Chapter 18Electrochemistry

Redox Reaction• Elements change oxidation number

e.g., single displacement, and combustion, some synthesis and decomposition

• Oxidation--oxidation number increases

• Reduction--oxidation number decreases Both must occur in a reaction--two half reactions

• oxidizing agent is reactant molecule that causes oxidation contains element reduced

• reducing agent is reactant molecule that causes reduction contains the element oxidized

Rules for Assigning Oxidation States1. Free elements have an oxidation state = 0

2. Monatomic ions have an oxidation state equal to their charge.

3. The sum of the oxidation states of all the atoms in a compound is 0.

4. The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.

5. The oxidation number of fluorine is always -1 in compounds with other elements.

Rules for Assigning Oxidation States

6. Chlorine, bromine and iodine always have oxidation numbers of -1 except when bonded to O or F.

7. The oxidation number of oxygen is almost always -2; the oxidation number of hydrogen is almost always +1.

Exceptions:

--When oxygen is in the form of a peroxide (O22-), the oxidation number is -1.

--When hydrogen forms a binary compound with a metal, the oxidation number is -1 and the compound is called a hydride.

Oxidation and Reduction

• Oxidation occurs when an atom’s oxidation state increases during a reaction

• Reduction occurs when an atom’s oxidation state decreases during a reaction

CH4 + 2 O2 → CO2 + 2 H2O-4 +1 0 +4 –2 +1 -2

oxidationreduction

Reducing agent Oxidizing agent

Identify the element that is oxidized and the element that is reduced in each of the following reactions. What is the oxidizing and the reducing agent in each reaction?

3 H2S + 2 NO3– + 2 H+ S + 2 NO + 4 H2O

MnO2 + 4 HBr MnBr2 + Br2 + 2 H2O

Common Oxidizing AgentsOxidizing Agent Product when Reduced

O2 O-2

H2O2 H2O

F2, Cl2, Br2, I2 F-1, Cl-1, Br-1, I-1

ClO3-1 (BrO3

-1, IO3-1) Cl-1, (Br-1, I-1)

H2SO4 (conc) SO2 or S or H2S

SO3-2 S2O3

-2, or S or H2S

HNO3 (conc) or NO3-1 NO2, or NO, or N2O, or N2, or NH3

MnO4-1 (base) MnO2

MnO4-1 (acid) Mn+2

CrO4-2 (base) Cr(OH)3

Cr2O7-2 (acid) Cr+3

Common Reducing AgentsReducing Agent Product when Oxidized

H2 H+1

H2O2 O2

I-1 I2

NH3, N2H4 N2

S-2, H2S S

SO3-2 SO4

-2

NO2-1 NO3

-1

C (as coke or charcoal) CO or CO2

Fe+2 (acid) Fe+3

Cr+2 Cr+3

Sn+2 Sn+4

metals metal ions

Balancing Redox Reactions

1. Assign oxidation numbers--determine element oxidized and element reduced

2. Separate the reaction into oxidation and reduction half- reactions.3. Balance half-reactions by mass

a. First balance elements other than H and Ob. Balance O using H2O c. Balance H using H+

4. Balance each half-reaction by charge by adding electrons to the reactants side of the reduction and the product side of the oxidation.

5. Multiply half-reactions by integers to make # electrons the same in both half-reactions

° Add half-reactions and cancel the electrons to produce a balanced equation.

° For reactions that occur in acidic solutions, skip to step 9.° For reactions that occur in basic solutions, add the same # of OH- as

H+ to both sides of the equation. 9. Check that reaction is balanced for mass and charge.

Practice - Balance the Equation H2O2 + KI + H2SO4 K2SO4 + I2 + H2O

Practice - Balance the Equation H2O2 + KI + H2SO4 K2SO4 + I2 + H2O+1 -1 +1 -1 +1 +6 -2 +1 +6 -2 0 +1 -2

oxidationreduction

ox: 2 I-1 I2 + 2e-1

red: H2O2 + 2e-1 + 2 H+ 2 H2Otot 2 I-1 + H2O2 + 2 H+ I2 + 2 H2O

1 H2O2 + 2 KI + H2SO4 K2SO4 + 1 I2 + 2 H2O

Electric Current Flowing Directly Between Atoms

Redox Reactions & Current

• Redox reactions involve the transfer of electrons from one substance to another.

• Therefore, redox reactions have the potential to generate an electric current.

• In order to harness the energy produced by moving electrons, we need to separate the half reactions.

Voltaic (or Galvanic) Cell

Electrochemical Cells• Electrochemistry -- the study of redox reactions

that produce or require an electric current.

• The conversion between chemical energy and electrical energy is carried out in an electrochemical cell

• Spontaneous redox reactions take place in a voltaic cell (galvanic cell).

• Non-spontaneous redox reactions can be made to occur in an electrolytic cell by the addition of electrical energy.

Electrodes• Anode

electrode where oxidation occurs

• Cathodeelectrode where reduction occurs

Voltaic (Galvanic) Cell

the salt bridge is required to complete the circuit and maintain charge balance

Current and Voltage• Current is the number of electrons that flow through the

system per secondunit = Ampere1 A of current = 1 Coulomb/second1 A = 6.242 x 1018 electrons/secondElectrode surface area dictates the number of electrons that can

flow• Potential difference is the difference in potential energy

between the reactants and products (between electrodes)unit = Volt1 V of force = 1 J (of energy)/Coulomb (of charge)The voltage that drive electrons through the external circuitAmount of force pushing the electrons through the wire is called

the electromotive force, emf

Cell Potential• The difference in potential energy between the

electrodes in a voltaic cell is called the cell potential

• The cell potential depends on the relative ease with which the oxidizing agent is reduced at the cathode and the reducing agent is oxidized at the anode

• The cell potential under standard conditions is called the standard emf, E°cell

25°C, 1 atm for gases, 1 M concentration of solutionsum of the cell potentials for the half-reactions

Cell Notation• Shorthand description of Voltaic cell

electrode | electrolyte || electrolyte | electrode• Oxidation half-cell on left, reduction half-cell on the

right• | = phase barrier

if multiple electrolytes in same phase, a comma is used rather than |

often use an inert electrode

• || = salt bridge

Fe(s) | Fe2+(aq) || MnO4(aq), Mn2+(aq), H+(aq) | Pt(s)

Standard Reduction Potential• The cell potential cannot be measured for a

half-reaction

• We need to compare them to an arbitrary standard.

• We select as a standard half-reaction the reduction of H+ to H2 (or the oxidation of H2 to H+) standard hydrogen electrode, SHEStandard conditions [H+]=1M, 25˚CPotential difference (E˚)= 0V

Half-Cell Potentials• SHE reduction potential is defined to be exactly 0 V

• Reduction potentials are tabulated for many half-reactions

• Change sign to get oxidation potential

• E°cell = E°oxidation + E°reduction

E°oxidation = E°reduction

When adding E° values for the half-cells, do not multiply the half-cell E° values, even if you need to multiply the half-reactions to balance the equation

Calculating E˚ for a cell reaction

Calculate E˚ for the reaction:

Cu2+ (aq) + Fe2+(aq) Fe2+(aq) + Cu(s)

1) Break up the redox reaction into half-reactions

2) Find the reduction potential (E˚)for each half-reaction

3) Change the sign on E˚ for the oxidation half-reaction

• Add E˚(anode)+ E˚(cathode)

• If the result is negative, the reaction occurs in the opposite direction.