Electrochemistry

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Electrochemistry

• Chapter 17

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Electrochemistry

• The study of the interchange of chemical and
  electrical energy.

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What are some examples of Redox Reactions?

• What are some examples of redox reactions?
• forest fire
• rusting steel
• combustion in auto engine
• metabolism of food in the body

• What everday uses depend on redox reactions?
• starting a car
• calculator
• digital watch
• portable radio
• portable CD player

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Review of Terms

• oxidation-reduction (redox) reaction involves a
  transfer of electrons from the reducing agent to
  the oxidizing agent.
• oxidation loss of electrons
• reduction gain of electrons

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OIL RIG

• Oxidation Is Loss.
• Reduction Is Gain.

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Redox

• Oxidation is the loss of electrons--oxidation
  number becomes more positive.
• Reduction is the gain of electrons--oxidation
  number becomes more negative.

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Redox

• metals nonmetals
• oxidation reduction
• reducing agents oxidizing agents
• metal ions nonmetal ions
• reduction oxidation
• oxidizing agents reducing agents

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Redox Reactions

• Loss and gain of electrons must be simultaneous.
• Loss and gain of electrons must be equal.
• Why must the loss and gain of electrons be equal?

Law of Conservation of Matter
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Redox Reactions

• Redox reactions are reactions in which electrons
  are transferred.
• Decomposition and synthesis reactions may be
  redox.
• Single replacement reactions are always redox.
• Double replacement reactions are never redox.
• Combustion reactions are always redox.

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Identifying Oxidation Reduction in a Reaction

• Identify the element which is oxidized and the
  one which is reduced.
• 1. 2Mg(s) O2(g) ---gt 2MgO(s)
• 2. 2Al(s) 3I2(s) ---gt 2AlI3(s)
• 3. 2Cu(s) O2(g) ---gt 2CuO(s)
• 4. 2Cs(s) F2(g) ---gt 2CsF(s)

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Half-Reactions

• The overall reaction is split into two
  half-reactions, one involving oxidation and one
  reduction.
• 8H MnO4? 5Fe2 ? Mn2 5Fe3 4H2O
• Reduction 8H MnO4? 5e? ? Mn2 4H2O
• Oxidation 5Fe2 ? 5Fe3 5e?

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Rules for Assigning Oxidation States

• 1. Oxidation state of an atom in an element 0
• 2. Oxidation state of monatomic element charge
• 3. Oxygen ?2 in covalent compounds (except in
  peroxides where it ?1)
• 4. H 1 in covalent compounds
• 5. Fluorine ?1 in compounds
• 6. Sum of oxidation states 0 in compounds
  Sum of oxidation states charge of the ion

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Charges Oxidation States

• Oxidation states are written as 2.
• Charges are written 2.

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Determining Oxidation States

• SF6 NO3-
• 6 -6 0 5 -6 -1 (-1 for
  each F) (-2 for each O)

The most electronegative element is assigned a
negative oxidation number--see electronegativity
chart on page 354.
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Determining Oxidation States
oxidation number rule 0 1 1,
-1 2 4, -2 3 1, -1
3 -3, 1, -1 5, 4, 7, 2

• substance
• Na(s)
• NaF(s)
• SO2(g)
• H2O2
• NH4Cl(s)

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Redox

• Oxidizing agent is the electron acceptor--usually
  a nonmetal.
• Reducing agent is the electron donor--usually a
  metal.
• CH4(g) 2O2(g) ----gt CO2(g) 2HOH(g)
• Carbon is oxidized.
• Oxygen is reduced.

-4 1 0 4 -2
1-21
CH4 is the reducing agent. O2 is the oxidizing
agent.
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Redox Reactions

• Identify the substance oxidized and the substance
  reduced as well as the oxidizing and reducing
  agents.
• PbO(s) CO(g) ---gt Pb(s) CO2(g)
• oxidized
• reduced
• oxidizing agent
• reducing agent

2 -2 2 -2 0 4 -2
Carbon Lead PbO CO
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Redox Reactions

• Identify the substance oxidized and the substance
  reduced as well as the oxidizing and reducing
  agents.
• 2PbS(s) 3O2(g) ---gt 2PbO(s) 2SO2(g)
• oxidized
• reduced
• oxidizing agent
• reducing agent

2 -2 0 2 -2
4 -2
sulfur oxygen O2 PbS
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Balancing by Half-Reaction Method

• 1. Write separate reduction, oxidation reactions.
• 2. For each half-reaction
• - Balance elements (except H, O)
• - Balance O using H2O
• - Balance H using H
• - Balance charge using electrons

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Balancing by Half-Reaction Method (continued)

• 3. If necessary, multiply by integer to equalize
  electron count.
• 4. Add half-reactions cancel identical species.
• 5. Check that both elements and charges are
  balanced.

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Balancing By Half-Reaction Acidic Solution

• H(aq) Cr2O72-(aq) C2H5OH(l) ---gt Cr3(aq)
  CO2(g) HOH(l)
• Red Cr2O72-(aq) ---gt Cr3(aq)
• Ox C2H5OH(l) ---gt CO2(g)
• Red 2(6e- 14H(aq) Cr2O72-(aq) ---gt
  2Cr3(aq) 7HOH(l))
• Ox C2H5OH(l) 3HOH(l) ---gt 2CO2(g) 12H(aq)
  12e-
• 16H(aq) 2Cr2O72-(aq) C2H5OH(l) ---gt
  4Cr3(aq) 11HOH(l) 2CO2(g)
• 12 12

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Redox Reactions

• Always add electrons to the side of the
  half-reaction with excess positive charge!

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Balancing By Half-Reaction Acidic Solution

• Cu(s) HNO3(aq) ---gt Cu(NO3)2(aq) NO(g)
  HOH(l)
• Ox Cu(s) HNO3(aq) ---gt Cu(NO3)2(aq)
• Red HNO3(aq) ---gt NO(g)
• Ox 3(Cu(s) 2HNO3(aq) ---gt Cu(NO3)2(aq)
  2H(aq) 2e-)
• Red 2(3e- 3H(aq) HNO3(aq) ---gt NO(g)
  2HOH(l) )
• 3Cu(s) 8HNO3(aq) ---gt 3Cu(NO3)2(aq) 4HOH(l)
  2NO(g)
• 0 0

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Half-Reaction Method - Balancing in Base

• 1. Balance as in acid.
• 2. Add OH? that equals H ions (both sides!)
• 3. Form water by combining H, OH?.
• 4. Check elements and charges for balance.

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Balancing By Half-Reaction Basic Solution

• Ag(s) CN-(aq) O2(g) ---gt Ag(CN)2-(aq)(Basic)
• Ox CN-(aq) Ag(s) ---gt Ag(CN)2-(aq)
• Red O2(g) ---gt
• Ox 4(2CN-(aq) Ag(s) ---gt Ag(CN)2-(aq) e-)
• Red O2(g) 4H(aq) 4e- ---gt 2HOH(l)
• 8CN-(aq) 4Ag(s) O2(g) 4H(aq) ---gt
  4Ag(CN)2-(aq) 2HOH(l)

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Balancing By Half-Reaction Basic Solution

• 8CN-(aq) 4Ag(s) O2(g) 4H(aq) 4OH-(aq)
• ---gt 4Ag(CN)2-(aq) 2HOH(l)
  4OH-(aq)
• 8CN-(aq) 4Ag(s) O2(g) 4HOH(l)
• ---gt 4Ag(CN)2-(aq) 2HOH(l)
  4OH-(aq)
• 8- 8-

2HOH(l)
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Galvanic Cell

• A device in which chemical energy is changed to
  electrical energy.

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Galvanic Cell

• A device in which chemical energy is changed to
  electrical energy. The basic parts are
• anode
• cathode
• electrochemical solution
• porous disk or salt bridge

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Anode and Cathode

• OXIDATION occurs at the ANODE.
• REDUCTION occurs at the CATHODE.
• AN OX RED CAT

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Electrons are transferred at the interface
between the electrodes and the solution. Porous
disk allows ion flow.
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A galvanic cell (Daniell Cell) involving Zn and
Cu electrodes. This cell was the energy source
for telegraphy during the War Between the States.
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Cu-Zn (Daniell Cell) on the microscopic level.
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Zinc electrode compared to a Standard Hydrogen
Electrode (SHE). The Zn has a potential of 0.76
V.
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Standard Reduction Potentials

• The E? values corresponding to reduction
  half-reactions with all solutes at 1M and all
  gases at 1 atm.
• Cu2 2e? ? Cu E? 0.34 V vs. SHE
• SO42? 4H 2e? ? H2SO3 H2O
• E? 0.20 V vs. SHE
• SHE Standard Hydrogen Electrode

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Cell Potential

• Cell Potential or Electromotive Force (emf) The
  pull or driving force on the electrons.

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Cell Potential Calculations

• To Calculate cell potential using Standard
  Reduction Potentials
• 1. One reaction and its cell potentials sign
  must be reversed--it must be chosen such that the
  overall cell potential is positive.
• 2. The half-reactions must often be multiplied
  by an integer to balance electrons--this is not
  done for the cell potentials.

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Cell Potential Calculations Continued

• Fe3(aq) Cu(s) ----gt Cu2(aq) Fe2(aq)
• Fe3(aq) e- ----gt Fe2(aq) Eo 0.77 V
• Cu2(aq) 2 e- ----gt Cu(s) Eo 0.34 V
• Reaction 2 must be reversed.

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Cell Potential Calculations Continued

• 2 (Fe3(aq) e- ----gt Fe2(aq)) Eo 0.77
  V
• Cu(s) ----gt Cu2(aq) 2 e- Eo - 0.34
  V
• 2Fe3(aq) Cu(s) ----gt Cu2(aq) 2Fe2(aq)
• Eo 0.43 V

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Ion selective electrodes are glass electrodes
that measures a change in potential when H
varies. Used to measure pH.
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Batteries

• A battery is a galvanic cell or, more commonly, a
  group of galvanic cells connected in series.

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A lead storage battery consists of a lead anode,
lead dioxide cathode, and an electrolyte of 38
sulfuric acid.
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Lead Storage Battery

• Anode reaction
• Pb(s) H2SO4(aq) ---gt PbSO4(aq) 2H(aq) 2e-
• Cathode reaction
• PbO2(s) H2SO4(aq) 2e- 2H(aq) ---gt
  PbSO4(aq) 2HOH(l)
• Overall reaction
• Pb(s) PbO2(s) 2H2SO4(aq) ---gt PbSO4(aq)
  2HOH(l)

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Common dry cell and its components.
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Mercury battery used in calculators.
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Fuel Cells

• . . . galvanic cells for which the reactants are
  continuously supplied.
• 2H2(g) O2(g) ? 2H2O(l)
• anode 2H2 4OH? ? 4H2O 4e?
• cathode 4e? O2 2H2O ? 4OH?

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emf and Work
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Free Energy and Cell Potential

• ?G? ?nFE?
• n number of moles of electrons
• F Faraday 96,485 coulombs per mole of
  electrons

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Concentration Cell

• . . . a cell in which both compartments have the
  same components but at different concentrations.

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The Nernst Equation

• We can calculate the potential of a cell in which
  some or all of the components are not in their
  standard states.

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Calculation of Equilibrium Constants for Redox
Reactions

• At equilibrium, Ecell 0 and Q K.

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Stoichiometry of Electrolysis

• How much chemical change occurs with the flow of
  a given current for a specified time?
• current and time ? quantity of charge ?
• moles of electrons ? moles of analyte ?
• grams of analyte
• 1 amp 1 C/s

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Electrolytic Calculations

• How many grams of copper can be plated out when a
  current of 10.0 amps is passed through a
  Cu2solution for 30.0 minutes?
• (10.0 C/s)(30.0 min)(60 s/1 min)(1 mol e-/96,485
  C)
• (1 mol Cu/2 mole e-)(63.5 g/1 mol) 5.94 g Cu

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Electrolytic Calculations

• How long must a current of 5.00 A be applied to a
  solution of Ag1 to produce 10.5 g of silver?
• (10.5 g Ag)(1 mol/107.86 g)(1 mol e-/1 mol Ag)
• (96,485 C/1 mole e-)(1 s/5.00 C)(1 min/60.0s)
• 
  31.3 min

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Electrochemical lt---gt Electrolytic

• Electrochemical lt---gt Electrolytic
• Spontaneous lt---gt Nonspontaneous
• Energy released lt---gt Energy absorbed
• Cu2(aq) Mg(s) lt---gt Cu(s) Mg2(aq)
• Electrochemical cell -- chemical energy to
  electrical energy.
• Electrolytic cell -- electrical energy to
  chemical energy.

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Corrosion

• Some metals, such as copper, gold, silver and
  platinum, are relatively difficult to oxidize.
  These are often called noble metals.
• About 1/5 of all iron and steel produced each
  year is used to replace rusted metal.

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Self-protecting Metals

• Some metals such as aluminum, copper, and silver
  form a protective coating that keeps them from
  corroding further.
• The protective coating for iron and steel flakes
  away opening new layers of metal to corrosion.

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Corrosion of Iron
Iron is oxidized at the anodic reaction and
oxygen is reduced at the cathodic reaction.
Dissolved ions are necessary to transfer
electrons between the anodic and cathodic areas.
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Prevention of Corrosion

• Coating--painting or applying oil to keep out
  oxygen and moisture.
• Galvanizing--dipping a metal in a more active
  metal -- galvanized steel bucket.
• Alloying -- mixing metals with iron to prevent
  corrosion -- stainless steel.
• Cathodic protection -- attaching a more active
  metal. Serves as sacrificial metal--used to
  protect ships, gas lines, and gas tanks.

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Electrolysis

• . . . forcing a current through a cell to produce
  a chemical change for which the cell potential is
  negative.

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Electrolysis of Water

• The electrolysis of water is
• DC
• 2HOH(l) -----gt 2H2(g) O2(g)
• Why would this process be important on a
• submarine?

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Schematic of Hall-Heroult process. Molten sinks
and is tapped off at the bottom.
Bauxite-cryolite mixture floats on top and is
electrolyzed by the carbon electrodes.
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Figure 17.11 Charles Martin Hall (1863-1914)
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Why did Napoleon III of France served his most
honored guests with aluminum silverware?