Chapter 20. A chemical reaction can perform two types of work: 1. Produce a gas to perform PV work...
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Transcript of Chapter 20. A chemical reaction can perform two types of work: 1. Produce a gas to perform PV work...
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Chapter 20
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A chemical reaction can perform two types of work:
1.Produce a gas to perform PV work2.Use movement of electrons from
redox reactions to perform electrical work
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A voltaic (galvanic) cell is a device in which the transfer of electrons takes place through an external pathway rather than directly between reactants.
By physically separating the reduction half of a redox reaction from the oxidation half, we create a flow of electrons through an external circuit.
Used to accomplish electrical work
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Two solid metals that are connected by the external circuit are called electrodes. Anode: Cathode:
Electrodes may or may not participate in the reaction Zn/Cu Pt or other conducting material
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Each of the two components of a voltaic cell is called a half-cell Oxidation half-cell Reduction half-cell
For a voltaic cell to work, the solutions in the two half-cells must remain electrically neutral Need migration of ions
Salt bridge or porous glass barrior
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Voltaic Cells
Anode Acceptor of electrons: Oxidation
Cathode Source of electrons: Reduction
Anions always migrate toward the anode and cations toward the cathode.
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Voltaic (Galvanic) Cell
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The following redox reaction is spontaneous: Cr2O7
2-(aq)+ 6I-(aq) 2Cr3+(aq) + 3I2(s)
A solution containing K2Cr2O7 and H2SO4 is poured into one beaker, and a solution of KI is poured into another. A salt bridge is used to join the beakers. A metallic conductor that will not react with either solution is suspended in each solution, and the two conductors are connected with wires through a voltmeter to detect an electric current. The resultant voltaic cell generates an electric current. Indicate the reaction occurring at the anode, the reaction at the cathode, the direction of electron migration, the direction of ion migration, and the signs of the electrodes.
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Describing a Voltaic Cell
The two half-reactions in a voltaic cell are Zn(s) Zn2+(aq)
+ 2e-ClO3
-(aq) + 6H+(aq) + 6e- Cl-(aq) + 3H2O(l)
(a) Indicate which reaction occurs at the anode and which at the cathode. (b) which electrode is consumed in the cell reaction? (c) Which electrode is positive?
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Cell EMF Under Standard ConditionsWhy do electrons transfer
spontaneously during redox reactions?
Electrons flow from the anode of a voltaic cell to the cathode because of a difference in potential energy. Potential energy higher at the anode
Electrons flow spontaneously toward the electrode with the more positive electrical potential.
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Cell EMF Under Standard Conditions
The difference in potential energy per electrical charge between two electrodes is measured in units of volts.
1V = 1 (J/C) Where V (volts), J (joule), and C
(coulomb)
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Cell EMF Under Standard Conditions
The potential difference between two electrodes provides a driving force that pushes electrons through the external circuit. Electromotive Force (emf)
Emf of a cell is denoted as Ecell (the cell potential)
For spontaneous reactions, the cell potential will be positive
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Standard EMF
Emf depends on The particular cathode and anode half reactions Concentrations of the reactants and products Temperature
Tabulated values of standard reduction potentials denoted Eo
red to calculate Eocell
Eocell = Eo
red (cathode) – Eored (anode)
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Standard Emf
Indirectly measure the standard reduction potential of a half-reaction
Reference point: 2H+ (aq, 1 M) + 2e- H2 (g, 1
atm)
Assigned a standard reduction potential of exactly zero volts
Called a standard hydrogen electrode (SHE)
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Standard Emf
When determining standard reduction potentials from other half-reactions, write the reaction as a reduction even though it is “running in reverse” as an oxidation reaction. Whenever an electrical potential is
assigned to a half-reaction, write the reaction as a reduction.
Eored are intensive properties
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Calculating Eored from
Eocell
For the Zn-Cu2+ voltaic cell, we have
Zn + Cu2+ Zn2+ + Cu Eocell =
1.10V
Given that Eored of Zn2+ to Zn is -0.76
V, calculate the Eored for the
reduction of Cu2+ to Cu
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A voltaic cell is based on the half-reactions:
In+ In3+ + 2e-Br2 + 2e- 2Br-
The standard emf for the cell is 1.46V and Eo
red for the reduction of bromine is +1.06V. Using this information, calculate Eo
red for the reduction of In3+ to In+.
Calculating Eored from
Eocell
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Calculating Eocell from
Eored
Using the standard reduction potentials listed in Table 20.1, calculate the standard emf for the following voltaic cells:
1.Cr2O72- + 14H+ + 6I- 2Cr3+ + 3I2 +
7H2O2.2Al + 3I2 2Al3+ + 6I-
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Standard EMF
For each of the half-cells in a voltaic cell, the standard reduction potential provides a measure of the driving force for reduction to occur.
The more positive the value of Eored,
the greater the driving force for reduction under standard conditions.
The more positive Eored value
identifies the cathode
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Determining Half-Reactions at ElectrodesA voltaic cell is based on the
following two standard half-reactions:
Cd2+ + 2e- Cd Sn2+ + 2e- SnBy using your chart, determine (a)
the half-reaction that occurs at the cathode and the anode, and (b) the standard cell potential
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Use Eored values to understand
aqueous reaction chemistryThe more positive the Eo
red value for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced and, therefore, to oxidize the other species. Better oxidizing agent
Strengths of Oxidizing and Reducing Agents
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The half-reaction with the smallest reduction potential is most easily reversed as an oxidation. The more negative the Eo
red, the stronger the ability to act as the reducing agent
Strengths of Oxidizing and Reducing Agents
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Strengths of Oxidizing and Reducing Agents
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Determining the Relative Strengths of Oxidizing AgentsUsing standard reduction potentials:Rank the following ions in order of
increasing strength as oxidizing agents: NO3
-, Ag+, Cr2O72-
Rank the following species from the strongest to the weakest reducing agent: I-, Fe, Al
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Free Energy and Redox Reactions
Determining the spontaneity of redox reactions
Eo = Eored (reduction process) –
Eored (oxidation
process)
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Spontaneous or Not?
Use standard reduction potentials to determine whether the following reactions are spontaneous under standard conditions.
1.Cu + 2H+ Cu2+ + H2
2.Cl2 + 2I- 2Cl- + I2
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Spontaneous or Not?
Use standard reduction potentials to understand the activity series of metals
Activity series of metals: strongest reducing agent at the top
Calculate standard emf forNi + 2Ag+ Ni2+ + 2Ag
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EMF and ΔG Relationship between G and EMF:
ΔG = -nFEWhere n = number of electrons transferred n
the reaction, G = Gibbs free energy, E = EMF, and F = Faraday’s constant
Faraday’s constant is the quantity of electrical charge on one mole of electrons (a faraday)
1 F = 96,485 C/mol = 96,485 J/V-mol
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Determining ΔGº and K
(a) Use standard reduction potentials to calculate ΔGº and K at 298K for the reaction:
4Ag + O2 + 4H+ 4Ag+ + 2H2O
(b) Suppose the reaction in part (a) was written: 2Ag + ½ O2 + 2H+ 2Ag+ + H2O What are values of Eº, ΔGº, and K when the reaction is written this way?
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For the reaction 3 Ni2+ + 2Cr(OH)3 + 10OH-
3Ni + 2CrO42-
+ 8H2O
(a) What is the value of n? (b) Given that ΔGº equals +87 kJ/mol, calculate K at a temperature of 298K
Determining n and K
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Cell EMF Under Nonstandard Conditions As a voltaic cell is discharged , the
reactants of the reaction are consumed and the products are generated The concentrations of these substances
changes EMF drops until E = 0, and the
concentration of reactants and products are at equilibrium
How does cell emf depend on the concentration of reactants and products?
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The Nernst Equation
Dependence of cell emf on concentration
Nernst Equation:
At 298K with units of volts, the equation simplifies to:
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The Nernst Equation
The Nernst equation helps us understand why the emf of a voltaic cell drops as the cell discharges
Increasing the concentration of reactants or decreasing the concentration of products increases the driving force (higher emf)
Decreasing the concentration of reactants or increasing the concentration of the products decreases the driving force (lower emf)
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Voltaic Cell EMF Under Nonstandard ConditionsCalculate the emf at 298K generated
by:
Cr2O72-(aq)+ 14H+(aq) + 6I-(aq)
2Cr3+(aq) + 3I2(s) + 7H2O(l)
When [Cr2O72-] = 2.0M, [H+] = 1.oM,
[I-] = 1.0M, [Cr3+] = 1.0x10-5M
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Calculate the emf at 298K generated by:
2Al (s)+ 3I2(s) 2Al3+ (aq) + 6I- (aq)
When [Al3+ ]= 4.0x10-3M and [I- ]0.010M
Voltaic Cell EMF Under Nonstandard Conditions
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Calculating Concentrations in a Voltaic cell If the voltage of the following cell is
+0.45V at 298K when [Zn2+] = 1.0M and PH2= 1atm, what is the concentration of H+?
Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)
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Batteries and Fuel Cells
A battery is a portable, self-contained electrochemical power source that consists of one or more voltaic cells.
When cells are connected in series, the battery produces a voltage that is the sum of the emfs of the individual cells. Multiple cells in series Multiple batteries in series
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Batteries and Fuel Cells
The substances that are oxidized at the anode and reduced at the cathode determine the emf of the battery. The usable life of the battery depends on
the quantities of these substances. Need a porous barrier between anode
and cathode compartments Primary and Secondary batteries
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Batteries and Fuel Cells
Lead-Acid Battery (12-v car battery, 6 voltaic cells in series that each produce 2V)
Alkaline Battery (most common primary battery)
Nickel-Cadmium, Nickel-Metal-Hydride, and Lithium-Ion Batteries (secondary batteries)
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Lead-Acid Batteries
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Corrosion
Corrosion reactions are spontaneous redox reactions in which a metal is attacked by some substance in its environment and converted to an unwanted compound. Oxidation is a thermodynamically
favored process in air at room temperature
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CorrosionPrevent corrosion by forming a
protective oxide layer that is impermeable to O2 and H2O
Examples: Al3+ forms protective Al2O3 layer Mg
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Corrosion of IronRusting of iron requires both oxygen
and water pH of solution, presence of salts, contact
with metals more difficult to oxidize than iron, and stress on the iron can accelerate rusting
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Corrosion of Iron
Cathodic protection: protecting a metal from corrosion by making it the cathode in an electrochemical cell.
The metal that is oxidized while protecting the cathode is called the sacrificial anode.
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Corrosion of Iron
Cathodic protection