1
Q
What must collide in order to react?
A
Reactants
2
Q
Define activation energy
A
The minimum amount of energy needed to break the chemical bond
3
Q
Describe the relationship between collision rates and number of collisions
A
More collisions = more reactions = higher collision rates
4
Q
Describe the relationship between concentration and collision rates
A
Increase concentration = increase collision rates
5
Q
Describe the relationship between temperature and collision rate
A
Increase temperature = increase collision rates
6
Q
What is the purpose of the potential energy diagram?
A
They show the change in potential energy as reactants are converted into products
7
Q
What is true of an exothermic reaction?
A
- The enthalpy of the reaction is less than 0
- Energy is leaving
- Reactants have more energy
- Downhill on phase diagram
8
Q
What is true of an endothermic reaction?
A
- The enthalpy of the reaction is greater than 0
- Energy is coming in
- The products have more energy
- The uphill of the phase diagram
9
Q
What is the relationship between activation energy and the reaction rate
A
A higher activation energy results in a slower reaction rate
10
Q
What three things are true of transition states?
A
- Partially formed and partially broken bonds
- Must go “uphill” to reach the transition state
- It is also called the activated complex
11
Q
Define molecularity
A
The number of reactants involved
12
Q
Define reaction mechanism
A
The entire series of individual chemical steps by which an overall chemical reaction occursW
13
Q
Define elementary step
A
Each step that occurs within a reaction mechanism, it cannot be broken down any further
14
Q
Define reaction intermediate
A
A species that forms in one elementary step and is consumed in another
15
Q
Are reaction intermediates the same as transition states?
A
No
16
Q
Define unimolecular steps
A
SIngle reactant produces one or more moles of product
17
Q
What is the rate law for a unimolecular step with the format: A –> products
A
Rate = k[A]
18
Q
What order is a unimolecular step?
A
Always first order
19
Q
Define bimolecular step
A
Two molecules or atoms collide to form an activated complex resulting in a product of products
20
Q
What is the rate law for a bimolecular step with the format: A + A –> products
A
Rate = k[A]^2
21
Q
What is the format for a bimolecular step with the format: A + B –> products
A
Rate = k[A][B]
22
Q
What is the order of a bimolecular step?
A
Always second order
23
Q
Define termolecular steps
A
Three molecules collide simultaneously resulting in a product of products
24
Q
What is the rate law for a termolecular step with the format: A + B + C –> products
A
Rate = k[A][B][C]
25
What is the rate law for a termolecular step with the format: A + A +B --> products
Rate = k[A]^2[B]
26
What is the rate law for a termolecular step with the format: A + A + A --> products
Rate = k[A]^3
27
What is the order of a termolecular step?
Always third order
28
Define rate determining step
- The step that occurs slower than all of the other steps
- Largest activation energy
29
What does the rate determining step allow us to predict?
The rate law
30
What are the three steps to finding the rate law of a multi-step reaction?
1. Identify the slow step and write the rate law for it
2. Check whether this step contains an intermediate species (if none, then you have your rate)
3. Find a term that can replace the intermediate in the rate law
31
What is the relationship between reaction rate and temperature
As you increase the temperature, you increase the reaction rate
32
What is the relationship between reaction rate and concentration
If the reaction order is greater than 0, then as you increase the concentration you increase the reaction rate because more collisions occur
33
What is true of the collision theory?
It deals with chemicals in the gas phase
34
What is the relationship between reaction rate and chemical nature of reactants?
As you move down alkali metals, they react faster with water
35
What is the relationship between reaction rate and physical state of the reactants?
As you increase surface area, you increase the reaction rate
36
Define catalyst
A substance that lowers the activation energy and therefore speeds up the reaction without being consumed
37
Define homogenous catalyst
A substance that speeds up the chemical reaction by being the same phase as the reactants
38
Define heterogenous catalyst
A substance that speeds up the chemical reaction by being a different phase than the reactnats
39
Define biological catalyst
Enzymes essentially
40
Define positive catalyst
Most catalysts are positive catalysts because they speed up a chemical reaction
41
Define negative catalyst
A catalyst that slows down the reaction
42
Define thermodynamics
The science of the relationship between heat, work, and energy
43
Define first law of thermodynamics
Energy can be transferred but not destroyed
44
Define spontaneous reaction
A process that proceeds in a given direction without needing to be driven by some outside source of energy
45
What type of process is vaporization?
Endothermic
46
What type of process is combustion?
Exothermic
47
What are spontaneous processes driven by?
The tendency of energy to disperse and of matter to become dispersed
48
Define entropy
The measure of the dispersal or disorganization of energy or matter in a system
49
What is the relationship between disorder and entropy?
As you increase disorder, you increase entropy
50
What is the difference between ΔH and ΔS?
ΔH measures total heat content while ΔS measures energy dispersal
51
What is the equation for entropy?
S = k ln W
* where k is the Boltzmann's constant and W is the number of possible micro states (no unit)
52
What is the second law of thermodynamics?
For any spontaneous process, the entropy of the universe increases
53
+ ΔS
Spontaneous
54
-ΔS
Non-spontaneous
55
ΔS = 0
equilibrium
56
What is the entropy equation in terms of the universe?
ΔS universe = ΔS system + ΔS surroundings
57
Create a chart to remember the sign of ΔS and what that means for the process type
58
What is the entropy change for a phase change?
ΔS phase change = q phase change / T phase change
ΔS phase change = ΔH phase change / T phase change
ΔS phase change = ΔH* x mol / T phase change
59
What is the entropy change for a chemical reaction?
ΔS reaction = ∑Sproducts - ∑Sreactants
60
What is the entropy change of the surroundings?
ΔS surroundings = q surroundings / T
ΔS surroundings = -q system / T
ΔS surroundings = -ΔH system / T
ΔS surroundings = -ΔHreaction * mol / T
61
What is the equation for the entropy change of the universe?
ΔS universe = ΔS system + ΔS surroundings
62
If we cannot determine spontaneity without enthalpy or entropy, what equation do we use?
Gibbs free energy
63
What does Gibbs Free Energy tell us?
How much energy are we storing
64
ΔG is a _____ function
state
65
What is true of temperature in the Gibbs Free Energy equation?
it remains constant
66
If ΔG is negative then the reaction is ______ and _____.
Spontaneous and product-favored
67
If ΔG is positive then the reaction is _______ and ______.
Non-spontaneous and reactant-favored
68
Make a chart to remember the temperature effects on spontaneity
69
What is the third law of thermodynamics?
- Essentially everything in the real world has entropy
- The entropy (S not ΔS) of a perfect crystal at absolute zero K is 0 J / mol K
70
Define standard molar entropy
The entropy gained by converting 1 mole of the substance from a perfect crystal at 0K to its standard state
71
Define standard entropy change
Change in entropy when reactant and products are in their standard states at the desired temperature
72
What is the standard state of a solid?
Pure substance at 1 atm
73
What is the standard state of a gas?
Pure, ideal gas at 1 atm
74
What is the standard state of a solution?
Aqueous solution with a concentration of 1M at 1atm
75
What is ΔS*rxn equation?
(∑n products S*products) - (∑n reactantsS*reactants)
76
What is the Gibbs free energy of formation equation?
ΔG*rxn = (∑n products G*f products) - (∑n reactants G*f reactants)
77
What is the difference between S and ΔS?
S= amount of entropy
ΔS= change in entropy
78
What does Gibbs free energy tell us?
- Predicts spontaneity
- Energy available for useful work
79
What does ΔH tell us?
- change in enthalpy
- measures heat exchange
- endothermic/exothermic
80
What does ΔS tell us?
change in entropy
81
What is the equation for forward/reverse reactions?
ΔH = Ea (forward) - Ea (reverse)
82
Do catalysts appear in the balanced equation?
No, catalysts do not appear in the balanced equation
83
What do catalysts do
Speed up a reaction by lowering the activation energy
84
How do catalysts appear in the rate law?
They appear as the rate-determining step
85
What is the rate- determining step?
The slowest step in the reaction
86
Are catalysts recycled?
Yes
87
In a reaction mechanism, how do you know which species is the catalyst?
It is consumed first in the reaction and produced later
88
In a reaction mechanism, how do you know which species is the intermediate?
It is produced first in the reaction and produced later
89
How do you determine if a species is a catalyst or intermediate?
Catalysts are used first produced later, intermediates are produced first and used later
90
How do you know which step is the rate-determining step?
Whichever step is the slow step will always be the rate-determining step. This means the overall rate will equal the rate of this step.
91
What role does the coefficient play in a rate law?
The coefficient becomes the exponent of the rate law
92
Difference between concentrated and diluted solution
Concentrated = more solute
Diluted = little solute
93
What is the molarity equation?
M = moles of solute / vol of solution (L)
94
What is the molality equation?
m = moles of solute / mass of solvent (kg)
95
What is the mole fraction equation?
x = moles of solute / total moles in solution (solvent + solute)
96
What is the mole percent equation?
mol % = moles of solute / total moles in solution x 100%
97
What is the volume percent equation?
% v / v - volume of solute / total vol of solution x 100
98
What is the parts per billion equation?
ppb = mass of solute (mg) / total mass of solution (kg) x 10^9
99
What is the parts per million equation?
ppm = mass of solute (mg) / total mass of solution (kg) x 10^6
100
What two factors does solubility depend on?
- Entropy
- IMFs
101
Make a chart to remember the phases of solute and solvent for each of the possible solution phases
102
What are the three steps of solution formation?
1. The solute is separated into its constituent particles (endothermic)
2. Solvent particles are separated from each other to make room for the solute particles (endothermic)
3. Solute particles and solvent particles are mixed together (exothermic)
103
Define enthalpy of hydration
Change in enthalpy when one mole of gaseous ions dissolves in water
104
How do you increase enthalpy of hydration?
- increase ionic charge
- decrease ionic size
105
If the ΔH solute < ΔH hydration what does that mean for ΔH solution
less than 0 , exothermic, increase temperature and the container feels hot
106
If the ΔH solute is > ΔH hydration what does that mean for ΔH solution?
greater than 0, endothermic, decrease temperature and the container feels cold
107
if ΔH solute is approximately equal to ΔH hydration what does that mean for ΔH solution?
ΔH solution is approximately equal to 0 so there is no noticeable change
108
Define saturated solution
Maximum possible amount of dissolved solute at equilibrium (solvent and solute = dynamic equilibrium)
109
What happens to a saturated solution?
- any solute added will not dissolve
- concentration will not change
110
Define unsaturated solution
- Less solute than saturated
- Dynamic equilibrium has not yet been reached
- Concentration is below the saturation concentration
111
What does the saturation concentration depend on?
Temperature and pressure
112
What do solubility curves show?
The maximum saturation/solubility
113
Above the line in a solubility curve means
Supersaturated
114
Below the line in a solubility curve means
Unsaturated
115
If ΔH solution < 0 what happens to solubility
Exothermic, solubility increases as temperature decreases
116
If ΔH solution is > 0 what happens to solubility?
Endothermic, solubility increases as temperature increases
117
ΔH solution is always negative for ____
gases
118
In gases, solubility always _____ as temperature ____
decreases, increases
119
Does pressure affect solubility of solids or liquids?
NO
120
What is the general idea of Henry's law?
As you increase the pressure of gas, you increase the solubility
121
What is the solubility of gases equation?
S = k_HP_gas
122
What is the equation for ΔH and reversible reactions?
ΔH = E_a forward - E_a reverse
123
What is the relationship between rate constant and activation energy?
The largest rate constant is the lowest activation energy
124
general chemistry 2
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