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18. Electrochemistry

electrochemistry: chemistry of flowing electrons

oxidation-reduction reaction

oxidation involves loss, reduction involves gain

example:

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Galvanic cell

galvanic cell: converts chemical energy into electrical energy

ions in liquid must match metal

metals connected with wire

voltmeter: measures voltage between two points in electric circuit

volt (V): number of electrons flowing past a point in an amount of time

V = 0 at equilibrium

salt bridge: piece of paper soaked in ionic compound

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anode: side of galvanic cell where oxidation occurs

cathode: side of galvanic cell where reduction occurs

Cell potential

cell potential (E): the pull of electrons through a wire (volts)

total cell potential must be positive for the cell to work

flip equation: E changes sign

do not multiply E if you multiply coefficients

half-cell: just the anode or just the cathode

hydrogen standard: E = 0

example:

reaction and total cell potential?

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flip Mg reaction (make positive)

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need to add the values of E

total cell potential: 0.71 V

balance electrons by multiplying equations by constant

voltage does not multiply

add equations together

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Mg is anode, Al is cathode

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ΔG: Gibb’s free energy (J)

n: moles of electrons (mol)

F: Faraday’s constant (96485 C/mol e⁻)

coulomb = C = J/V

faraday (unit): moles of electrons

E: cell potential (V)

changing the amount of metal does not do anything, only changing concentrations

Nonstandard cell potential

standard cell potential (E°): cell potential under standard conditions (1 M)

nonstandard cell potential (E): cell potential not under standard conditions

Nernst equation:

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E: nonstandard cell potential

E°: standard cell potential

R: gas constant (J/K)

T: temperature (K)

n: electrons (mol)

F: Faraday’s constant (96485 C/mol e⁻)

Q: reaction quotient

adding to the left increases E, adding to the right decreases E

example:

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if [Al³⁺] = 2.0 M, [Mn²⁺] = 1.0 M: E would be less

if [Al³⁺] = 1.0 M, [Mn²⁺] = 3.0 M: E would be greater

Electrolysis

electrolysis: adding electricity to reverse the reaction in a galvanic cell

e.g. charging devices, plating metals

ampere (amp): speed of electron current (coulombs per second)

use Faraday’s constant to convert to moles of electrons

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I: flow of electrons (A)

q: charge (C)

t: time (s)

example:

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, add 10.0 amps of electricity for 30 min; how much Cu(s) is made (g)?

amperes → coulombs:

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coulombs → moles of electrons:

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moles of electrons → moles of copper:

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moles → grams

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Galvanic cell

Cell potential

Nonstandard cell potential

Electrolysis

Gallery

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