redox reactions and electrochemistry

Redox Reactions and Electrochemistry

9.3 Oxidation Numbers ( only first half, assigning ox numbers)

10.1 Galavanic or Voltaic Cellsa) Anode/Cathode/Salt Bridge

b) Cell Notations

c) Determining Cell Potential/Cell Voltage/Electromotive force (emf)

Application

V. Corrosion

VI. Batteries

VII. Fuel Cells

VIII.Electrolytic Cells

Electrochemistry: Interconversion of electrical and chemical energy using redox reactions

Redox (Oxidation-Reduction) Reaction: Type of electrontransfer reaction. One substance gives up electrons;

the other accepts electrons.

OIL RIG

•Oxidation Half-Reaction; Oxidation Involves Loss of electrons

•Reduction Half-Reaction; Reduction Involves Gain of electrons

gge

e

Net Redox Rxn; 2Mg + O2 -> 2 Mg+2 + 2 O-2

Oxidation numberThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferredto the more electronegative atom.

1. Oxidation number equals ionic charge for monoatomic ions in ionic compound

2. Metal ions in Family A have one, positive oxidation number; Group IA metals are +1, IIA metals are +2

Li+, Li = +1; Mg+2, Mg = +2

4.4

CaBr2; Ca = +2, Br = -1

Oxidation number,continuedThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferredto the more electronegative atom.

3. The oxidation number of a transition metal ion is positive, but can vary in magnitude.

4. Nonmetals can have a variety of oxidation numbers,both positive and negative numbers which can vary in magnitude.

4.4

5. Free elements (uncombined state) have an oxidation number of zero. Each atom in O2, F2, H2, Cl2, K, Be has the same oxidation number; zero.

7. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.

IF; F= -1; I = +1

8. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1 or when it’s in elemental form (H2; oxidation # =0).

HF; F= -1, H= +1

NaH; Na= +1, H = -1

6. The oxidation number of fluorine is always –1. (unless fluorine is in elemental form, F2)

H2O ; H=+1, O= -2SO3; O = -2; S = +6

9. The oxidation number of oxygen is usually –2. In H2O2 and O2

2- it is –1, in elemental form (O2 or O3) it is 0.

HCO3-

O = -2 H = +1

3x(-2) + 1 + ? = -1

C = +4

Oxidation numbers of all the atoms in HCO3

- ?

4.4

NaIO3

Na = +1 O = -2

3x(-2) + 1 + ? = 0

I = +5

IF7

F = -1

7x(-1) + ? = 0

I = +7

K2Cr2O7

O = -2 K = +1

7x(-2) + 2x(+1) + 2x(?) = 0

Cr = +6

Oxidation numbers of all the elements in the following ?

4.4

+1+1

-2-6+5

+5 +1+2

-2-14+12

+6 +1+6

-2-2-4

-2 -1-1+1

+1H N O 3 C2H6OK2Cr2O7 AgI

+1+2

-2-8+5

+5H2PO4

–

Application to Electrochemistry• An electric cell converts chemical energy into electrical energy

– Alessandro Volta invented the first electric cell but got his inspiration from Luigi Galvani. Galvani’s crucial observation was that two different metals could make the muscles of a frog’s legs twitch. Unfortunately, Galvani thought this was due to some mysterious “animal electricity”. It was Volta who recognized this experiment’s potential.

– An electric cell produces very little electricity, so Volta came up with a better design:

• A battery is defined as two or more electric cells connected in series to produce a steady flow of current– Volta’s first battery consisted of several bowls of brine (NaCl(aq))

connected by metals that dipped from one bowl to another– His revised design, consisted of a sandwich of two metals

separated by paper soaked in salt water.

• Alessandro Volta’s invention was an immediate technological success because it produced electric current more simply and reliably than methods that depended on static electricity.

• It also produced a steady electric current –something no other device could do.

• Electric cells are composed of two electrodes – solid electrical conductors and at least one electrolyte (aqueous electrical conductor)

• In current cells, the electrolyte is often a moist paste (just enough water is added so that the ions can move). Sometimes one electrode is the cell container.

• The positive electrode is defined as the cathode and the negative electrode is defined as the anode

– The electrons flow through the external circuit from the anode to the cathode.To test the voltage of a battery, the red(+) lead is connected to the cathode (+ electrode), and the black(-) lead is connected to the anode (- electrode)

• A voltmeter is a device that measures the energy difference, per unit charge, between any two points in an electric circuit (called electric potential difference)– I.e. A 9V battery releases 6X as much energy compared with the electrons from a 1.5V

battery.

– The voltage of a cell depends mainly on the chemical composition of the reactants in the cell

• An ammeter is a device that measures the rate of flow of charge past a point in an electrical circuit (called electric current)– The larger the electric cell, the greater current that can be produced

Electric potential difference is like the

potential energy difference between 1kg of water at the top of the dam and 1kg of water at the bottom of the

dam.

Electric potential difference is like the

potential energy difference between 1kg of water at the top of the dam and 1kg of water at the bottom of the

dam.

Electric current (or flow of electrons) is like the flow of the water. A larger drain

would be like a larger electric cell, allowing more water (or

electrons) to flow.

Electric current (or flow of electrons) is like the flow of the water. A larger drain

would be like a larger electric cell, allowing more water (or

electrons) to flow.

Primary cells cannot be recharged, but are relatively

inexpensive

Primary cells cannot be recharged, but are relatively

inexpensive

Secondary cells can be recharged using electricity,

but are expensive

Secondary cells can be recharged using electricity,

but are expensive

Fuel cells produce electricity by the reaction

of a fuel that is continuously supplied.

More efficient, and used for NASA vehicles, but still too expensive for general or commercial

applications

Fuel cells produce electricity by the reaction

of a fuel that is continuously supplied.

More efficient, and used for NASA vehicles, but still too expensive for general or commercial

applications

Gets Smaller -> <- Gets Larger

Voltaic Cell Animation

Anode; Site of OxidationCathode; Site of Reduction

AnOx or both vowelsRed Cat or both consonants

Direction of electron flow; anode to cathode (alphabetical)

Salt Bridge; Maintains electrical neutrality+ ion migrates to cathode- ion migrates to anode

Cell Notation

1. Anode

2. Salt Bridge

3. Cathode

Anode | Salt Bridge | Cathode

| : symbol is used whenever there is a different phase

19.2

Cell Notation

Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)

[Cu2+] = 1 M & [Zn2+] = 1 M

Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)anode cathode

Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s)

anode cathodeSalt bridge

More detail..

K(NO3)

Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)

Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt

Electrochemical Cells

The difference in electrical potential between the anode and cathode is called:

• cell voltage

• electromotive force (emf)

• cell potential

000reductionoxidationCell EEE

UNITS: Volts Volt (V) = Joule (J) Coulomb, C

Standard Electrode Potentials

Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.

V

Standard hydrogen electrode (SHE)

eatm

Reduction Reaction

Determining if Redox Reaction is Spontaneous

• + E°CELL ; spontaneous reaction

• - E°CELL; nonspontaneous reaction

More positive E°CELL ; stronger oxidizing agent ormore likely to be reduced

• E0 is for the reaction as written

• The half-cell reactions are reversible

• The sign of E0 changes when the reaction is reversed

• Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0

• The more positive E0 the greater the tendency for the substance to be reduced

Standard Potentials (cont.)• If in constructing an electrochemical cell, you need to

write the reaction as a oxidation instead of a reduction, the sign of the 1/2 cell potential changes.

Zn+2 + 2e- Zn E° = -0.76 V

Zn Zn+2 + 2e- E° = +0.76 V

• 1/2 cell potentials are intensive variables. As such, you do NOT multiply them by any coefficients when balancing reactions.

Writing Galvanic Cells (cont.)

Shorthand Notation

Zn|Zn+2||Cu+2|Cu

Anode Cathode

Salt bridge

Predicting Galvanic Cells

• Given two 1/2 cell reactions, how can one construct a galvanic cell?

• Need to compare the reduction potentials of the two half cells.

• Turn the reaction for the weaker reduction (smaller E°1/2) and turn it into an oxidation. This reaction will be the anode, the other the cathode.

Predicting Galvanic Cells (cont.)

• Example. Describe a galvanic cell based on the following:

Ag+ + e- Ag E°1/2 = 0.80 V

Fe+3 + e- Fe+2 E°1/2 = 0.77 V

Weaker reducing agent – turn it around

Ag+ + Fe+2 Ag + Fe+3 E°cell = 0.03 V

E°cell > 0….cell is galvanic

Another Example

• For the following reaction, identify the two half cells, and use these half cells to construct a galvanic cell

3Fe+2(aq) Fe(s) + 2Fe+3(aq)+2 0 +3 oxidation

reduction

Fe+2(aq) + 2e- Fe(s) E° = -0.44 V

Fe+3(aq) + e- Fe+2(aq) E° = +0.77 V

Another Example (cont.)

Fe+2(aq) + 2e- Fe(s) E° = -0.44 V

Fe+3(aq) + e- Fe+2(aq) E° = +0.77 V

weaker reduction – turn it around

Fe(s) Fe+2(aq) + 2e- E° = +0.44 V

2 x

2Fe+3(aq) + Fe(s) 3Fe+2(aq) E°cell = 1.21 V

Corrosion – Deterioration of Metals by Electrochemical Process

Corrosion – Deterioration of Metals by Electrochemical Process

Cathodic Protection

Abbreviated Standard Reduction Potential Table

Batteries

19.6

Leclanché cell

Dry cell

Zn (s) Zn2+ (aq) + 2e-Anode:

Cathode: 2NH4+ (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)

Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

Batteries

Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-Anode:

Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)

Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

Mercury Battery

19.6

Batteries

19.6

Anode:

Cathode:

Lead storagebattery

PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e- PbSO4 (s) + 2H2O (l)

Pb (s) + SO42- (aq) PbSO4 (s) + 2e-

Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq) 2PbSO4 (s) + 2H2O (l)

Fuel Cell vs. Battery

• Battery; Energy storage device– Reactant chemicals already in device

– Once Chemicals used up; discard (unless rechargeable)

• Fuel Cell; Energy conversion device– Won’t work unless reactants supplied

– Reactants continuously supplied; products continuously removed

Fuel Cell

A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning

Anode:

Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)

2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-

2H2 (g) + O2 (g) 2H2O (l)

Types of Electrochemical Cells

• Voltaic/Galvanic Cell; Energy released from spontaneous redox reaction can be transformed into electrical energy.

• Electrolytic Cell; Electrical energy is used to drive a nonspontaneous redox reaction.