Chapter 20: Electrochemistry Flashcards Preview

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Flashcards in Chapter 20: Electrochemistry Deck (36)
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1

Oxidation

Loss of electrons

2

Reduction

Gain of electrons

3

Balancing oxidation-reduction equations: Acidic solutions

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

4

Balancing redox reactions: Basic solutions

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

5

Electrical current

Flow of electric charge

6

Electrochemical cell definition

Generation of electricity through redox reactions carried out in this device

7

Voltaic (Galvantic) cell definition

Electrochemical cell that produces electrical current from a spontaneous chemical reaction

8

Electrolytic cell definition

Consumes electrical current to drive non spontaneous reactions

9

Voltaic Cell Components

- 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)

10

Ampere

(A) AKA amps; measurement of electrical current

11

Potential difference and its significance

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

12

Cell potential (Ecell or cell emf)

- 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

13

E˚cell

- Standard cell potential (1M, 1 atm, or 25˚C)
E˚cell = E˚final - E˚initial | E˚cathode - E˚anode

14

Cell diagram/line notation

- 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

15

Standard electrode potential

Electrode in each half-cell's own individual potential

16

Standard hydrogen electrode

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

17

Summary of Standard Electrode Potentials

- 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

18

Calculating Standard Potential

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)

19

Predicting Spontaneous Reactions

- 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

20

Predicting whether metal will dissolve in acid

- 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

21

Faraday's Constant

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

22

Relating E˚cell and K

E˚cell = 0.0592V/n(logK)

23

Cell Potential and Concentration (Nernst)

Cell = E˚cell - (0.0592V/n)(logQ)

24

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

- Ecell = E˚cell
- Ecell > E˚cell
- Ecell < E˚cell

25

Dry cell battery

Fairly common; do not contain large amount of liquid

26

Lead-Acid Storage Batteries

Found in most cars; 6 electrochemical cells wired in series

27

Fuel cells

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

28

Electrolysis

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

29

Summary:
- Electrochemical cells
- Voltaic cells
- Electrolytic cells

- 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

30

Electrolysis of pure molten salt

Anion is oxidized and cation is reduced

31

Electrolysis of Mixtures of Cations or Anions

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

32

Electrolysis of Aqueous Solutions

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

33

Electrolysis with Overvoltage

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

34

Corrosion

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

35

Rusting:
- Anodic regions
- Cathodic regions
- Promoted by?

- 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

36

Sacrificial electrode

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