Final Exam Flashcards
(139 cards)
1
Q
4 types of IMFs
A
Van der Waals
- dipole-dipole
- dispersion
- hydrogen bonds
Other
- ion-dipole
2
Q
dipole-dipole
A
attraction between polar molecules
3
Q
hydrogen bonds
A
H bound to N, O, F
subset of dipole-dipole
H acquires highly positive charge
4
Q
dispersion forces
A
result from instantaneous dipoles & induced dipoles
5
Q
IMF present in all molecules
A
dispersion forces
6
Q
what kinds of molecules have stronger dispersion forces?
A
larger
linear
7
Q
ion-dipole interactions
A
ions interacting with neutral polar molecules
8
Q
determines magnitude of ion-dipole interactions
A
cations are stronger
ions with smaller nuclei are stronger
9
Q
cohesion
A
attraction of molecules to one another
10
Q
adhesion
A
attraction of molecules to another surface
11
Q
1 atm = ____ torr
A
760
12
Q
4 types of crystals
A
ionic
covalent
molecular
metallic
13
Q
properties of ionic crystals
A
hard, brittle, high MP, poor conductor, often white/grainlike
14
Q
properties of covalent crystals
A
hard, brittle, high MP, poor conductor
15
Q
properties of molecular crystals
A
soft, low MP, poor conductors
more likely to be organic
16
Q
properties of metallic crystals
A
strong, malleable, good conductors, variable MPs
17
Q
substance exists in equilibrium of all 3 phases
A
triple point
18
Q
equilibrium between liquid and gas terminates
A
critical point
19
Q
3 phase change curves
A
- Melting curve
- Vapor pressure curve
- Sublimation curve
20
Q
state functions (9)
A
mass
pressure
temperature
volume
particle number
entropy
enthalpy
Gibbs free energy
internal energy
21
Q
C + ___ = K
A
273
22
Q
(+) Gibbs means…
A
non-spontaneous
endergonic
23
Q
(-) Gibbs means…
A
spontaneous
exergonic
24
Q
standard conditions (6)
A
- 1 atm
- 1.00 M
- pH 7 (biological)
- 25° C
- Pure solids & liquids
- Most stable allotropic form of elements at std conditions
25
3 laws of TD
1. Energy cannot be created or destroyed, only _transduced_ (changed)
2. The total entropy of the universe is always increasing
3. The entropy of a perfect crystalline substance is zero at 0° K (absolute zero)
26
S° ↑ with… (5)
molar mass & molecular complexity, temperature, volume, and when a reaction produces more gas molecules than it consumes
27
overall reaction order =
sum of exponents in the rate law
28
order of a reaction depends on…
concentration of reactants only
29
elementary reaction
single collision of reactant molecules
30
intermediates
species that appear in the mechanism, but not in the products or reactants
31
molecularity
number of reactant molecules involved in the collision
32
rate-determining step
slowest step in the sequence
33
Rate-determining step must have…
the same rate law as that determined by experimental data for the overall reaction
34
heterogeneous catalyst
reactants & catalysts are in different phases
35
homogeneous catalyst
reactants & catalysts are in the same phase - usually in solution
36
If Kc \>\> 1, then equilibrium…
lies to the right
𝚫G° is negative
37
If Kc \<\< 1, then equilibrium…
lies to the left
𝚫G° is positive
38
If Kc is between ___ and ___ , product concentration roughly equals reactants’
0.01 and 100
39
Law of mass action
Kc = Qc at equilibrium
40
If Qc \> Kc…
reactants favored
41
If Kc \> Qc…
products favored
42
In the Kp equation, 𝚫n =
moles gaseous product - moles gaseous reactant
43
If you reverse the reaction, Kc…
becomes 1/Kc = Kc\*
44
If you multiply the equation, Kc…
is raised to the power of what you multiplied by
45
if you divide the reaction, Kc…
the root of the divisor is taken on Kc
46
if you add 2 reactions, the Kc's…
are multiplied together
47
* Increase [reactants] =
* Increase [products] =
* Decrease [reactants] =
* Decrease [products] =
* Increase [reactants] = Q \< K, shift to right
* Increase [products] = Q \> K, shift to left
* Decrease [reactants] = Q \> K, shift to left
* Decrease [products] = Q \< K, shift to right
48
* Decrease volume =
* Increase volume =
* Decrease pressure =
* Increase pressure =
* 𝚫 in pressure or volume =
* Adding an inert gas to the reaction at fixed volume & pressure =
* Decrease volume = shift to side with fewer particles
* Increase volume = shift to side with more particles
* Decrease pressure = shift to side with more particles
* Increase pressure = shift to side with fewer particles
* 𝚫 in pressure or volume = no change if both sides have equal particle number
* Adding an inert gas to the reaction at fixed volume & pressure = no change
49
* Increase heat in an exothermic rxn =
* Decrease heat in an endothermic rxn =
* Increase heat in endothermic rxn =
* Decrease heat in endothermic rxn =
* Increase heat in an exothermic rxn = shift to left
* Decrease heat in an endothermic rxn = shift to right
* Increase heat in endothermic rxn = shift to right
* Decrease heat in endothermic rxn = shift to left
50
amphoteric
acts as acid or base
51
Kw =
Kw = [H3O][OH] = 10-14 @ 25° C
52
Hydrohalic acid strength
HF \<\< HCl \< HBr \< HI
53
oxoacids
contain H, O and a central nonmetal atom
54
examples of oxoacids
Chloric acid HClO3; bromic acid HBrO3; carbonic acid H2CO3; nitrous acid HNO2; nitric acid HNO3; phosphorous acid H3PO3; phosphoric acid H3PO4; sulfuric acid H2SO4
55
more oxygen in an oxoacid =
stronger acid
56
more electronegative central atoms in an oxoacid =
stronger acid
57
carboxylic acids
include C bound to 2 O and an R group
58
examples of strong acids
* Hydrochloric acid HCl
* Hydrobromic acid HBr
* Hydroiodic acid HI
* Nitric acid HNO3
* Chloric acid HClO3
* Perchloric acid HClO4
* Sulfuric acid H2SO4
59
examples of strong bases
* Group 1A hydroxides
* LiOH
* NaOH
* KOH
* RbOH
* CsOH
* Group 2A hydroxides
* Ca(OH)2
* Sr(OH)2
* Ba(OH)2
60
Percent ionization =
100([H3O]eq / [HA]0)
61
larger Ka =
stronger acid
62
lewis acid
accepts pair of electrons
63
lewis base
donates pair of electrons
64
when can the H-H equation no longer be used?
at the equivalence point and beyond
65
to find a suitable buffer…
find an acid w/ pKa within a unit of the pH
66
“acid + base → …….”
salt + water
67
not included in Ksp equation
solids
68
molar solubility:
mass solubility:
mol/L (M)
g/L
69
what does common ion effect have on the ICE table?
changes the initial value of the ion present in the “other” solution
the spectator ion doesn't matter
70
oxidation
loss of e-
71
reduction
gain of e-
72
charge of H with a nonmetal
positive
73
charge of H with a metal
negative
74
oxidation state rule about oxygen
usually 2-, except in molecules with O-O bonds
75
elements in bottom left corner more likely to have ___ charge
positive
76
11 oxidation state rules
1. 0 charge in substance with atoms of one element
2. simple ions' charge is their OS
3. H+ with a nonmetal
4. H- with a metal
5. Group 1A = 1+
6. Group 2A = 2+
7. Oxygen usually 2-, except in O-O molecules
8. Group VIIA = 1-
9. sum of OS in neutral compound is 0
10. sum of states in polyatomic ion is that ion's charge
11. elements in bottom left more likely to be +
77
10 steps to balance redox rxns
1. split into ½ reactions
2. balance all atoms except O and H
3. balance O with H2O
4. balance H with H+
5. balance charge with e-
6. (bases) balance H+ with OH- on BOTH sides
7. (bases) combine H+ with OH- to form water
8. cancel all e- via multiplication
9. combine equations - cancel anything you can
10. check charges & balancing
78
cell notation
(anode reactant) | (anode product) || (cathode reactant) | (cathode product)
79
cell potential
Ecell
80
E =
| (redox)
electron motive force
81
cathode
reductive process
gain of electrons
becoming more negative
82
anode
oxidative process
loss of electrons
becoming more positive
83
rule about Ecell under std conditions
must be +
84
E of H2
0 V
85
strongest reducing agents are…
strongest oxidizing agents are…
most negative anodes
most positive cathodes
86
n (electrochem) =
moles of electrons transferred
87
3-way relationship in electrochem
E° — 𝚫G° — K
88
if we are setting up a galvanic cell, we want E° to be…
positive (spontaneous)
89
used for non-std conditions in galvanic cells
Nernst equations
90
2 types of Nernst equations
2 for non-std temp (involving RT)
1 for std temp (omits RT)
91
to find variables for Nernst equation…
find ½ reactions
use E° values given for each
balance to find n (electrons transferred)
use [products]/[reactants] to find Q (omitting water)
92
if Q \> 1, E __ E°
E \< E°
93
if Q \< 1, E __ E°
E \> E°
94
radioactivity
all ___ elements
spontaneous emission of particles or electromagnetic energy
elements \>83
95
proton symbol
1/1 P
1/1 H
96
neutron symbol
1/0 n
97
electron symbol
0/-1 e
0/-1 B
98
positron symbol
0/+1 e
99
alpha particle symbol
4/2 a
| (He)
100
top number in notation for subatomic particles
mass #
101
bottom number in notation for subatomic particles
atomic #
102
“belt of stability”
stable nuclei - 2, 8, 20, 50, 82, 126 protons/neutrons
103
4 types of nuclear decay
beta emission
electron capture
positron emission
alpha decay (releases He)
104
gamma radiation
emits photons
105
notation for nuclear transmutation
14/7 N + 4/2 H → 17/8 O + 1/1 H
**14/7N (a, p) 17/8O**
parent nucleus (bombarding particle, emitted particle) product nucleus
106
higher E; more likely to be \_\_\_\_\_\_
lower E; more likely to be \_\_\_\_\_\_\_
reduced
oxidized
107
E of strongest reducing agents
more negative
108
E of strongest oxidizing agents
more positive
109
“what is the name of a thin film of water contracting and adhering to the wall of a glass cylinder?”
capillary action
110
“which of the following has the highest surface tension at a given temp?”
two molecules with O by double bonds
111
“location at which two phases can exist in equilibrium”
phase boundary
112
what forces hold each type of solid together?
ionic crystals: attractions between anions and cations
metallic crystals: IMFs & sharing of electrons throughout
covalent crystals: covalent bonds
molecular crystals: IMFs
113
O – C – O
IMFs
dispersion
114
3H - C —— C - 2H - OH
IMFs
dispersion
hydrogen bonds
dipole-dipole
115
H —— S — 3Cl
IMFs
dispersion
dipole-dipole
116
“highest boiling point”
NH3 or CH4
NH3
117
“highest vapor pressure”
smaller molecule - C3H8
118
“highest surface tension”
CH3COOH
OH group present
119
“order of reaction of decomposition of sulfuryl chloride (SO2Cl2)”
first (see units)
120
“for reaction A + 2B → C, which expression is correct?”
𝚫[C]/𝚫t = -½𝚫B/𝚫t
121
steps to find rate law given a table of experimental data
take initial concentration of one trial divided by another (where other reactant is constant concentration)
take same given initial rates divided by one another
**the rate quotient = the concentration quotient*****x***
repeat for other reactant(s)
122
equation to use for finding concentration of dimerizing butadiene after 3 hours
integrated rate law for second order rxns
(0.0052 M)
123
number of elementary steps =
number of “hills” on the potential energy profile
124
“what will decrease the entropy of an inert gas system?"
decreasing volume
125
entropy increases with…
more particles
more energy
higher temperature
greater molecular mass
more even distribution of particles among cells
126
what is larger: joule or kilojoule
kilojoule
127
joule to kilojoule
divide by 1000
128
kilojoule to joule
multiply by 1000
129
conjugate base of sulfuric acid?
HSO4 -
130
“Kc of decomposition of HI”
0.26
131
“concentration of N2 for nitrogen fixation reaction”
2.7 M
132
“[H3O] concentration in formic acid”
4.2 x 10^-3
133
“Kp of nitric oxide and bromine”
134
134
"Ka of butyric acid"
1.5 x 10^-5
135
question to go straight to QF after ICE table
[H3O]
piric acid in the leather industry
0.15 M
136
“cartoon of weak acid at equilibrium”
“solvent water molecules omitted for clarity”
Ka = 10^-3
137
"polyprotic acid malonate ion [] "
2 x 10^ -6 M
138
pH of malonic acid problem
1.8 (-log0.016)
139
difference between the mass of an atom and the sum of the masses of its constituent nucleons
mass defect
