Chapter 18 Oxidation-Reduction Reactions & Electrochemistry.
Electrochemistry Chapter 17. Electrochemistry The branch of chemistry that links chemical reactions...
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Transcript of Electrochemistry Chapter 17. Electrochemistry The branch of chemistry that links chemical reactions...
![Page 1: Electrochemistry Chapter 17. Electrochemistry The branch of chemistry that links chemical reactions to the production or consumption of electrical energy.](https://reader033.fdocuments.net/reader033/viewer/2022052509/56649cba5503460f94982b5f/html5/thumbnails/1.jpg)
ElectrochemistryChapter 17
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Electrochemistry
• The branch of chemistry that links chemical reactions to the production or consumption of electrical energy.
• In chemistry, electrical energy is stored in electrons. – In other words, electrochemistry is based upon
the principles oxidation-reduction reactions.
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Redox Revisited
• Earlier this year, we examined Redox reactions and how to write acid-base half reactions
• To review principles of redox reactions:– A redox reaction is the sum of two half reactions, the
reduction and oxidation reactions. – Reduction and oxidation reactions happen
simultaneously, so the number of electrons gained and lost must match exactly
– Oxidation=loss of electrons– Reduction=gain of electrons.
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Redox revisited
• Lets look at a common reaction of Zinc metal immersed in a copper sulfate solution.
The entire reaction is represented as a single replacement reaction where the blue CuSO4 solution becomes clear as the Zinc replaces the copper:
Zn + CuSO4 → Cu + ZnSO4
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Redox Revisited
Zn + CuSO4 → Cu + ZnSO4
The Redox half reactions are then represented as:
Zn(s) → Zn+2(aq) + 2e-
Cu+2(aq) + 2e- → Cu(s)
Zn(s) + Cu+2(aq) + 2e- → Cu(s) + Zn+2
(aq) + 2e-
So:Zn(s) + Cu+2
(aq) → Cu(s) + Zn+2(aq)
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Electrochemical (Galvanic) cells• An apparatus that
converts chemical energy into electrical work, or vice versa– Contains two
compartments– Bridge that allows flow
of energy (electrons)• “salt bridge” • Usually a piece of tubing
filled with an electrolyte
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Electrochemical (Galvanic) cells
• Anode-compartment in which the oxidation half reaction takes place
• Cathode-compartment which the reduction half reaction takes place. – We represent the reactions that take place using
cell diagrams• Cell diagrams are symbols that show how the
components of an electrochemical cell are connected.
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Electrochemical (Galvanic) cellsSalt bridge replaced with porous disk to allow ion flow and minimum mixing of solutions. Oxidizing agent “pulls” electrons through wire from reducing agent. Let there be light.
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Homework
• Pg 879-880 # 13-19.
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Cell Potential
• The “pull” on the electrons is called the cell potential (E °cell), or “electromotive force” (emf), is measured in volts.– Volt: 1 joule of work per coulomb of charge
transferred. 1 J/C– coulomb: defined as the charge transported by a
constant current of one ampere in one second:
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Standard Reduction Potentials
• E °--standard reduction potentials in Volts. – E °cell= E °cathode- E °anode
• Pay close attention to sign of E for certain reactions• If the reaction is in reverse, change the sign for the
reduction potential. • See Table 17.1 on page 843 in your book. • All reduction potentials are given with all solutes at
1M and all gases at 1 atm pressure. • E °cell is always positive
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Standard Reduction Potentials
Zn + CuSO4 → Cu + ZnSO4
The Redox half reactions are then represented as:Zn(s) → Zn+2
(aq) + 2e- -E °anode=.76
Cu+2(aq) + 2e- → Cu(s) E °cathode=.34
Zn(s) + Cu+2(aq) → Cu(s) + Zn+2
(aq) E °cell = 1.1V• This cell produces 1.1 volts of electrical energy
(work).
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Standard Reduction Potentials
• Another example, consider the galvanic cell based on the reaction:
Al+3(aq) + Mg(s) → Al(s) + Mg2+
(aq)
Give the redox half reactions, make sure to balance the reactions (equal # of electrons), and calculate E °cell (E ° for half reactions on table 17.1)
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Standard Reduction Potentials
Al+3(aq) + Mg(s) → Al(s) + Mg2+
(aq)
The Redox half reactions are then represented as:3(Mg(s) → Mg+2
(aq) + 2e)- -E °anode =2.37V
2(Al+3(aq) + 3e- → Al(s)) E °cathode=-1.66V
2Al+3(aq) + 3Mg(s) → 2Al(s) + 3Mg+2
(aq) E °cell = .71V• This cell produces .71 volts of electrical energy
(work).
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Standard Reduction Potentials
• Sometimes there are multiple possibilities for redox potentials, in this case, pay close attention to what the equation they give states the cell is based on.
• Example: a galvanic cell is based on the reaction:– MnO4
-(aq) + H+
(aq) + ClO3-(aq) → ClO4
-(aq) + Mn+2
(aq) + H2O(l)
• There are multiple oxidation reactions for MnO4-(aq) , so you must
consult table 17.1 and match the reactants and products.
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Representing Cells with Line Notations
• Consider key components of this galvanic cell:
Zinc solid electrode
Zinc ions in solution
Spectator ions
Spectator ions
Copper ions in solution
Copper solid electrode
Salt bridge
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Steps to Representing Cells with Line Notations
• Rule #1…list everything• Separate the cathode/anode with the double line
notation (II) that represents the salt bridge• Separate substances in different states of matter in
the same compartment with a single line (I), separate substances in the same state in the same compartment with a comma.
• Dispense spectator ions (usually water)• When there is no solid electrode listed, assume
there is Platinum (Pt) in both.
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Representing Cells with Line Notations
• Consider key components of this galvanic cell:
Zinc solid electrode
Zinc ions in solution
Spectator ions
Spectator ions
Copper ions in solution
Copper solid electrode
Salt bridge
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homework
• Pg 880, #’s 25-35 odd
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Determining Spontaneity in Galvanic cells
• 1. Any reduction reaction is spontaneous when paired with the reverse of any reaction below it on Table 17.1.
• 2.If the Cell potential calculated is negative, the reaction is not spontaneous (yes, I know I told you that the cell potential is always positive…in a cell that works).
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Cell Potential, Electrical Work, And Free Energy (∆G)
• Potential Difference (V)= work (J)/charge©– When a cell produces a current (V), the cell
potential is positive, and the current can be used to do work.
– Since the work does not stay in the system (the cell), the sign for work is negative. So:
q=charge in coulombs transferred from anode to cathode
q
w E
qEw
qEw
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Cell Potential, Electrical Work, And Free Energy (∆G)
qEw Work measured in Joules
Coulombs based on Faraday Constant
Cell potential difference in V or J/C
The Faraday Constant: the charge on one mole of electrons is 96,485 Coulombs of charge . When 1.33 moles is transferred:
nFq = 1.33 mol e- x 96,485 C/mol e-
This equation calculates amount of work done. HOWEVER, since some work and energy is always lost to the surroundings as heat, there is a way to calculate Maximum work
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Cell Potential, Electrical Work, And Free Energy (∆G)
• All galvanic cells have a maximum potential that they never reach because of energy lost as heat. To calculate maximum work, use the maximum potential in your equation. And then the Maximum work equals ∆G.
maxmaxmax nFEqEGw
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Dependence of a Cell on Concentration
• Simply put…if the product concentration is raised above 1.0M, E °cell will be less than what is listed 17.1
• If the reactant concentration is above 1.0M, E °cell will be greater than listed in table 17.1
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Dependence of a Cell on Concentration
• The dependence of cell potential on concentration relies directly on the dependence of free energy on concentration. Remember that
• And since• And • Then:
)ln(QRTGG cellcell nFEqEG max
nFEG
)ln(QnF
RTEE “Nerst Equation”
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Nerst Equation
At 25°C:
E=E°-
n= moles of electrons
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homework
• Pg 881, #’s 39-59 odd.