Definitely on the Test Flashcards
1
Q
le chatelier’s principle
A
if a reaction in equilibrium is disturbed, the reaction shifts left in order to minimize the disturbance
2
Q
acid equilibrium expression
A
K(a) = [H3O+][A-]/[HA]
3
Q
rate and temperature
A
higher temp. = higher rxn rate because there will be more collisions (need correct orientation)
4
Q
units for rate laws (1st order)
A
s ^-1
5
Q
spontaneity
A
the tendency to occur without being driven by an external force
6
Q
entropy
A
energy and matter tend to disperse in an orderly manner
7
Q
predicting change in entropy
A
- phase change (s - l +S; l - q +S, q - l -S)
- state of matter (solid < liquid < gas)
- atomic weight (He < Ne < Ar < Kr, Xe)
8
Q
enthalpy & ∆S
A
if ∆H is - (exothermic) ∆S increases
if ∆H is + (endothermic) ∆S decreases
9
Q
∆S equation
A
∆S = ∆S(final) — ∆S(inital) [J/K or J/Kmol]
10
Q
direction of chemical change
A
entropy helps us determine the direction a rxn will proceed in (direction that increases entropy of the universe)
11
Q
second law of thermodynamics
A
for any spontaneous process, the entropy of the universe increases (∆S > 0)
12
Q
∆S and ∆H
A
entropy is more important than enthalpy at certain temps
13
Q
equation for entropy of the universe
A
∆S(universe) = ∆S(system) + ∆S(surroundings)
14
Q
temperature effect on ∆S
A
entropy depends on temperature
magnitude of ∆S(surr) is proportional to ∆H(sys)
15
Q
equation for temperature effect on ∆S
A
∆S(surr) = -∆Hsys/T
16
Q
equation for gibbs free energy
A
∆G(sys) = ∆H(sys) - T∆S(sys)
17
Q
gibbs free energy
A
∆G is a value that can tell us about the spontaneity of a reaction
18
Q
gibbs and spontaneity
A
if ∆G is + , ∆S(uni) is — (nonspontaneous)
if ∆G is - , ∆S(unis) is + (spontaneous)
19
Q
STP
A
25° C, 1 atm, 1.0 M
20
Q
∆S° equation
A
∆S°(rxn) = ∑n∆S(products) — ∑n∆S(reactants)
21
Q
∆G° equation
A
∆G°(rxn) = ∑n∆G(products) — ∑n∆G(reactants)
22
Q
entropy for an element
A
∆S° is NOT 0!
23
Q
calculating ∆G at non-standard conditions
A
∆G = ∆G° + RTlnQ
24
Q
calculating ∆G at equilibrium
A
∆G° = -RTlnK
∆G° = K @ equilibrium
25
intermolecular forces (ion-ion)
essentially an ionic bond (not an IMF) [opposites attract]
26
intermolecular forces (ion-dipole)
charged ions are attracted to the opposite "partial charge" on a nearby polar molecule [STRONGEST]
27
intermolecular forces (hydrogen)
dipole-dipole interaction involving "H" atoms. [2nd STRONGEST]
rule to form hydrogen IMF:
1. H must be bonded to F, O, or N
2. H is attracted to a lone electron pair
28
intermolecular forces (dipole-dipole
partial positive & negative charges on polar molecules attract [3rd STRONGEST]
29
intermolecular forces (london dispersion)
induced dipole moment. depends on the molecular weight of the molecule. More weight = stronger LD force. (all molecules have LD forces) [4th STRONGEST]
30
how a liquid boils
it undergoes a phase change (all imf forces must be broken)
31
phase change (L - G)
vaporization (boiling)
32
phase change (L - S)
freezing
33
phase change (S - G)
sublimation
34
phase change (G - L
condensation
35
phase change (S - L)
melting
36
phase change (G - S)
deposition
37
vapor pressure and temperature
as temp. increases, so does vapor pressure
38
boiling point
the temp. where vapor pressure of has is equal to atmospheric pressure
39
equilibrium
the rate of the forward reaction is equal to the rate of the reverse reaction
(concentrations of reactants/products stay constant)
40
equilibrium constant (K)
K = [products]/[reactants]
41
equilibrium constant (K) extra info
K depends on temperature and nature of reactants
pure liquids and solids are not involved in the expression
K is unitless
42
reverse of a chemical rxn
K(reverse) = 1/K
43
multiplying rxn by a factor
K^factor
44
adding two rxns
K(overall) = K(1) * K(2)
45
K and pressure
K is large (> 1) [Pro] > [Reac] rxn lies to the right [ΔG -]
K is small (< 1) [Pro] < [Reac] rxn lies to the left [ΔG +]
46
factors that change equilibrium (volume)
V increases (P decreases) rxn shifts left
V decreases (P increases) rxn shift right
47
factors that change equilibrium (temperature) [exothermic]
[adding product] rxn shifts left
[removing product] rxn shifts right
48
factors that change equilibrium (temperature) [endothermic]
[adding reactant] rxn shifts right
[removing reactant] rxn shifts left
49
strong acid
an acid that dissociates completely into ions solution
50
strong base
a base that reacts completely with H+ ions in solution.
51
weak acid
an acid that is only slightly dissociated in solution
52
weak base
A base that does not completely react with H+ ions in solution.
53
base equilibrium expression
K(b) = [OH-][HB+]/[B]
54
large K(a)
acid is relatively strong weak acid (weak conjugate base)
55
small K(a)
acid is relatively weak weak acid (strong conjugate base)
56
K(w) equation
K(w) = K(a) * K(b)
57
auto ionization of water
K(w) = [H3O+][OH-] = 1.0x10^-14
58
template for weak acid rxn
HA + H2O ⇄ H3O+ + A-
59
template for weak base rxn
B + H2O ⇄ OH- + HB+
60
K(a) and pK(a)
pK(a) = -log[Ka]
K(a) = 10^-pKa
61
pH and pOH
pH = -log(H+) [H3O+]
pOH = -log(OH-)
14 = pH + pOH
62
pH scale
ph > 7 basic
pH = neutral
ph < 7 acidic
63
list of strong acids
HCl, HBr, HI,
H2SO4
HClO4
HNO3
64
list of strong bases
LiOH, NaOH, KOH,
Ca(OH)2, Sr(OH)2, Ba(OH)2
65
buffer solution
a buffer is a solution that resists a change in pH when a strong acid/base is added, or if the solution is diluted
66
henderson - hasselbalch equation
easy way to calculate pH of buffer
pH = pKa + log [A-]/[HA]
Ka = Kw/Kb
67
titrations
a controlled chemical rxn that is used to calculate the concentration of an unknown substance
68
oxidation
the loss of electrons (OIL)
69
reduction
the gain of electrons (RIG)
70
rules for redox reactions
both oxidation and reduction must occur
of e- lost must = # of e- gained
71
rules for balancing redox reactions
1. break the reaction into two half rxns
2. balance all atoms other than O and H
3. add H2O to balance O atoms
4. add H+ to balance H atoms
a. if in acid go to step 5
b. if in base, add OH- to both sides
5. add e- to correct side to balance
6. multiply half rxns to get e- equal
7. add the two half rxns
72
galvanic (voltaic) cell
redox rxn that is always spontaneous (E +) [G-]
73
salt bridge
1. each rxn is separated into a half cell
2. e- transfer occurs at electrodes (solids)
3. ions move through solution
4. oxidation occurs at the anode
5. reduction occurs at the cathode
6. salt bridge connects the two cells and allows charge to flow between solutions
74
direction of ions/e- in salt bridge (galvanic)
e- flow FROM anode TO cathode
NO3- flows into anode solution
K+ flows into cathode solution
anode solution goes into metal
cathode metal goes to cathode solution
75
calculating e cell
E°cell = E°red — E°ox (in V)
76
e° cell at equilibrium
E°cell = (0.0592/n) * logK
77
nernst equation
used when not at equilibrium
E = E° — (0.0592/2) * logQ
78
what does nernst equation mean?
Q > 1 (Ecell decreases) [rxn shifts left]
Q = 1 (ΔE = 0V)
Q < 1 (Ecell increases) [rxn shifts right]
79
electrolytic cell
a nonspontaneous electric cell which requires a potential greater than E°cell to drive the rxn
80
anode + cathode (electrolytic)
anode is +
cathode is —
[reduction still happens at cathode]
81
electrical current
current is the flow of charge
82
current (i)
i = charge / time
83
ampere
coulomb / second
84
what factors affect rxn rates (physical state of matter)
gas = slow
liquid = faster
solid = fastest
85
what factors affect rxn rates (conc. of reactants)
higher [conc.] = fast
lower [conc.] = slow
86
what factors affect rxn rates (temp.)
higher temp. = fast
lower temp. = slow
87
what factors affect rxn rates (presence of catalyst)
catalysts speed up the reaction
88
rxn rate
R = Δ[conc] / Δt
89
rxn rate of reactant
R = — Δ[reactants] / Δt
90
rxn rate of product
R = + Δ[products] / Δt
91
general rate of rxn
aA + bB = cC
R = (-1/a)*Δ[A]/Δt = (-1/b)*Δ[B]/Δt = +(1/c)*Δ[C]/Δt
92
rate laws
rate = k[A]^m [B]^n
93
what rate constant (k) means
large k = FAST
small k = SLOW
94
activation energy
the minimum amount of energy required to start a chemical reaction (break bonds)
95
Ea (exothermic)
final potential energy is lower than starting energy
96
Ea (endothermic)
final potential energy is higher than starting energy
97
arrhenius equation
k=Ae^(-Ea/RT)
98
single step rxn mechanism
reaction occurs in a single step
2 molecules = bimolecular
3 molecules = termolecular
4 molecules = quad-molecular
99
multi-step rxn mechanism
rxn happens in multiple steps
(will have an intermediate species that doesn't appear in the overall rxn)
100
orders of multi-step rxns
A = product (1st order)
A + B = product (2nd order)
A + A = product (2nd order)
A + A + B = product (3rd order)
101
catalyst
substance that speeds up the rate of a chemical reaction by lowering the activation energy (is not used up in the reaction) [will appear in the beginning and end of reaction but not in the overall]
102
integrated rate law (0th order)
[A]t — [A]0 = -kt
103
integrated rate law (1st order)
ln[A]t — ln[A]0 = -kt
104
integrated rate law (2nd order)
1/[A]t — 1/[A]0 = kt
105
units for rate laws (0th order)
M / s
106
units for rate laws (2nd order)
1 / M * s
107
units for rate laws (1st order)
s^-1
