ElectrochemistryPart 1

Ch. 20 in Text

(Omit Sections 20.7 and 20.8)

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I) Overview

electrochemistry deals with: electricity and chemistry electron transfers oxidation: loss of electrons reduction: gain of

electrons changes in oxidation

number single replacement,

synthesis, and decomposition

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A) Oxidation

the loss of negatively charged electrons results in an INCREASE in oxidation number

often involves metals because it CAUSES

reduction, the species oxidized is known as the reducing agent

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B) Reduction

the gain of negatively-charged electrons results in a DECREASE in oxidation number

often involves nonmetals

because it CAUSES oxidation, the species reduced is known as the oxidizing agent forums.joeuser.com

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C) An Example

the following reaction occurs in a nickel-cadmium battery:

Cd(s) + NiO2(s) + 2H2O(l)

→ Cd(OH)2(s) + Ni(OH)2(s)

failblog.orgHW: 20.6

D) Half-Reactions

show a redox reaction as 2 separate processes

atoms and charge must be conserved

oxidation represents electrons as products

reduction represents electrons as reactants

the number of electrons lost in oxidation must equal that gained in reduction sciencenotes.wordpress.com

Cd(s) + NiO2(s) + 2H2O(l) → Cd(OH)2(s) + Ni(OH)2(s)

E) Balancing Equations by the Half-Reaction Method

may be in acidic or basic solution

see textbook

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II) Voltaic Cells (Galvanic Cells) a spontaneous redox

reaction performs electrical work

a half cell consists of a metal strip placed in a solution of its ion; when connected they form a voltaic cell

the metal strips are connected with a wire for the movement of electrons

the solutions are connected with a salt bridge for the movement of ions (complete circuit)

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electrons move from the more active metal (anode) where oxidation occurs to the less active metal (cathode) where reduction occurs

ions move through the salt bridge to maintain electrical neutrality

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the anode gets smaller over time since the solid metal is breaking down into aqueous ions

the cathode gets

bigger over time since the aqueous ions are building up as a solid metal

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III) Cell EMFA) Potential Energy the flow of electrons from

anode to cathode is spontaneous due a difference in potential energy between the two metals

the potential difference drives the reaction

the potential difference in energy per unit charge is measured in volts

1 V = 1 J/C where J = Joules and C = Coulombs todayscampus.com

B) EMF

the potential difference between the two electrodes is also known as the electromotive force or emf

standard emf or standard cell potential is measured at 1 atm and 25 ºC for 1 M solutions

denoted Eºcell

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C) SHE although the standard cell potential

could be measured for ALL different metal combinations, it is easier to measure all standard cell potential values in reference to a COMMON standard cell

the standard hydrogen electrode (SHE) is designated as this standard

2H+(aq, 1 M) + 2 e- → H2(g, 1 atm) is given a cell potential of 0 V (arbitrarily)

all other half-cells are measured against this standard mikeblaber.org

D) Standard Reduction Potentials

by convention, the potential of an electrode is chosen to be the potential for reduction to occur (in relation to the SHE)

thus, tabulated standard emf potential values = Eºred

Eºcell = Eºred(cathode) - Eºred(anode)

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the more positive the Eºred, the greater the driving force for reduction

the more negative the Eºred, the greater the driving force for oxidation

FYI, changing the coefficient in a half-rxn to balance the net equation does NOT change the value of Eºred!!!

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Ex) Calculate the standard emf for a cell involving the following rxn:

2Al(s) + 3I2(s) → 2Al+3(aq) + 6I-(aq)

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HW: 20.24, 20.26 (a) & (b), 20.32

E) Oxidizing/Reducing Agents

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hsc.csu.edu.auHW: 20.40

ElectrochemistryPart 2

Ch. 20 in Text

(Omit Sections 20.7 and 20.8)

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IV) SpontaneityA) Basics

a “+” value for E°cell means that the rxn is spontaneous

a “-” value for E°cell means that the rxn is nonspontaneous

Ex) Determine the spontaneity of the following:

Cu(s) + 2H+(aq) → Cu+2(aq) + H2(g)

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B) Free Energy Change

ΔG = -nFE where: ΔG is the free energy change in

J/mol n is the # of electrons

transferred in the rxn (no units) F is Faraday’s constant (1 F =

96,500 C/mol = 96,500 J/V•mol) E is the emf in V

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Ex) If E°cell = 0.43 V for the following rxn, what is the standard free energy change?

4Ag(s) + O2(g) + 4H+(aq) → 4Ag+(aq) + 2H2O(l)

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HW: 20.46 (a), 20.50

V) Concentration and Cell EMFA) The Nernst Equation

E = E° - RT ln QOR nF

E = E° - 0.0592 V log Q n (at 25 °C)

we can use the above equations to determine the voltage under nonstandard conditions

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Ex) If E°cell = 1.10 V for the following reaction, what is Ecell if [Cu+2] = 5.0 M and [Zn+2] = 0.050 M at 298 K?

Zn(s) + Cu+2(aq) → Zn+2(aq) + Cu(s)

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HW: 20.54

B) Concentration Cells

emf generated solely by a difference in concentration is called a concentration cell

although the E°cell = 0, this occurs under nonstandard conditions

driven by equalization of concentrations

the more dilute solution is the anode and the more concentrated solution is the cathode

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Link

Ex) Write half reactions for the previously depicted cell and calculate the voltage at 25 °C.

debate.orgHW: 20.58

C) Cells and Equilibrium

cells eventually “die” out as reactants are converted to products, Q increases, and thus, E decreases

when E = 0, ΔG = 0 and we are equilibrium

log K = nE°/0.0592 at 298 K

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Ex) O2(g) + 4H+(aq) + 4Fe+2(aq) → 4Fe+3(aq) + 2H2O(l)

If E° = 0.46 V at 298 K, what is the value of the equilibrium constant?

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VI) Electrolytic CellsA) Set Up

electrical energy is used to cause a nonspontaneous rxn to occur

a power source (like a battery) must be added to force the rxn to occur

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B) Quantitative Aspects

I = q/t where: I is the current in amperes

(amps), A q is the charge in Coulombs t is time in seconds knowing the current can allow

us to convert to grams of solid plated in an electrolytic cell

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Ex) What mass of Al is produced in 1.00 hr by the electrolysis of molten AlCl3 if the electrical current is 10.0 A?

failblog.orgHW: 20.82 (a)

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