Chem 112B Final Flashcards
1
Q
Collision theory
A
The rate of a chemical reaction is proportional to the number of collisions between reactant molecules.
The more often reactant molecules collide, the more often they react with one another, and the faster the reaction rate
2
Q
In a reaction profile, peaks are… and valleys are…
A
transition states, intermediate states
3
Q
Reaction energy (ΔE) is
A
the energy of products minus the energy of the rectants
4
Q
Activation energy (Ea) is
A
the energy needed to get to the peak, can be either TS-reactant or TS-intermediate
5
Q
In an elementary reaction, the method of initial rates
Rate = k[A]^x{B}^y
A
x and y are the same as a and b
6
Q
In a multi-step reaction, the exponents
A
must be determined experimentally
7
Q
Reaction mechanism consists of
A
elementary steps
8
Q
Rate law is always based on the
A
slowest step
9
Q
How do I know if a reaction is elementary or not?
A
All reaction steps in a reaction mechanism are elementary
10
Q
If the first step is the fast step, will it always be an equilibrium reaction?
A
Yes. For fast first step, use equilibrium to solve for concentration of intermediate
11
Q
1st order integrated rate law
A
plotting ln[A] vs time gives a straight line. Half-life only depends on k
12
Q
2nd order integrated rate law
A
plotting 1/[A] vs time gives a straight line. Half life depends on both k and [A]0.
13
Q
Plotting lnk vs 1/T gives a
A
straight line
slope = -Ea/R
y-intercept = lnA
14
Q
At two different temperatures, we can determine Ea or ratio of k using which two equations
A
lnk = -Ea/RT + lnA
lnk1/k2 = Ea/R [1/T2 - 1/T1]
15
Q
At a fixed temp, we can use difference in Ea to determine..
A
ratio of k
kcat/kuncat= e (,..)
16
Q
H-bond donor
A
strongly electronegative atom such as N, O, or F
17
Q
H-bond acceptor
A
electronegative atom of a neighboring molecule or ion that contains a lone pair
18
Q
To be phosphorylated, a side chain must
A
contain alcohol or amine group
19
Q
if the half-life is constant, the order is
A
1st order
20
Q
Intermediate
A
Produced, then consumed
21
Q
Catalyst
A
Consumed, then produced
22
Q
Rate of forward process=
A
Rate of reverse process
23
Q
For elementary reactions, Keq=
A
kf/kr, where kf are rate constants for the forward and reverse reaction, respectively
24
Q
K»_space; 1
A
Products predominate
25
K << 1
Reactants predominate
26
Keq for the reverse reaction is the
inverse of Keq for the forward reaction
27
When a reaction has been multiplied by a number, Keq for the new reaction is the
original Keq raised to a power of the multiplication factor
28
If the net reaction can be obtained by the addition of 2 or more steps, Keq is the
product of the individual Keq
29
Q
reaction quotient
30
Q is calculated using...Keq is calculated using...
Any concentrations, equilibrium concentrations
31
Q< Keq
Reaction shifts right (produce more products, consume more reactants)
32
Q > Keq
Reaction shifts left (produce more reactants, consume more products)
33
At equilibrium, added reactants or products will cause the reaction to shift to
consume towards to reestablish equilibrium (balance scale)
34
Add reactants or remove products
Reaction shifts right
35
Remove reactants or add products
Reaction shifts left
36
Since pressure is increased by disturbance, reaction will shift in direction that
reduces pressure
37
Decreasing the volume shifts the reaction to the side with
less moles of gas
38
If a reaction has the same moles of gas on both sides, changing pressure/volume does
NOT have any effect on equilibrium
39
When heat is added to a system at equilibrium, reaction shifts in the direction that
absorbs heat
40
Spontaneity
process proceeds in forward direction to an appreciable extent by itself
41
Free energy of formation at standard conditions can be calculated using
ΔG°f = ΔH°f - (298K)ΔS°
42
Calculate ΔG° at 25°
ΔG° = ΔH° - (298K)ΔS°
43
Calculate ΔG° at temps besides 25°
ΔG° = ΔH° - TΔS°
ΔG° = -RTlnK
44
ΔG > 0, Q > K
System is nonspontaneous in forward direction
45
ΔG < 0, Q < K
System is spontaneous in forward direction (system does work)
46
Endothermic ΔH
positive
47
Exothermic ΔH
negative
48
Molar entropy increases when
Gaseous, largest molar mass, most molecular complexity, volume increases, pressure decreases
49
Increased Kd means
Weaker binding
50
Decrease in entropy of a reaction
Decrease is numbers of gaseous or aqueous molecules
51
Increase in entropy of a reaction
Solids dissolve
Endothermic (s->l->g)
Volume increase
52
7 strong acids
HClO4, HI, HBr, HCl, H2SO4, HNO3, HClO3
53
8 strong bases
KOH, LiOH, NaOH, Ca(OH)2, Ba(OH)2, RbOH, Sr(OH)2, CsOH
54
Group 1A and heavy group 2A hydroxides are
strong bases
55
Amines are
weak bases
56
Carboxylic acids and thiols are
weak acids
57
The conjugate base of a strong acid is
a negligible base
58
The conjugate base of a weak acid is a
weak base
59
For weak acids and bases, the stronger the acid, the
weaker its conjugate base
60
Within the same group,
larger atoms means stronger acid
61
Within the same period, higher electronegativity means
stronger acids
62
The more electrognegative the central atom, he higher the
number of oxygen atoms attached, the stronger the acid
63
Non-metal oxides/hydroxides are
acidic
64
For pure water,
[OH] = [H]
65
Strong acids and strong bases are strong electrolytes, so they
dissociate completely in water
66
Larger Ka, smaller pKa
Stronger acid
67
Smaller Kb, larger pKb
Weak base
68
Addition of a common ion decreases
ionization
69
Buffer
solution that contains a weak acid and salt of its conjugate base
solution that contains a weak base and salt of its conjugate acid
70
To make/produce buffer
Combine weak acid and salt of conjugate base/ vice versa
OR use neutralization to combine excess weak acid with a strong base
71
Optimal pH of a buffer =
pKa
72
More concentrated buffer solution has
greater capacity
73
Equivalence point is when
stoichiometrically equivalent quantities of acid and bASE REACTED
74
At half eq point, pH=
pKa
75
Solubility
Maximum amount of solid that will dissolve under a given condition, forms a saturated solution
76
If Q > Ksp
precipitate will form
77
If Q < Ksp
More solid can be dissolved
78
Chelating ligands have
two or more donor atoms
79
Larger Kf (more stable) is mostly due to the
entropy effect
80
Electrode of a voltaic cell
GRedCat, LAnOx
81
Salt bridge of voltaic cell
to neutralize excess charges (anions to anode, cations to cathode)
82
External circuit of a voltaic cell
e- flow from anode to cathode
83
Oxidizing agent
Gains e-, gets reduced
84
Reduction agent
Loses e-, gets oxidized
85
The more positive E°red
the more spontaneous and one will undergo reaction (at cathode)
86
Determine which electrode is cathode and anode
* Write reduction half rxns,
* Compare E° red , the more positive (spontaneous) one will undergo reduction (at cathode)
87
Calculate standard cell potential
* Write oxidation and reduction half rxns,
* E° cell = E° ox + E° red
88
Connect free energy (and work) to cell potential
* ∆G = –nFEcell and ∆G o = –nFE°cell
* Ecell is (+), ∆G is (–), rxn/process is spontaneous
89
Describe electrolysis and electrolytic cell (Drive nonspontaneous rxns by
applying electrical energy.)
* Consist electrodes in either aqueous solutions or molten salts.
* Oxidation still at anode, reduction still at cathode.
* Be able to differentiate voltaic cell and electrolytic cell.
90
Predict products at cathode or anode of electrolytic cell
* Identify all species in solution
* Write half reactions with species as reactants
* Compare potentials to predict products
91
Do quantitative calculations involving electrolysis
* Write balanced half reactions
* Use datasheet conversions for amps, volts, Faradays, etc.
92
Linear combinations of wavefunctions to create molecular orbitals
* Same number of MOs created: bonding and antibonding
* Each MO holds maximum of 2 electrons
93
Bonding orbitals (lower energy)
* Sigma (σ) : e density on axis
* Pi (π): e density above and below
* Subscript of original AO
94
Anti-bonding orbitals (higher energy):
* Node between nuclei
* Star superscript
95
Fill MOs of diatomic molecules/ions:
* Fill lower MOs first.
* For degenerate MOs, singly occupy first before doubly up.
96
Predict bond order using MO diagram:
* BO = ½(#BE - #AE)
* Bond order of zero means substance is very unstable (does not exist)
97
Coordination compound is a
neutral compound of the metal complex ion + counter
ion(s)
98
Inside the bracket of the coordination compound is the
metal complex (ion),
which has a transition metal cation + ligands covalently bond (coordinated) to it.
99
Outside the bracket is the
counter ion(s)
100
The charge of the metal complex =
– (total charge of counter ions)
101
In water, coordination compound fully dissociates into
the metal complex and counter ions
102
The coordination number of the metal is
the number of donor atoms (from the ligand) that are bonded/coordinated to the metal cation
103
Describe metal complex
* Transition metal act as Lewis acids (accept e – pair);
* Ligands act as Lewis bases (donate e – pair)
104
Crystal field theory
Δ = Crystal field splitting energy
Large Δ ->Strong Field -> Low spin
Small Δ -> Weak Field -> High spin
105
Paramagnetic
have unpaired electrons
106
Diaamagnetic
No unpaired electrons
107
Units of 0 order
Ms^-1
108
Units of 1st order
s^-1
109
Units of 2nd order
M^-1s^-1
110
4 factors that affect reaction rates
Inc contact area = faster
Inc concentration = faster
Inc temp = faster
Catalyst lowers Ea = rate faster
111
Small Km,
strong binding
112
Inc kCat, small Km,
Efficient enzyme
113
Q < Keq
Reaction proceeds reverse, more reactant produced
114
Q > Keq
Reaction proceeds forward, more product produced
115
Endothermic heat is the
reactantE
116
Exothermic heat is the
product
117
Kb =
Kon/Koff
118
Kd =
Koff/Kon
119
Large Koff=
weak binding
120
Large Kb=
Strong binding
121
Small Kd =
Strong binding
122
nm to um
divide by 1000
123
Ki =
Koff/Kon
124
Competitive inhibition
Vmax not affected
Km increases
125
Noncompetitive inhibition
Vmax decreases
Km stays same
126
Fewer H-bonds raises
Ki
127
Weaker interactions, Ki
increases
128
For a process to be spontaneous, ΔS must be
+
129
When ΔH and ΔS is negative,
ΔG is negative at low T (spontaneous)
ΔG is positive at high T (nonspontaneous)
130
When ΔH and ΔS is positive,
ΔG is negative at high T (spontaneous)
ΔG is positive at low T (nonspontaneous)
131
When ΔH is negative and ΔS is positive,
ΔG is negative (spontaneous)
132
When ΔH is positive and ΔS is negative,
ΔG is positive (nonspontaneous)
133
Arrhenius acids produce
H+ when dissolved in water
134
Arrhenius bases produce
OH- when dissolved in water
135
Bronsted acids
donate H+
136
Bronsted bases
accept H+
137
Ka increases
acidity increases
138
Stability of conjugate base factors
larger atom
electronegative
resonance
inductive effect (electronegativity)
139
Lewis acids
accept e- (incomplete octets)
140
Lewis bases
donate e- (at least 1 lone pair)
141
Ligand is the molecule or ion that
bonds to a metal or ion
142
Coordination number
The number of atoms bonded to the metal (# of bonds metal forms)
143
Acidic
H > OH
144
Basic
H < OH
145
Cation
positive charge
146
Anion
negative charge
147
