Voltaic (Galvanic) Cell
the salt bridge is required to complete the circuit and maintain charge balance
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Current and Voltage• Current is the number of electrons that flow through the
system per secondunit = Ampere1 A of current = 1 Coulomb/second1 A = 6.242 x 1018 electrons/secondElectrode surface area dictates the number of electrons that can
flow• Potential difference is the difference in potential energy
between the reactants and products (between electrodes)unit = Volt1 V of force = 1 J (of energy)/Coulomb (of charge)The voltage that drive electrons through the external circuitAmount of force pushing the electrons through the wire is called
the electromotive force, emf
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Cell Potential• The difference in potential energy between the
electrodes in a voltaic cell is called the cell potential
• The cell potential depends on the relative ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode
• The cell potential under standard conditions is called the standard emf, E°cell
25°C, 1 atm for gases, 1 M concentration of solutionsum of the cell potentials for the half-reactions
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Cell Notation• Shorthand description of Voltaic cell
electrode | electrolyte || electrolyte | electrode• Oxidation half-cell on left, reduction half-cell on the
right• | = phase barrier
if multiple electrolytes in same phase, a comma is used rather than |
often use an inert electrode
• || = salt bridge
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Fe(s) | Fe2+(aq) || MnO4(aq), Mn2+(aq), H+(aq) | Pt(s)
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Standard Reduction Potential• The cell potential cannot be measured for a
half-reaction
• We need to compare them to an arbitrary standard.
• We select as a standard half-reaction the reduction of H+ to H2 (or the oxidation of H2 to H+) standard hydrogen electrode, SHEStandard conditions [H+]=1M, 25˚CPotential difference (E˚)= 0V
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Half-Cell Potentials• SHE reduction potential is defined to be exactly 0 V
• Reduction potentials are tabulated for many half-reactions
• Change sign to get oxidation potential
• E°cell = E°oxidation + E°reduction
E°oxidation = E°reduction
When adding E° values for the half-cells, do not multiply the half-cell E° values, even if you need to multiply the half-reactions to balance the equation
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Calculating E˚ for a cell reaction
Calculate E˚ for the reaction:
Cu2+ (aq) + Fe2+(aq) Fe2+(aq) + Cu(s)
1) Break up the redox reaction into half-reactions
2) Find the reduction potential (E˚)for each half-reaction
3) Change the sign on E˚ for the oxidation half-reaction
• Add E˚(anode)+ E˚(cathode)
• If the result is negative, the reaction occurs in the opposite direction.