CHEMISTRY - SILBERBERG 8E CH.21 - …lightcat-files.s3.amazonaws.com/packets/chemistry-4...CHEMISTRY...

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CHEMISTRY - SILBERBERG 8E

CH.21 - ELECTROCHEMISTRY

CONCEPT: OXIDATION-REDUCTION REACTIONS

Chemists use some important terminology to describe the movement of electrons.

• In ______________ reactions we have the movement of electrons from one reactant to another.

L

E

O

G

E

R

Agent Agent

Rules for Assigning an Oxidation Number (O.N.)

A. General Rules

1. For an atom in its elemental form (Na, O2, S8, etc.): O.N. = 0

2. For an ion the O.N. equals the charge: Na+ , Ca2+ , NO3 –

B. Specific Rules

1. Group 1A: O.N. = +1

2. Group 2A: O.N. = +2

3. For hydrogen: O.N. = +1 with nonmetals

O.N. = -1 with metals and boron

4. For Fluorine: O.N. = -1

5. For oxygen: O.N. = -1 in peroxides (X2O2 , X = Group 1(A) element)

O.N. = − 12

in superoxides (XO2 , X = Group 1(A) element)

O.N. = - 2 in all other compounds

6. Group 7A O.N. = -1 (except when connected to O)



CONCEPT: OXIDATION-REDUCTION REACTIONS (PRACTICE)

EXAMPLE: In the following reaction identify the oxidizing agent and the reducing agent:

a. 2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (g)

PRACTICE: What is the oxidation number of each underlined element?

a. P4 b. BO33-

c. AsO42- d. HSO4

–

PRACTICE: In the following reaction identify the oxidizing agent and the reducing agent:

a. Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O



CONCEPT: BASIC REDOX CONCEPTS OXIDATION-REDUCTION (REDOX) reactions deal with the transfer of electrons from one reactant to another.

Lose

Electrons

Oxidation

Gain

Electrons

Reduction

Reducing Agent (Reductant)

} }} Element

becomes

more positive

Oxidation

Number

Increases } Element

becomes

more negative

Oxidation

Number

Decreases

Oxidizing Agent (Oxidant)

Li (s) Li+ (aq) + e – Cl2 (g) + 2 e – 2 Cl – (aq)

Li (s) + Cl2 (g) Li+ (aq) + 2 Cl – aq)

Electrical Charge

The units for electrical charge are measured in ________________ (C).

} Charge of 1 electron Faraday Constant

}(1.602×10−19C) ⋅ (6.022×1023mol−1) = 9.647×104C

1mole e−charge mole

e –

q = n ⋅ F Faraday Constant

Electrical Current

The units for electrical current are in __________ (A).

Electrical Voltage

The relationship between work and voltage can be expressed as:

The relationship between Gibbs Free Energy and electric potential can be expressed as:

Ohm's Law

The units for resistance are in __________ (Ω).

Power

Power represents work done per unit of time. The units for power are in __________ (W).

w = E ⋅ qWork Voltage Charge

ΔG = − n ⋅ F ⋅ E GibbsFree Energy

mole e –

Faraday Constant

Voltage

I = ER

Voltage

ResistanceCurrent

P = E ⋅ IPower Voltage Current

Current

Charge

TimeI = q

t



CONCEPT: BALANCING REDOX REACTIONS

Now that you have refreshed your memory in terms of redox reactions you now have to master balancing them. Generally,

you will need to balance a redox reaction in an acidic or basic solution.

Balancing A Redox Reaction in Acidic Reactions: STEP 1: Write the equation into 2 half-reactions.

STEP 2: Balance elements that are not oxygen or hydrogen.

STEP 3: Balance Oxygens by adding H2O.

STEP 4: Balance Hydrogens by adding H+.

STEP 5: Balance overall charge by adding electrons (e –) to the more _________________ side.

STEP 6: Combine the half-reactions and cross out reaction intermediates.

Balancing A Redox Reaction in Basic Reactions: Follow Steps 1-6 from above.

STEP 7: Balance H+ by adding OH – ions to both sides of the chemical reaction.



CONCEPT: BALANCING REDOX REACTIONS

Now that you have refreshed your memory in terms of redox reactions you now have to master balancing them. Generally, you will need to balance a redox reaction in an acidic or basic solution.

Balancing A Redox Reaction in Acidic Reactions: STEP 1: Write the equation into 2 half-reactions.

STEP 2: Balance elements that are not oxygen or hydrogen.

STEP 3: Balance Oxygens by adding H2O.

STEP 4: Balance Hydrogens by adding H+.

STEP 5: Balance overall charge by adding electrons (e –) to the more _________________ side and make sure both half

reactions have the same number of electrons.

STEP 6: Combine the half-reactions and cross out reaction intermediates.

Balancing A Redox Reaction in Basic Reactions: Follow Steps 1-6 from above.

STEP 7: Balance H+ by adding OH – ions to both sides of the chemical reaction.



PRACTICE: BALANCING ACIDIC REDOX REACTIONS

EXAMPLE: Balance the following reaction in an acidic solution.

NO2– + MnO4

– NO3 – + Mn2+

PRACTICE: Balance the following reaction in an acidic solution.

Cl2(g) + S2O32-(aq) Cl-(aq) + SO42-(aq)



PRACTICE: BALANCING BASIC REDOX REACTIONS

EXAMPLE: Balance the following reaction in a basic solution.

Mo3O93- (aq) + Br – (aq) Mo (aq)

+ BrO4

2- (aq)

PRACTICE: Balance the following reaction in a basic solution.

XeO2 (aq) H2Xe (aq) + XeO4



CONCEPT: NERNST EQUATION The Nernst Equation reveals the quantitative connection between the concentrations of compounds and cell potential.

E = Eo − RTnFln AB

b

AAa

Eo = Standard Cell Potential

R = Gas Constant = ______________________

n = Number of electrons transferred

F = Faraday's constant = __________________

A = Activity

______ represents the cell potential under non-standard conditions, while ______ represents it under standard conditions.

At 25 oC, RTF

=8.314 J

K ⋅mol⎞

⎠⎟ 298.15 K)(

⎛

⎝⎜

9.649×104 Cmol

⎞

⎠⎟

⎛

⎝⎜

= 0.0257 JC= 0.0257 V

The Nernst Equation then becomes,

ECell = ECello −

0.0257 Vn

lnQ

By multiplying ln by 2.303 we can obtain the log function.

ECell = ECello −

0.05916 Vn

logQ

The cell potential calculated from Nernst equation is the maximum potential at the instant the cell circuit is connected. As the cell discharges and current flows, the electrolyte concentrations will change, Q increases and Ecell deceases.

• Over time the reaction will reach equilibrium at then Q = K and cell potential will equal zero.

ECell = ECello −

RTnF

⎞

⎠⎟

⎛

⎝⎜ lnK = 0 ΔG = ΔGo −RTlnK = 0



CONCEPT: ELECTRICAL CURRENT

EXAMPLE 1: Gold can be plated out of a solution containing Au3+ based on the following half reaction:

Au3+ (aq) + 3 e - Au (s)

a) What mass of gold is plated by a 41 minute flow of 6.8 A current?

EXAMPLE 2: A solution of Mn+5 is used to plate out Mn in an electrochemical cell. If a total of 1.13 g of Mn is plated out in a

total time of 1600 seconds, what was the electrical current used? (MW of Mn is 54.94 g/mol)



PRACTICE: ELECTRIC CURRENT (CALCULATIONS 1)

EXAMPLE: If steady current of 15 amperes is provided by a stable voltage of 12 Volts for 600 seconds, answer each of the following questions.

a) Calculate the total charge that passes through the circuit in this time.

b) Calculate the total number of moles of electrons that pass through the circuit in this time.

c) Calculate the total amount of energy that passes through the circuit in this time.

d) Calculate the power that the battery provides during this process.



CONCEPT: GALVANIC VS. ELECTROLYTIC CELLS

With REDOX reactions we will now deal with a new variable: a reaction’s cell potential, which uses the variable _________ .

• The greater this variable then the more likely ______________________ will occur.

• The smaller this variable then the more likely ______________________ will occur.

Galvanic/Voltaic Cell: ____________ or ____________ electricity so it’s a ________________ .

Galvanic/Voltaic Cell

__________________ ( – ) __________________ ( + )

Cd2+ (aq) + 2 e – Cd (s) Eo = - 0.40 V

Ni2+ (aq) + 2 e – Ni (s) Eo = - 0.25 V

For all redox reactions under standard conditions it is possible to determine if a reaction is spontaneous or not.

∆SUniverse ∆Go K Eo Reaction Under Standard State Conditions

Cell Type

Positive Negative >1 Positive

Negative Positive <1 Negative

0 0 1 0

Anode ___________

Producing ñ Voltage

[Anode] ______

[Cathode] ______

Ionization Energy __________

Cathode __________

Electron Affinity __________



CONCEPT: ELECTROLYTIC CELLS In terms of spontaneity the following correlations between the following variables can be made:

ΔGo

K

Eo

ΔSo

Q vs. K

Reaction Classification

Cell Type

< 0

> 1

> 0

> 0

Q < K

> 0

< 1

< 0

< 0

Q > K

= 0

= 1

= 0

= 0

Q = K

Electrolytic Cell: A non-spontaneous electrochemical cell that _______________ electricity and so requires a battery.

Ionization Energy __________

Electron Affinity __________

Electrolytic Cell

__________________ ( – ) __________________ ( + )



PRACTICE: GALVANIC VS. ELECTROLYTIC CELLS (CALCULATIONS)

EXAMPLE: A certain electrochemical cell involves a five electron change and has a value of Keq = 3.0 x 1016 at 298 K. The

value of ∆Ho for the reaction is -68.3 kJ/mol. Calculate the values of ∆Go, ∆Eo, for a standard electrochemical cell

constructed based on this reaction and also ∆So for the reaction.

PRACTICE: Given the following standard reduction potentials,

Hg22+(aq) + 2 e– 2 Hg(l) E° = +0.789 V

Hg2Cl2(s) + 2 e– 2 Hg(l) + 2 Cl-(aq) E° = +0.271 V

determine Ksp for Hg2Cl2(s) at 25 °C.



PRACTICE: CELL NOTATION (CALCULATIONS 1)

EXAMPLE: The cell notation for a redox reaction is given as the following at (T = 298 K):

Zn (s) Ι Zn2+ (aq, 0.37 M) ΙΙ Ni2+ (aq, 0.059 M) Ι Ni (s)

a) Write the balanced half-reactions occurring at the anode and the cathode.

b) Write out the complete balanced redox reaction.

c) Determine the Eo cell.

d) Calculate the maximum electrical work that can be produced by this cell.

e) Calculate the reactant quotient, Q, for this cell and the cell potential under non-standard conditions.



PRACTICE: GALVANIC VS. ELECTROLYTIC CELLS (CALCULATIONS 2)

EXAMPLE: Answer each of the following questions based on the following half reactions:

HALF REACTIONS Eo (V)

Cl2 (g) + 2 e – 2 Cl – (aq) + 1.36

l2 (g) + 2 e – 2 l – (aq) + 0.535

Pb2+ (aq) + 2 e – Pb (s) - 0.126

V2+ (aq) + 2 e – V (s) - 1.18

a) Which is the strongest oxidizing agent?

b) Which is the strongest reducing agent?

c) Will I – (aq) reduce Cl2 (g) to Cl – (g)?



Electric current, or flow of electrons, is measured in Amperes (A). One Ampere is the delivery of one coulomb (C) of charge per second. What mass of Zinc (in g) is oxidized (to Zn2+) by a dry cell battery that supplies 125 mA of current for two hours (recall that Faraday’s constant is the charge in coulombs on a mole of electrons)?

a. 0.03 b. 0.1 c. 0.3 d. 1.0 e. 3.0



Consider the combustion of formaldehyde CH2O and the data below for the following five questions.

I. CH2O(g) + O2(g) → CO2(g) + H2O(l)

II. 4H+ + 4e- + O2(g) → 2H2O(l) ΔE° = 1.23V

Which compound is reduced in reaction I?

a. CH2O b. H2O c. O2 d. CO2 e. None of these

What is the change in oxidation number of the carbon in reaction I?

a. -4 b. -1 c. 0 d. 1 e. 4

What is the ΔG° for the combustion of formaldehyde (kJ/mol)?

a. 224 b. -521 c. 96 d. -150 e. more date is required

How many electrons are required to balance the half reaction CH2O → CO2 in acidic solution?

a. 1 b. 2 c. 3 d. 4 e. 5

What is the standard half cell potential (in volts) for the CH2O oxidation?

a. 1.7 b. 0.76 c. 0.12 d. 3.33 e. 1.69



The purpose of a galvanic cell is to:

a. Transduce chemical energy to electrical energy.

b. Purify solids.

c. Allow for oxidation without reduction.

d. To consume electricity.



During the process for electrolysis of water a current is passed through water and produces hydrogen gas

and oxygen gas. Which of the following statements is true?

a. O2 gas is produced at the anode.

b. H2 gas is produced at the cathode.

c. In the reaction, H2 moles are twice the O2 moles.

d. All of the following are correct.



Which statement is false?

a. Reduction occurs at the cathode.

b. A reducing agent will lose electrons.

c. Cations migrate to the cathode in both electrolytic and electrochemical cells.

d. Li (s) is the strongest oxidizing agent; F2 is the strongest reducing agent.



Which of the following reactions may occur at the anode?

a. Ga3+ (aq) + 3 e– Ga (s)

b. Cu2+ (aq) + 2 e– Cu (s)

c. 2 Cl – (aq) Cl2 (g) + 3 e–

d. Co (s) + e– Co+ (aq)



Define a salt bridge.

A) A pathway, composed of salt water, that ions pass through.

B) A pathway in which no ions flow.

C) A pathway between the cathode and anode in which ions are reduced.

D) A pathway between the cathode and anode in which ions are oxidized.

E) A pathway by which counterions can flow between the half-cells without the solutions in the half-cell

totally mixing.



What statement is NOT true about standard electrode potentials?

A) E°cell is positive for spontaneous reactions.

B) Electrons will flow from more negative electrode to more positive electrode.

C) The electrode potential of the standard hydrogen electrode is exactly zero.

D) E°cell is the difference in voltage between the anode and the cathode.

E) The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative

to the standard hydrogen electrode and therefore has a positive E°.



Use the standard reduction potentials below to determine which element or ion is the best reducing agent.

Pd2+ (aq) + 2 e – Pd (s) E° = + 0.90 V

2 H+ (aq) + 2 e – H2 (g) E° = 0.00 V

Mn2+ (aq) + 2 e – Mn (s) E° = – 1.18 V

a) Pd (s) b) H+ (aq) c) Mn2+ (aq) d) H2 (g)



Consider an electrochemical cell where the following reaction takes place:

Na2O (aq) + Ba (s) 2 Na (s) + BaO (aq)

What is the cell notation for this cell?

What is the ratio of oxidizing agent to reducing agent?



CONCEPT: BATTERIES

Batteries can be thought of as portable galvanic or voltaic cells that generate electricity.

Dry-Cell Batteries

These types of batteries are called dry-cell because they lack large amounts of liquid water. In a typical dry-cell battery the zinc sleeve acts like the ________________:

Zn (s) Zn2+ (aq) + 2 e – Eo = - 0.76 V The ______________ portion has a carbon tube immersed in a damp paste of MnO2 that also possesses NH4Cl.

2 MnO2 (s) + 2 NH4+ (aq) + 2 e – Mn2O3 (s) + 2 NH3 (g) + H2O (l) Eo = 0.74 V

The counterpart to the more acidic dry-cell battery is the more commonly used _________________ battery.

Zn (s) + 2 OH – (aq) Zn(OH)2 (s) + 2 e –

2 MnO2 (s) + 2 H2O (l) + 2 e – 2 MnO(OH) (s) + 2 OH – (aq)



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