Chapter 20: Electrochemistry

Question 1

Oxidation

Answer 1

Loss of electrons

Question 2

Reduction

Answer 2

Gain of electrons

Question 3

Balancing oxidation-reduction equations: Acidic solutions

Answer 3

1) Assign oxidation states and identify substances being oxidized or reduced
2) Separate overall reaction into two half-reactions
3) Balance each reaction: (a) balance all elements other than H&O, (b) balance O by adding H2O, (c) balance H by adding H+
4) Balance half reaction wrt charge by adding e-
5) Make number e- in both half-reactions equal by multiplying one or both by small whole number
6) Add half-reactions; cancel electrons as necessary
7) Verify

Question 4

Balancing redox reactions: Basic solutions

Answer 4

1) Assign oxidation states; identify substances being oxidized/reduced
2) Separate into half-reactions
3) Balance all elements except H&O, balance O with H2O, balance H by adding H+, neutralize H+ by adding enough OH- to neutralize each H+; add sam number OH- to each side of the equation
4) Balance wrt charge (electrons)
5) Make number of electron equal for both
6) Add equations

Question 5

Electrical current

Answer 5

Flow of electric charge

Question 6

Electrochemical cell definition

Answer 6

Generation of electricity through redox reactions carried out in this device

Question 7

Voltaic (Galvantic) cell definition

Answer 7

Electrochemical cell that produces electrical current from a spontaneous chemical reaction

Question 8

Electrolytic cell definition

Answer 8

Consumes electrical current to drive non spontaneous reactions

Question 9

Voltaic Cell Components

Answer 9

• Half-cell (one of the two electrodes)
• Anode: Electrode where oxidation occurs
• Cathode: Electrode where reduction occurs (electrons flow into cathode)
• Salt bridge (inverted, U-shaped tube with strong electrolytes to connect half-cells)

Question 10

Ampere

Answer 10

(A) AKA amps; measurement of electrical current

Question 11

Potential difference and its significance

Answer 11

• Measure of difference in PE (joules) per unit charge (coulombs)
• Large potential difference corresponds to strong tendency for electron flow

Question 12

Cell potential (Ecell or cell emf)

Answer 12

• Potential difference between two electrodes
• Depends on relative tendencies of cell to undergo oxidation and reduction; concentrations of reactants and products in the cell and temperature

Question 13

E˚cell

Answer 13

• Standard cell potential (1M, 1 atm, or 25˚C)

E˚cell = E˚final - E˚initial | E˚cathode - E˚anode

Question 14

Cell diagram/line notation

Answer 14

• Represent electrochemical cell with compact notation
  1) Write oxidation of half-rxn. on the left; reduction on right. Double vertical line represents salt bridge.
  2) Substances in different phases are separated by single vertical line
  3) For some redox rxns, reactants/products of one or both of half-rxn. may be in same phase; separate reactions and products from each other with comma in line diagram

Question 15

Standard electrode potential

Answer 15

Electrode in each half-cell’s own individual potential

Question 16

Standard hydrogen electrode

Answer 16

• Half-cell electrode normally chosen to have potential of zero

Question 17

Summary of Standard Electrode Potentials

Answer 17

• Electrode in any half-cell with greater tendency to undergo reaction is positively related to SHE
• Electrode in any half-cell with lesser tendency negatively charged
• Cell potential of any electrochemical cell is difference between electrode potentials of cathode and anode
• Ecell is + for spontaneous reactions and - for non spontaneous reactions

Question 18

Calculating Standard Potential

Answer 18

1) Separate into oxidation and reduction reactions
2) Use standard electrode potentials; add half reactions, calculate E˚cell
3) E˚cell = E(cathode) - E(anode)

Question 19

Predicting Spontaneous Reactions

Answer 19

• Half-reaction with more positive electrode potential will attract more electrons and undergoes reduction
• Half-reaction with more negative repels and undergoes oxidation
• Any reduction reaction is spontaneous when paired with reverse of any reactions listed below on table

Question 20

Predicting whether metal will dissolve in acid

Answer 20

• Most acids dissolve metals by reduction of H+ ions to hydrogen gas and corresponding oxidation of metal to its ion
• Generally metals whose half-reactions are listed below reduction of H+ to H2 dissolve; metals above do not

Question 21

Faraday’s Constant

Answer 21

• Quantifies charge that flows in an electrochemical reaction; change in charge of coulombs of 1 mol of electrons (C/e-)
  ∆G˚ = -nFE˚cell

Question 22

Relating E˚cell and K

Answer 22

E˚cell = 0.0592V/n(logK)

Question 23

Cell Potential and Concentration (Nernst)

Answer 23

Cell = E˚cell - (0.0592V/n)(logQ)

Question 24

Nernst Equation:
Q = 1
Q < 1
Q > 1

Answer 24

• Ecell = E˚cell
• Ecell > E˚cell
• Ecell < E˚cell

Question 25

Dry cell battery

Answer 25

Fairly common; do not contain large amount of liquid

Question 26

Lead-Acid Storage Batteries

Answer 26

Found in most cars; 6 electrochemical cells wired in series

Question 27

Fuel cells

Answer 27

Reactants constantly flow through battery; generate electrical current as they undergo redox (fuel provided from external source)

Question 28

Electrolysis

Answer 28

Electrical current drives otherwise non spontaneous reaction (electrical current supplied causes reverse reaction to occur)

Question 29

Summary:

• Electrochemical cells
• Voltaic cells
• Electrolytic cells

Answer 29

• Oxidation at anode; reduction at cathode
• Anode is source of electrons and has negative charge; cathode draws electrons and has positive charge
• Electrons drawn away from anode; which must be connected to positive terminal of external power source (+), electrons driven to the cathode; must be connected to negative terminal of the power source (-)
• -> In all, cathode is sure of reduction and anode is site of oxidation

Question 30

Electrolysis of pure molten salt

Answer 30

Anion is oxidized and cation is reduced

Question 31

Electrolysis of Mixtures of Cations or Anions

Answer 31

• Cation with more positive electrode potential is reduced first
• Anion most easily be oxidized (one with more negative electrode potential) is oxidized first

Question 32

Electrolysis of Aqueous Solutions

Answer 32

• Complicated by electrolysis of water itself
• Cations of active metals cannot be reduced from aqueous solutions by electrolysis
• Li+, K+, Na+, Mg2+, Ca2+, Al3+

Question 33

Electrolysis with Overvoltage

Answer 33

Additional voltage that must be applied in order to get non spontaneous reaction to occur

Question 34

Corrosion

Answer 34

Usually gradual; oxidation of metals exposed to oxidizing agents of the environment (simply because oxygen has a strong tendency to undergo reduction)

Question 35

Rusting:

• Anodic regions
• Cathodic regions
• Promoted by?

Answer 35

• Oxidation reaction stands to occur at this site
• Electrons produced at anodic regions travel to cathodic region to react with oxygen and H+ ions dissolved in moisture
• Electrolytes and presence of acids promote rusting

Question 36

Sacrificial electrode

Answer 36

• Composed of a metal that OXIDIZES more easily than iron; oxidizes in place of iron