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Chemistry for Engineers Finals Notes > GE-CHEM 1103 Module 4.3 > Flashcards

GE-CHEM 1103 Module 4.3 Flashcards

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1
Q

the study of the
relationships between electricity
and chemical reactions. It includes
the study of both spontaneous and
nonspontaneous processes.

A

Electrochemistry

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2
Q

To keep track of what
loses electrons and
what gains them, we
assign

A

oxidation
numbers

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3
Q

If the oxidation number
increases for an
element

A

element
is oxidized

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4
Q

If the oxidation number
decreases for an
element

A

element
is reduced

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5
Q

Synopsis of Assigning Oxidation Numbers

A
  • Elements = 0
  • Monatomic ion = charge
  • F: –1
  • O: –2 (unless peroxide = –1)
  • H: +1 (unless a metal hydride = –1)
  • The sum of the oxidation numbers equals
    the overall charge (0 in a compound).
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6
Q

LEO

A

Loses Electron Oxidized

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7
Q

GER

A

Gains Electron Reduced

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8
Q

causes something else to be
oxidized

A

oxidizing agent

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9
Q

causes something else
to be reduced

A

reducing agent

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10
Q
  • The oxidation and reduction are written
    and balanced separately.
  • We will use them to balance a redox
    reaction.
A

Half-Reactions

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11
Q

Balancing Redox Equations: The
Half-Reactions Method (For acidic solution)

A
  1. Make two half-reactions (oxidation
    and reduction).
  2. Balance atoms other than O and H.
    Then, balance O and H using H2O/H+.
  3. Add electrons to balance charges.
  4. Multiply by common factor
    to make electrons in half-reactions
    equal.
  5. Add the half-reactions.
  6. Simplify by dividing by common factor
    or converting H+
    to OH–
    if basic.
  7. Double-check atoms and charges
    balance!
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12
Q

Balancing in Basic Solution

A
  • A reaction that occurs in basic solution can
    be balanced as if it occurred in acid.
  • Once the equation is balanced, add OH–
    to
    each side to “neutralize” the H+
    in the
    equation and create water in its place.
  • If this produces water on both sides,
    subtract water from each side so it appears
    on only one side of the equation.
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13
Q
  • In spontaneous redox
    reactions, electrons
    are transferred and
    energy is released.
  • That energy can do
    work if the electrons
    flow through an
    external device.
A

Voltaic Cells

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14
Q

oxidation occurs

A

anode

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15
Q

reduction occurs

A

cathode

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16
Q

When electrons flow, charges aren’t balanced so a contraption is used to keep the charges balanced

A

a salt bridge, usually a U-shaped tube that contains a salt/agar solution

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17
Q

electrons leave
the anode and flow through
the wire to the cathode

A

cell

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18
Q

formed in the
anode compartment

A

Cations

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19
Q

electrons reach the
cathode

A

cations in solution
are attracted to the now
negative cathode

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20
Q

gain electrons
and are deposited as metal
on the cathode

A

cations

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21
Q

flow
spontaneously
one way in a
redox reaction,
from high to low
potential energy.

A

electrons

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22
Q

potential difference between the anode
and cathode in a cell

A

electromotive force (emf)

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23
Q

Other name of electromotive force (emf)

A

cell potential and is
designated Ecell

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24
Q

It is measured in volts (V). One volt is one
joule per coulomb (1 V = 1 J/C).

A

cell potential

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25
Q

Standard condition of a reduction potentials

A

1 M, 101.3 kPa,
25°C

26
Q

reference for the reduction potentials is
called

A

standard
hydrogen electrode
(SHE)

27
Q

the reduction potential
for hydrogen

A

0 V

28
Q

The cell potential at standard conditions can be found through this equation:

A

E°cell = E°red(cathode) - E°red(anode)

29
Q

cell potential is based on the potential
energy per unit of charge

A

intensive property

30
Q

The more positive the
value of E° red

A

the greater the tendency
for reduction

31
Q

most positive
reduction potentials

A

The strongest
oxidizers

32
Q

most negative
reduction potentials

A

The strongest
reducers

33
Q

produce a positive
cell potential, or emf.

A

Spontaneous redox reactions

34
Q

Equation for spontaneous redox reactions

A

E° = E° red (reduction) – E° red (oxidation)

35
Q

positive emf corresponds to
negative ΔG

A

Since Gibbs free energy is the measure of
spontaneity

36
Q

How is Free Energy, Redox, and K related

A

ΔG° = –nFE° = –RT ln K

37
Q

Nernst equation

A

E = E° – (RT/nF) ln Q
OR E = E° – (2.303 RT/nF) log Q

38
Q

Nernst equation (Using standard thermodynamic temperature and the constants R and F)

A

E = E° – (0.0592/n) log Q

39
Q

a cell could be created that has the same substance at both electrodes

A

concentration cell

40
Q

as long as the concentrations are
different

A

E will not be 0

41
Q

Some Applications of Cells

A

– Batteries
– Prevention of corrosion
– Electrolysis

42
Q

Examples of batteries

A
  • Lead–acid battery
  • Alkaline battery
  • Ni–Cd and Ni–metal hydride batteries
  • Lithium-ion batteries
43
Q

reactants and products are
solids, so Q is 1 and the potential is independent of concentrations; however, made with lead and sulfuric acid (hazards)

A
  • Lead–acid battery
44
Q

most common primary battery

A
  • Alkaline battery
45
Q

lightweight, rechargeable; Cd is toxic and heavy, so hydrides are replacing it

A
  • Ni–Cd and Ni–metal hydride batteries
46
Q

rechargeable, light; produce
more voltage than Ni-based batteries

A
  • Lithium-ion batteries
47
Q

the energy created can be
converted to electrical energy

A

fuel is burned ( this conversion is only 40% efficient, with the remainder lost as heat)

48
Q

The direct conversion of chemical to electrical energy is expected to be more efficient and is the basis for it

A

fuel cells

49
Q

are NOT batteries; the source of energy
must be continuously provided

A

fuel cells

50
Q
  • In this cell, hydrogen
    and oxygen form water.
  • The cells are twice as
    efficient as combustion.
  • The cells use hydrogen
    gas as the fuel and
    oxygen from the air.
A

Hydrogen Fuel Cells

51
Q

it is an oxidation and its common name is rusting

A

Corrosion

52
Q

occurs because zinc is more easily
oxidized, so that metal is sacrificed to keep the iron from rusting

A

Cathodic protection

53
Q

A method to prevent corrosion is used for
underground pipes wherein the anode is oxidized before the pipe

A

sacrificial anode

54
Q

Use of electrical energy to create chemical
reactions

A

electrolysis

55
Q

1 coulomb

A

1 ampere × 1 second

56
Q

equation for Q (relation to electrolysis and stoichiometry)

A

= It = nF

57
Q

Q represents

A

charge (C)

58
Q

I represents

A

current (A)

59
Q

t represents

A

time (s)

60
Q

n represents

A

moles of electrons that
travel through the wire in
the given time

NOTE: n is different than that
for the Nernst equation!

61
Q

F represents

A

= Faraday’s constant (96,485 C/mol)

Chemistry for Engineers Finals Notes (6 decks)

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