Chapter 16: Electrochemistry (TEST 3)

OXIDATION STATES: Group I metals have an oxidation state of what in all of their compounds?

+1 EX: Na = +1 in NaCl

OXIDATION STATES: Group II metals have an oxidation state of what in all of their compounds?

+2 EX: Mg = +2 in MgCl2

What is corrosion? Provide an example.

-Corrosion is usually defined as the degradation of metals by a naturally occurring electrochemical process. -Example: The formation of rust on iron, tarnish on silver, and the blue-green patina that develops on copper.

What is a cathode?

-Negatively charged electrode -the electrode at which reduction occurs

What is an anode?

-Positively charged electrode -the electrode at which oxidation occurs

Iron Rusting

-Redox Reaction: Fe ® Fe2+(aq) + 2e- O2 + 2H2O + 4e-® 4OH-(aq)+ 2e- •Presence of salt facilitates the redox reaction ØPrecipitation Reaction Fe2+(aq) + 2OH-(aq) ® Fe(OH)2(s)

•General Reduction Half Reaction: A + ne → B

-The more negative the standard reduction potential, the more reducing the product B is. -The more positive the standard reduction potential, the more oxidizing the reactant A is.

SALT BRIDGE EXPLAINED

-The salt bridge is a porous material consisting of a concentrated salt solution. •NaNO3 or KNO3 are frequently used. -As the oxidation occurs, a surplus of positive ions builds up at the anode. -Anions flow toward the anode to neutralize the build-up of positive charge. •-The area near the cathode becomes deficient in positive ions. •-Cations flow toward the cathode to "replace" the cations that are being consumed. -This flow of ions enables electrical neutrality to be achieved, the circuit to be complete, and current to flow!

What is the "salt bridge"?

-is a porous material consisting of a concentrated salt solution. -NaNO3 or KNO3 are frequently used.

What are the steps for balancing redox reactions?

1) Assign oxidation numbers. 2) Split reaction into two half-reactions: oxidation & reduction. 3) Balance two half-reactions by transferred electrons. 4) Balance all elements except H and O. 5) Balance O by adding H2O. 6) Balance H by adding H+.(in acidic condition) 7) Add OH− to neutralize the H+ ions (in basic condition). *Double Check: atoms and electrons must be balanced.

4.) Rust forms on iron through the redox reaction 4Fe + 3O2 2Fe2O3 . Identify the correct oxidizing agent. A. Iron is the oxidizing agent B. O2 is the oxidizing agent C. Iron is the reducing agent and also oxidizing agent.

B. O2 is the oxidizing agent

12.) In an electrolytic cell, reduction occurs at the _____. A. anode terminal B. cathode terminal C. negative terminal

B. cathode terminal

7.) A voltaic cell must be split into two separate cells, an oxidation and reduction half-cells. The two half-cells must be connected by a _____ so ions can freely flow between the two half-cells to balance the charges or the voltaic cell cannot function. A. metallic wire B. salt bridge C. membrane

B. salt bridge

Electrolytic Cells vs. Galvanic Cells

ELECTROLYTIC CELLS: the anode terminal is the positive terminal; the cathode terminal is the negative terminal. GALVANIC CELLS: the anode terminal is the negative terminal; the cathode terminal is the positive terminal. No matter which cells, electrons always flows from anode to cathode, Always oxidation at anode and reduction at cathode!

What is electrolysis?

Electrolysis is a process where an external circuit imposes a voltage sufficient to drive an otherwise nonspontaneous reaction to happen.

Balancing Oxidation-Reduction Reactions (Review of Chapter 7): LEO says GER

LEO: Loss of Electrons is Oxidation GER: Gain of Electrons is Reduction

Free metals tend to be reducing or oxidizing?

REDUCING -the more reactive the metal (with low electronegativity) the more reducing it will be.

Rechargeable (Secondary) Batteries

Rechargeable (Secondary) Batteries are based on conveniently reversible cell reactions that allow recharging by an external power source.

The E°red value of all other half reactions have been assigned relative to what?

SHE

Single-Use (primary) Batteries

Single-Use (primary) Batteries are designed for single-use applications and cannot be recharged.

OXIDATION STATES: Monatomic ions have an oxidation state that are equal to their ____________

charge EX: In NaCl Na = +1 and Cl = −1

Is E°red an extensive or intensive property?

intensive property

OXIDATION STATES: Free elements have an oxidation state of

oxidation state = 0. EX: Na = 0 and Cl2 = 0 in 2 Na(s) + Cl2(g)

is the reducing agent reduced or oxidized

oxidized

do the oxidizing agent and reducing agent appear as reactants or products in the overall redox reaction?

reactants EX: 2H2 + O2 2H2O 2H2 and O2 would be the agents

E°, ΔG°, and K are all measures of ........

reaction spontaneity

is the oxidizing agent reduced or oxidized?

reduced

OXIDATION STATES: In their compounds, nonmetals have oxidation states according to the pic

see pic FLUORINE -1 HYDROGEN +1 OXYGEN -2 GROUP 7A -1 GROUP 6A -2 GROUP 5A -3

Relationship Between E°, ΔG° and K

see slide #28 picture

The Nernst Equation

see slide #32

Balancing Oxidation-Reduction Reactions (Review of Chapter 7): Oxidation and Reduction occur _________________________

simultaneously (at the same time)

What allows one to determine what is being oxidized and what is being reduced?

the assignment of oxidation numbers

OXIDATION STATES: The sum of the oxidation states of all the atoms in a polyatomic ion equals what?

the charge on the ion EX: N = +5 and O = −2 in NO3-, (+5) + 3(−2) = −1

What are redox reactions?

the chemical processes involving the transfer of electrons between reactants

What is electrochemistry?

the study of redox reactions

Relationship Between E° and ΔG°

ΔG°=-nFE°cell ° indicates standard conditions (gases at 1 atm (or 1 bar); aqueous specices at 1 M) ΔG° is the standard free energy change E°cell is the standard cell potential n is the number of moles of electrons transferred in the balanced redox reaction. F is called the Faraday constant, this is the charge of a mole of electrons. F = 96485 J/mol.V -Notice that ΔG° and E° have opposite signs.

Galvanic Cell Summary

•"Compartmentalizing" the oxidation and reduction half reactions into half-cells, enables a spontaneous redox reaction to be a source of electrical energy. •In one half-cell, the anode, oxidation occurs. •In the other half-cell, the cathode, reduction occurs. •Each half-cell consists of an electrode dipped into an aqueous solution. •The half-cells are joined by an external wire (electron movement) and salt bridge (ion movement).

Around 1800, chemists began exploring ways to transfer electrons via what?

•Around 1800, chemists began exploring ways to transfer electrons via an external circuit rather than directly via intimate contact of redox reactants.

Balancing Oxidation-Reduction Reactions (Review of Chapter 7): REDUCTION

•Gain of electrons •Electrons are consumed •The species is reduced •Oxidation number decreases

16.3: Standard Reduction Potentials (Ered°)

•In a circuit, the flow of charge is a result of an electrical potential difference between two points in the circuit. •In a galvanic cell, this electrical potential difference (cell potential, Ecell ) is the driving force to move electrons from anode to cathode via an external wire. (Units of Ecell is V or mV)

Balancing Oxidation-Reduction Reactions (Review of Chapter 7): OXIDATION

•Loss of electrons •Electrons are produced •The species is oxidized •Oxidation number increases

CELL NOTATION: •Describes what happens in a galvanic cell. Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)

•Oxidation on the left (Cu oxidized to Cu2+) •Reduction on the right (Ag+ reduced to Ag) •Single vertical line is a phase boundary •Liquid-metal or liquid-gas, etc. •Double line is the salt bridge •Sometimes the concentration of the ion(s) is included. Cu(s)│1M Cu(NO3)2(aq)║1M AgNO3(aq)│Ag(s)

Standard Potentials (E°)

•Standard potentials (E°) are measured with •All aqueous concentrations at 1M •The pressure of all gases at 1 atm •Temp. held constant (usually at 25 °C)

Standard Hydrogen Electrode (SHE)

•The E° for a half reaction cannot be measured, only E°cell of the entire redox reaction can be measured. •A standard reduction potential (E°red , or simply E°) has been set. •The value of E° for the reduction of 2H+ ions to H2(g) has been assigned to be 0.000 V. 2H+ (aq, 1M) + 2e- ® H2(g, 1 atm) •This is called the standard hydrogen electrode (SHE).

Effect of Concentration on Potential (when potentials aren't in standard conditions) •Voltage will increase and the reaction will become more spontaneous if ......

•The concentration of a reactant is increased •The concentration of a product is decreased

Effect of Concentration on Potential (when potentials aren't in standard conditions) •Voltage will decrease and the reaction will become less spontaneous if

•The concentration of reactant is decreased •The concentration of product is increased

Reduction results in _______________

•a decrease in oxidation number. EX: 2Ag+(aq) + 2e- → 2Ag(s)

Oxidation results in ____________

•an increase in oxidation number. EX: Cu (s) → Cu2+ (aq) + 2 e-

In the last two centuries, the field of electrochemistry has evolved to what?

•to yield rechargeable batteries and refillable batteries for electric vehicles

Relationship between E° and K: •Redox reactions eventually reach a state of equilibrium.

-see equations on slide #29 •This equation applies only at 25 °C •Note that: •If E° is positive, K is greater than 1 •If E° is negative, K is less than 1

If the standard potential (E°cell) of a redox reaction is negative

-then the reaction is non-spontaneous at standard conditions.

If the standard potential (E°cell) of a redox reaction is positive

-then the reaction is spontaneous at standard conditions. -This is why E°cell for galvanic cells are always positive.

OXIDATION STATES: The sum of the oxidation states of all the atoms in a compound (A compound is a substance formed when two or more chemical elements are chemically bonded together) is __________

0 EX: Na = +1 and Cl = −1 in NaCl, (+1) + (−1) = 0

14.) The cathode terminal in an electrolytic cell is the _____ terminal. A. (-) B. (+) C. neutral

A. (-)

15.) Based on the reduction potential data, what is the standard cell potential for the following electrochemical cell reaction: Zn(s) + Cu2+(aq)Zn2+(aq) + Cu(s)? E°red = -0.763 V for Zn2+(aq) + 2e- Zn(s) E°red = +0.340 V for Cu2+(aq) + 2e- Cu(s) A. +1.103 V B. -0.423 V C. +0.423 V

A. +1.103 V

18.) The standard cell potential is +0.460 V for the following electrochemical cell reaction: Cu(s) + 2Ag+(aq)Cu2+(aq) + 2Ag(s). What is the ∆G°rxn? A. -88.8 kJ/mol B. -44.4 kJ/mol C. +44.4 kJ/mol

A. -88.8 kJ/mol

23.) What is the oxidation state of P in K3PO4? A. 5 B. 4 C. 3

A. 5

5.) Which is the reducing agent in the combustion of methanol? A. CH3OH B. O2 C. CO2

A. CH3OH

1.) Assuming the following pair of half-reactions below takes place in an acidic solution, write a balanced equation for the overall reaction.Ca Ca2++2e- ; 2e- +F2 2F- A. Ca + F2 Ca2+ + 2F- B. Ca2+ + 2F- Ca + F2 C. Ca + F2 Ca2+ + 2F- +2e-

A. Ca + F2 Ca2+ + 2F-

26.) Determine the cell potential under the stated conditions for the electrochemical reaction described. State whether each is spontaneous or nonspontaneous under the set of conditions at 298.15 K. Hg(l) S2 (aq, 0.10 M ) 2Ag (aq, 0.25 M ) 2Ag(s) HgS(s) HgS(s) +2e Hg(l) +S2-(aq) Eo = -0.70 V Ag+ (aq) + e Ag(s) Eo = 0.7996 V A. Ecell = 1.43 V, spontaneous B. Ecell = 1.50 V, spontaneous C. Ecell = 0.0996 V, spontaneous

A. Ecell = 1.43 V, spontaneous

16.) Based on the following reduction potential data, which is the strongest reducing agent? E°red = +1.78 V for H2O2(aq) + 2H+(aq) + 2e-2H2O(l) E°red = +0.800 V for Ag+(aq) + e- Ag(s) E°red = -0.130 V for Pb2+(aq) + 2e- Pb(s) A. Pb(s) B. Ag(s) C. H2O(l)

A. Pb(s)

29. Define electrolysis. A. Process using electrical energy to cause a nonspontaneous process to occur B. Process using electrical energy to cause a spontaneous process to occur C. Process using chemical energy to cause a nonspontaneous process to occur

A. Process using electrical energy to cause a nonspontaneous process to occur

8.) In a voltaic cell, oxidation occurs at the _____. A. anode terminal B. cathode terminal C. positive terminal

A. anode terminal

13.) Electron flow in an electrolytic cell is always from _____ to _____. A. anode to cathode B. cathode to anode C. anode to negative terminal

A. anode to cathode

9.) Electron flow in a voltaic cell is always from _____ to _____. A. anode to cathode B. cathode to anode C. anode to negative terminal

A. anode to cathode

11.) In a voltaic cell, electrons are received at the _____. A. cathode terminal B. negative terminal C. anode terminal

A. cathode terminal

What can all metals under H do?

All metals under H can react with acids to produce hydrogen gas

10.) The cathode terminal in a voltaic cell is the _____ terminal. A. (-) B. (+) C. neutral

B. (+)

22.) What is the oxidation state of oxygen in K2O2? A. 0 B. -1 C. -2

B. -1

19.) What is the ∆G°rxn for a 2 electron transfer electrochemical cell with a standard cell potential of +0.770 V. A. -74.3 kJ/mol B. -149 kJ/mol C. +74.3 kJ/mol

B. -149 kJ/mol

2.) Assuming the following pair of half-reactions below takes place in an acidic solution, write a balanced equation for the overall reaction. Li Li++e- ; 2e- +Cl2 2Cl- A. Li + Cl2 Li+ + 2Cl- B. 2Li + Cl2 2Li+ + 2Cl- C. 2Li + Cl2 Li+ + 2Cl-

B. 2Li + Cl2 2Li+ + 2Cl-

3.) Assuming the following pair of half-reactions below takes place in an acidic solution, write a balanced equation for the overall reaction. Ag Ag+ + e- ; MnO4- +4H+ + 3e- + MnO2 + 2H2O A. Ag + MnO4- +4H+ MnO2 + 2H2O B. 3Ag +MnO4- +4H+ 3Ag+ + MnO2 + 2H2O C. Ag + MnO4- +4H+ Ag+ + MnO2 + 2H2O

B. 3Ag +MnO4- +4H+ 3Ag+ + MnO2 + 2H2O

21.) Cell phones batteries are typically rechargeable. How do they differ from non-rechargeable batteries? A. A rechargeable battery has the ability to undergo a double replacement reaction while non-rechargeable batteries do not. B. A rechargeable battery has the ability to undergo a reversible redox reaction while non-rechargeable batteries do not. C. A rechargeable battery has the ability to undergo a redox reaction while non- rechargeable batteries do not.

B. A rechargeable battery has the ability to undergo a reversible redox reaction while non-rechargeable batteries do not.

17.) Based on the following reduction potential data, which is the strongest oxidizing agent? E°red = +1.78 V for H2O2(aq) + 2H+(aq) + 2e- 2H2O(l) E°red = +0.800 V for Ag+(aq) + e- Ag(s) E°red = -0.130 V for Pb2+(aq) + 2e- Pb(s) A. Pb2+(aq) B. H2O2(aq) C. Ag+(aq)

B. H2O2(aq)

20.) Which strong acid is commonly used in automobile batteries? A. HCl B. H2SO4 C. HNO3

B. H2SO4

27.) Use the standard reduction potenital to calculate equilibrium constants for the following reaction. CdS(s) Cd2+ (aq) +S2- (aq) at 377 K Cd2+ (aq) +2e Cd (s) Eo = -0.4030 V CdS(s) +2e Ca (s) +S2- (aq) Eo = -1.17 V A. K=1.7 x10 -10 nonspontaneous B. K=3.1 x10-21 nonspontaneous C. K=1.5 x10-8 nonspontaneous

B. K=3.1 x10-21 nonspontaneous

24.) Which of the following is a strong oxidizing agent? A. N2 B. O2 C. H2

B. O2

30. Based on the given standard free energy change and electron stoichiometry values below, calculate a corresponding standard cell potential. (see HW) A. B. C. D.

C.

28.) Use the standard reduction potenital to calculate equilibrium constants for the following reaction. (see HW) A. K=1.0 x10 -10 spontaneous B. K=1.6 x10-11 nonspontaneous C. K=1.0 x10-14 nonspontaneous

C. K=1.0 x10-14 nonspontaneous

25.) Which of the following is a strong reducing agent? A. F2 B. N2 C. Mg

C. Mg

6.) Automobiles from northern states tend to show more signs of rust formation than southern states. The reason for this is that the rate of corrosion is affected by _____, which is more common during winter months especially in northern states. A. the cold air and clouds B. rain clouds C. snow, ice, and salt mixtures

C. snow, ice, and salt mixtures

A redox reaction is spontaneous if:

E°cell is positive. ΔG° is negative. K is greater than 1.

Free nonmetals tend to be reducing or oxidizing?

OXIDIZING -The more reactive the nonmetal (with high electronegativity), the more oxidizing it will be.

•Something must cause the species to be oxidized.

This is the oxidizing agent (oxidant)

Something must cause the species to be reduced.

This is the reducing agent (reductant)

What does is mean to be a free metal or free nonmetal?

When electricity flows, the electrons are considered "free" only because there are more electrons than there should be

all metals above H ____________________ wait this don't make any sense cause there are no metals above H on the periodic table??

all metals above H cannot react with acids to produce H2. EX?: Cu can react with HNO3 but only produce NO, not H2.