Chapter 11: Phenomena Phenomena: Two electrochemical cells were - - PowerPoint PPT Presentation

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Chapter 11: Phenomena Phenomena: Two electrochemical cells were - - PowerPoint PPT Presentation

Feb 29, 2024 33 likes •309 views

Chapter 11: Phenomena Phenomena: Two electrochemical cells were constructed. Experiments were done by changing mass, concentration, and/or temperature. Examine the data to determine patterns in how these variables affect the voltage of the

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Chapt pter er 11: Electr troc

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Chapter 11: Phenomena

Phenomena: Two electrochemical cells were constructed. Experiments were done by changing mass, concentration, and/or temperature. Examine the data to determine patterns in how these variables affect the voltage of the cell.

Half Cell 1: Zn(s) & Zn2+(aq) Half Cell 2: Cu(s) & Cu2+(aq) Cell Reaction: Zn(s)+Cu2+(aq) Zn2+(aq)+Cu(s) Exp. Mass Zn(s) [Zn2+] Volume Zn2+ Mass Cu(s) [Cu2+] Volume Cu2+ Temp Voltage 1 1.0 g 1.0 M 0.50 L 1.0 g 1.0 M 0.50 L 298 K 1.10 V 2 1.0 g 1.0 M 0.50 L 1.0 g 5.0 M 0.50 L 298 K 1.12 V 3 2.0 g 1.0 M 0.50 L 1.0 g 1.0 M 0.50 L 298 K 1.10 V 4 2.0 g 1.0 M 0.50 L 3.4 g 1.0 M 0.50 L 298 K 1.10 V 5 5.0 g 1.0 M 0.50 L 2.5 g 1.0 M 0.50 L 415 K 1.07 V 6 1.0 g 0.2 M 0.50 L 1.0 g 1.0 M 0.50 L 298 K 1.12 V 7 3.4 g 5.0 M 0.50 L 1.0 g 1.0 M 0.50 L 415 K 1.05 V Half Cell 1: Ca(s) & Ca2+(aq) Half Cell 2: Ag(s) & Ag+(aq) Cell Reaction: Ca(s)+2Ag+(aq) Ca2+(aq)+2Ag(s) Exp. Mass Ca(s) [Ca2+] Volume Ca2+ Mass Ag(s) [Ag+] Volume Ag+ Temp Voltage 1 1.0 g 1.0 M 0.50 L 1.0 g 1.0 M 0.50 L 298 K 3.67 V 2 1.0 g 1.0 M 0.50 L 1.0 g 5.0 M 0.50 L 298 K 3.71 V 3 1.0 g 1.0 M 1.00 L 1.0 g 1.0 M 0.50 L 298 K 3.67 V 4 1.0 g 1.0 M 1.00 L 1.0 g 1.0 M 1.70 L 298 K 3.67 V 5 1.0 g 1.0 M 2.50 L 1.0 g 1.0 M 2.50 L 415 K 3.57 V 6 1.0 g 0.2 M 0.50 L 1.0 g 1.0 M 0.50 L 298 K 3.69 V 7 1.0 g 5.0 M 1.70 L 1.0 g 1.0 M 0.50 L 415 K 3.55 V

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Chapter 11

Electrochemistry

  • Redox Reaction Review
  • Galvanic Cells
  • Thermo of

Electrochemistry

  • Nernst Equation
  • Batteries / Fuel Cells
  • Electrolytic Cells

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Big Idea: Electron transfer in a chemical reaction is both material and concentration

  • specific. If the process

is spontaneous the transfer of electrons can be used to produce a current and drive electrical devices.

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Redox Reaction Review

Assigning Oxidation Numbers

 The oxidation number (ON) of an element uncombined with

another element is zero: Na(s) ON = 0, and H2(g) ON = 0.

 For monatomic ions, the charge is the ON: Na+ ON = +1.  The ONs of elements in group 1 equal 1 (ex. Lithium ON = +1)

ONs of elements in group 2 equal 2 (ex. Magnesium ON = +2), when the atoms are in a compound.

 The ON of fluorine is always -1 in compounds.  The ON of the other elements in group 7 usually equal -1

when the atoms are in a compound.

 The ON of oxygen is usually -2 in compounds. Exceptions are

fluorine compounds and peroxide (a compound that contains an O-O single bond).

 Hydrogen's ON is +1 when combined with non metals and -1

when combined with metals.

 The sum of the ON’s of all the atoms in a species is equal to its

total charge.

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Redox Reaction Review

Assigning Oxidation Numbers

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NaCl

Element Oxidation Number Reason Na Cl

Fe2(SO4)3

Element Oxidation Number Reason SO4 O S Fe

Br2

Element Oxidation Number Reason Br

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Redox Reaction Review

Determining Which Element is Oxidized and Which is Reduced

 Step 1: Assign oxidation numbers.  Step 2: Use oxidation numbers and ‘OIL RIG’ to

identify which element is oxidized and which is reduced.

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Not

  • te: If the question asks which “substance is oxidized;” instead of giving the

element that is oxidized give the entire compound.

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Redox Reaction Review

Oxidizing Agent: A species that removes

electrons from a species being oxidized in a redox reaction.

Reducing Agent: The species that supplies

electrons to a substance being reduced in a redox reaction.

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Not

  • te: Oxidizing agent is the species being reduced.

Note: Reducing agent is the species being oxidized.

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Redox Reaction Review

Balancing Redox Reactions in Acidic Conditions

 Step 1: Write unbalanced half reactions.  Step 2: Balance half reactions except for O and

H.

 Step 3: Balance O by using H2O.  Step 4: Balance H by using H+.  Step 5: Balance electrons in each half reaction.  Step 6: Multiply half reactions by an integer so

that number of electrons match, then add the reactions together.

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Redox Reaction Review

Balancing Redox Reactions in Basic Conditions

 Step 1: Balance the reaction as if it were in

acidic conditions.

 Step 2: Determine the number of H+ in the

balanced equation.

 Step 3: Add the same number OH- as there are

H+ to BOTH sided of the equation.

 Step 4: The H+ and OH- on one side of the

reaction will combine and form H2O.

 Step 5: Simplify your reaction (combine waters) if

necessary.

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Galvanic Cell

 Electrochemical Cell: Device in which an electric current

is produced by either a spontaneous reaction or is used to bring about a non spontaneous reaction.

 Galvanic Cell: An electrochemical cell in which a

spontaneous chemical reaction is used to generate an electrical current.

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Electrode: Metal contacts. Electrolyte: Ionically conducting medium. Salt Bridge: Allows for the flow

  • f ions but prohibits reactions

from taking place. Anode: Is where oxidation

  • ccurs.

Cathode: Is where reduction

  • ccurs.
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Galvanic Cell

Short Hand Cell Notation

1)

The anode is written on the left and cathode on the right.

2)

Phase interfaces are separated with a ‘|’

3)

Same phase species are separated by a ‘,’

4)

For same phases, the species that is capable of being the

  • xidizing agent is written 1st followed by the species that is

capable of being the reducing agent.

 Example:  Fe3+(aq) + e-  Fe2+(aq) Fe3+ is the oxidizing agent  (anode) Fe3+(aq), Fe2+(aq)||  (cathode) ||Fe3+(aq), Fe2+(aq) 5)

The salt bridge is represented by ||

6)

Never include water

7)

Anode order of phase: solid | gas | liquid | aqueous

8)

Cathode order of phase: aqueous | liquid | gas | solid

9)

An inert electrode always goes on the outside (usually made up of Pt(s) or C(gr))

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Student Question

Galvanic Cell

Write the balanced reaction for the given cell

Pt(s)|H2(g)|H+(aq)||Co3+(aq),Co2+(aq)|Pt(s)

a)

H2(g)+Co3+(aq) 2H+(aq)+Co2+(aq)

b)

Pt(s)+H2(g)+2Co3+(aq)2H+(aq)+2Co2+(aq)+Pt(s)

c)

2H+(aq)+Co2+(aq)H2(g)+Co3+(aq)

d)

H+(aq)+Co2++e-(aq) H2(g)+Co3+(aq)

e)

None of the above

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Thermo of Electrochemistry

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Standard Reaction Potentials at 298 K Half -Reaction E°(V) F2 + 2e-  2F- 2.87 Ce4+ + e-  Ce3+ 1.70 MnO4

  • + 4H+ + 3e-  MnO2 + 2H2O

1.68 IO4

  • + 2H+ + 2e-  IO3
  • + H2O

1.60 MnO4

  • + 8H+ + 5e-  Mn2+ + 4H2O

1.51 Au3+ + 3e-  Au 1.50 Cl2 + 2e- 2Cl- 1.36 Cr2O7

2- + 14H+ + 6e-  2Cr3+ + 7H2O

1.33 O2 + 4H+ + 4e-  2H2O 1.23 MnO2 + 4H+ + 2e-  Mn2+ + 2H2O 1.21 IO3

  • + 6H+ + 5e-  ½I2 +3H2O

1.20 Br2 + 2e- 2Br- 1.09 VO2

+ + 2H+ + e-  VO2+ + H2O

1.00 AuCl4

  • + 3e-  Au + 4Cl-

0.99 NO3

  • + 4H+ + 3e- NO + 2H2O

0.96 ClO2 + e-  ClO2

  • 0.95

2Hg2+ + 2e-  Hg2

2+

0.91 Ag+ + e-  Ag 0.80 Hg2

2+ + 2e-  2Hg

0.80 Fe3+ + e-  Fe2+ 0.77 MnO4

  • + e-  MnO4

2-

0.56 I2 + 2e-  2I- 0.54 Cu+ + e-  Cu 0.52 Cu2+ + 2e-  Cu 0.34 Standard Reaction Potentials at 298 K Half –Reaction E°(V) Hg2Cl2 + 2e-  2Hg + 2Cl- 0.27 AgCl + e-  Ag + Cl- 0.22 SO4

2- + 4H+ + 2e-  H2SO3 + H2O

0.20 Cu2+ +e-  Cu+ 0.16 2H+ + 2e-  H2 0.00 Fe3+ + 3e-  Fe

  • 0.04

Pb2+ + 2e-  Pb

  • 0.13

Sn2+ + 2e  Sn

  • 0.14

Ni2+ + 2e-  Ni

  • 0.23

PbSO4 + 2e-  Pb + SO4

2-

  • 0.35

Cd2+ +2e-  Cd

  • 0.40

Fe2+ + 2e-  Fe

  • 0.44

Cr3+ + e-  Cr2+

  • 0.50

Cr3+ + 3e-  Cr

  • 0.73

Zn2+ + 2e-  Zn

  • 0.76

2H2O + 2e-  H2 + 2OH-

  • 0.83

Mn2+ + 2e-  Mn

  • 1.18

Al3+ + 3e-  Al

  • 1.66

H2 + 2e-  2H-

  • 2.23

Mg2+ + 2e-  Mg

  • 2.37

La3+ + 3e-  La

  • 2.37

Na+ + e-  Na

  • 2.71

K+ +e-  K

  • 2.92

Li+ + e-  Li

  • 3.05
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Thermo of Electrochemistry

All reaction of referenced to:

H+(aq) + 2e-  H2(g) E˚ = 0 V

Things to remember:

 If you flip a reaction you change the sign of E˚  If you multiply a reaction by a constant DO NOT DO

ANYTHING TO E˚

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Exa xampl ple: e: F2 + 2e-  2F- E˚= 2.87 2F- F2 + 2e- E˚= -2.87 Example: F2 + 2e-  2F- E˚= 2.87 2F2 + 4e-  4F- E˚= 2.87 Not

  • te: The book gives you the following equation:

𝐹𝑑𝑓𝑚𝑚

°

= 𝐹° 𝑑𝑏𝑢ℎ𝑝𝑒𝑓 − 𝐹° 𝑏𝑜𝑝𝑒𝑓 DO NOT USE THIS EQUATION IF YOU ARE USING THE FLIPPING METHOD

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Student Question

Thermo of Electrochemistry

Calculate E° of the following cell Cr(s)|Cr3+(aq)||Br-(aq)|Br2(l)|Pt(s)

Helpful Information: Cr3+ + 3e-  Cr E° = -0.73 V Br2 + 2e-  2Br- E° = 1.09 V

a) 2.55 V b) 1.82 V c) -1.82 V d) -2.55 V e) None of the above

Bonus: Determine the balanced reaction for this cell.

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Thermo of Electrochemistry

Calculate the solubility product of Hg2Cl2 at 298.15 K. Solubility Product Ksp: Reactions of Interest: Hg2Cl2 + 2e-  2Hg + 2Cl- E˚= 0.27 V Hg2

2+ + 2e- 2Hg

E˚= 0.79 V

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Student Question

Thermo of Electrochemistry

The equilibrium constant for the following general reaction is 10. at 25°C. Calculate E° for the cell. X2(s) + Y+(aq)  X2+(aq) + Y(s) (unbalanced)

a) 0.015V b) 0.030 V c) 0.060 V d) 0.045 V e) None of the above

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Student Question

Nernst Equation

Calculate the emf, at 25ºC, of the cell Zn(s)|Zn2+(aq,1.5 M)||Fe3+(aq, 0.0010 M)|Fe(s)

Helpful Information: Fe3+ + 3e-  Fe(s) Eo = -0.04 V Zn2+ + 2e-  Zn(s) Eo = -0.76 V

a) 0.72V b) 0.66 V c) 2.1 V d) 0.69 V e) None of the above

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Batteries/ Fuel Cells

 Primary Cells: Galvanic cell with the reactants

sealed inside at the manufacturer. They can not be recharged.

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Type: Primary Cell/Dry Cell Anode Electrode: Zn Cathode Electrode: C Cathode Material: MnO2, Electrolyte: Carbon black, and NH4Cl Voltage :1.5 V

(A) Zn  Zn2+ + 2e- (C) 2MnO2 + 2NH4

+ + 2e-  Mn2O3 + 2NH3 + H2O

Zn+2NH4

++2MnO2 Zn2++Mn2O3+2NH3+ H2O

Zinc Carbon

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Batteries/ Fuel Cells

 Secondary Cell: Galvanic cell that must be

charged before it can be used.

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Type: Secondary Cell Anode Electrode: Pb Cathode Electrode: Pb covered with PbO2 Electrolyte: H2SO4 Voltage : 12 V

(A) Pb + HSO4

– PbSO4 + H+ + 2e–

(C) PbO2 + 3H+ + 2HSO4

– + 2e– PbSO4 + 2H2O

Pb + PbO2 + 2H+ + 2HSO4

–  2PbSO4 + 2H2O

Lead Acid

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Batteries/ Fuel Cells

 Fuel Cell: A Galvanic cell in which the reactants

are continuously supplied.

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Type: Fuel Cell Anode Electrode: Platinum Cathode Electrode: Platinum Electrolyte: KOH Voltage : 1.2 V

(A) 2H2 +4OH-  4H2O + 4e- (C) 4e- + O2 +2H2O  4OH- 2H2 + O2 2H2O

Alkaline

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Electrolytic Cell

General Difference between Galvanic and Electrolytic Cells

 Electroplating: The deposition of a thin film of

metal on an object.

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Galvanic Cell Electrolytic Cell E > 0 E < 0 No external power External power 2 Compartments 1 Compartment 2 Electrolytes 1 Electrolyte Room temp. & press. High temp. & press.

Not

  • te: Metal or graphite coated plastic is placed at the cathode.

Not

  • te: An aqueous solution of the salt of the plating material is the electrolyte.
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Electrolytic Cell

How much voltage must be supply to plate Zinc

  • nto copper?

Zn2+(aq) + 2e-  Zn(s) E˚ = -0.76 V Cu2+(aq) + 2e-  Cu(s) E˚ = 0.34 V

 Overpotential: Additional voltage over the emf

required to run a reaction.

22

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Electrolytic Cell

Faraday’s law of

Electrolysis: The amount of product formed or reactant consumed by an electrical current is stoichiometrically equivalent to the amount of electrons supplied.

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Student Question

Electrolytic Cell

What time is required to plate a metal tray (24.0 cm×12.0 cm) with a coating (thickness 0.00200 cm) of silver (density =10.54

𝑕 𝑑𝑛3) using a

current 7.65 A? Neglect the amount of silver required to coat the edges.

Helpful Information: The reaction of interest is Ag+ + e-  Ag, and 𝑁𝐵𝑕 = 107.9

𝑕 𝑛𝑝𝑚

a) 1420 s

b) 912 s

c) 1130 s d) 708 s e) None of the above

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Take Away From Chapter 11

 Big Idea: Electron transfer in a chemical reaction

is both material and concentration specific. If the process is spontaneous the transfer of electrons can be used to produce a current and drive electrical devices.

 Redox Reaction Review

 Be able to assign oxidation numbers  Be able to identify element oxidized/reduced  OIL RIG  Be able to identify oxidizing/reducing agent  Be able to balance redox equations (Ch. 4 - 83)  Acidic conditions (no OH-)  Basic conditions (no H+)

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Numbers correspond to end of chapter questions.

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Take Away From Chapter 11

 Galvanic Cell

Know that a galvanic cell is a spontaneous reaction (3)

Be able to draw a galvanic cell (19&21)

Identify anode

Identify cathode

Identify electron flow

Identify ion migration through salt bridge

 Thermo of Electrochemistry

 Know that for galvanic cell Ecell is positive  Be able to calculate E°cell for a system (22&24)  Be able to select the best oxidizing/reducing agent from a list

  • f standard potentials (25,26,27,28,29,&30)

 Be able to calculate ΔG, wmax , and K of a cell (42,43,44,45,

47,&100)

∆𝐻 = −𝑜𝐺𝐹 = 𝑥𝑛𝑏𝑦

𝐹° = 𝑆𝑈

𝑜𝐺 ln(𝐿)

26

Numbers correspond to end of chapter questions.

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Take Away From Chapter 11

 Nernst Equation

 Be able to calculate E at conditions other than the standard

state (10,55,56,59,65,&98)

 𝐹 = 𝐹° − 𝑆𝑈 𝑜𝐺 ln(𝑅)

 Batteries / Fuel Cell  Electrolytic Cell

 Know that an electrolytic cell is a non spontaneous reaction  Be able to calculate the products of electrolysis.

(72,73,74,75,&84)

 𝑜 = 𝐽𝑢

𝐺  Be able to determine when competitive reactions (water) will

effect electrolytic cell (81)

27

Numbers correspond to end of chapter questions.