Chem ch 17

driving force for (or energy accompanying) a specific type of charge transfer process (V = J/C)

Potential (voltage) =

transfer of electrons between redox reactants

Potential (voltage) = In an electrochemical cell

difference in potential between two half-cells

Cell potential (Ecell) =

Ecell = Ered,cathode – Ered,anode

E°cell = E°red,cathode – E°red,anode

half-cells are under standard-state conditions: 1 M, 1 bar, 298 K)

(standard cell potentials =

All reduction reactions are assigned a potential relative to the

standard hydrogen electrode (SHE) (0 V)

Standard Reduction Potentials Table: List of reduction half-reactions and

their associated voltage (relative to the SHE)

Calculate the cell potential (E°cell): Spontaneous is

Positive

Calculate the cell potential (E°cell): nonspontaneous

negative

E°cell > 0

reaction is spontaneous

E°cell < 0

reaction is nonspontaneous

What substance in this table is most easily reduced

Highest positive value

What substance in this table is most easily oxidized?

Lowest negative value

ΔG° = standard free energy change

(max work performed by a system)

ΔG° =

standard free energy change

n =

moles of e- transferred

F =

Faraday’s constant (96500 C/mol e-) (charge in Coulombs per mole of electrons)

E°cell =

standard cell potential

Potentials under Nonstandard Conditions

Ecell = E°cell – (𝑹𝑻/𝒏𝑭)𝒍𝒏𝑸 Q= reaction quotient

voltage imposed on a system to drive an otherwise nonspontaneous reaction (e.g. recharging batteries, electroplating (jewelry, utensils, auto parts))

Electrolysis =

Electrolysis = example

Downs cell (electrolysis of molten NaCl(l))

Electrolysis requires a positive potential of magnitude

greater than the (negative) cell potential!

Electrolysis: Complications in aqueous solution: For aqueous solutions of ionic compounds, anode and cathode half-reaction may involve electrolysis of

water species (H2O, H+, OH-) OR solute species (cations or anions from the ionic compound)

rate of flow for a charged species (A, ampere = 1 C/s)

Electrical current =

Quantifying Electrolysis Q= I t

Q = charge transferred

I = constant current (A, C/s)

t = time interval (s)

Quantifying Electrolysis Q= nF

Q = charge transferred

n = mol e-

F = Faraday’s constant (96500 C/mol e-)

Use these equations for “How much is produced?” or “How long will it take?”-type of problems

Q= nF or Q= I t

more positive E°red =

more readily reduced = better oxidizing agent

more negative E°red =

more readily oxidized = better reducing agent

Use the Nernst equation to find cell potentials under non-standard conditions

Ecell = E°cell – (𝑹𝑻/𝒏𝑭)𝒍𝒏𝑸

How do electrolytic cells compare to galvanic cells?

They’re the opposite! (involve a nonspontaneous reaction & require an applied potential)

chemical reaction that involves changes in the oxidation numbers/oxidation states for one or more elements

Redox reaction (reduction-oxidation reaction) =

charge an atom would have in a compound if it were ionic

(sum of the oxidation numbers for all atoms in the molecule = charge on the molecule)

Oxidation number =

In Ionic compounds the oxidation numbers =

ionic charges

Covalent compounds—pretend like

partial charges on molecules were fully transferred (like they were ionic)

Balancing Redox Reactions Key:

one atom must be oxidized and one atom must be reduced

oxidation =

increase in oxidation # (whatever is oxidized is the reducing agent)

reduction =

decrease in oxidation # (whatever is reduced is the oxidizing agent)

Balancing Redox Reactions: 8 Steps

1. Write the redox half-reactions.

2. Balance all species except H & O.

3. Balance half-reactions by adding H2O.

4. Balance half-reactions by adding H+ for H.

5. Balance charges with electrons (e-).

6. Balance e- by multiplying eqs. by coefficients.

7. Add the two reactions and simplify.

8. If rxn happens in basic solution, add OH- to

neutralize H+.

In electrochemistry, many reaction are aqueous…and are harder to balance by simple inspection. In these cases, use the

half-reaction method to help (i.e. begin with the oxidation and reduction reactions/equations separately).

device in which direct transfer of electrons is prevented by separation of reactants via an external circuit that connects them

Electrochemical cell =

electrochemical cell in which a spontaneous redox reaction takes place

Galvanic/voltaic cell =

(each beaker) containing “half” of the redox conjugate pair (single reactant)

half-cell =

where oxidation happens

anode =

where reduction happens

cathode =

tube filled with inert electrolyte solution that connects the two half-cells

salt bridge =

What is the purpose of the salt bridge?

maintains charge balance

Why is a salt bridge needed?

completes the electrical circuit

Cell notation/schematic

• Half-cell reactions are shown on opposite sides of the two parallel vertical lines (∥)

• Interfaces between components are separated by a single vertical line (⏐)

• Same-phase components are separated by commas

Cell notation/schematic Why does Pt(s) show up on one side?

inert electrode or active electrode

inert electrode =

needed when neither component of the half-cell can function as an electrode

active electrode =

electrode made of one material of the redox couple (e.g. Mg(s))

What is a galvanic/voltaic cell

Electrochemical cell in which a spontaneous redox reaction takes place

Ionic compound oxidation number The sum must be 0

if there is no net charge on the outside of the compound

Ionic compound oxidation number follow the

Number coming from the periodic table

Covalent compound oxidation number In elemental state, not compounded to anything…

Compound has an oxidation state of 0

Reducing agent is the

oxidized species

oxidizing agent is the

reduced species

A spontaneous redox reaction a copper wire is placed in

silver nitrate, The liquid turns blue and

Pt(s) is platinum what is it used for in the cell notation schematic

Conduct electricity in 1metal and aqueous solution

Not involved in the redox reaction facilitates electron flow

To start the process of figuring out a galvanic cell the anode will be

The species in the elemental state

When making a half cell the e- goes on the side that is more

positive

lead on the periodic table

Pb

Silver on the periodic table

Ag

Gold on the periodic table

Au

Tin on the periodic table

Sn

Iron on the periodic table

Fe

Copper on the periodic table

Cu

Cell potential equation

Ecell = E cath — E anode =

The better oxidizing agent is the

element further down on the standard reduction potential list

The better reducing agent is the

element further up on the standard reduction potential list

What two concentration based adjustments could you make to increase cell potential to the reaction?

Remove products or increase reactant concentrations. This will move the reaction into the forward direction