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Chem 109 > Exam 3 > Flashcards

Exam 3 Flashcards

1
Q

Chemical Reaction

A

process that changes 1 or more substances into 1or more new substances

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2
Q

What do rates of reaction indicate about equilibrum?

A

system at equilibrium when rate of forward reaction = rate of reverse reaction

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3
Q

What is K

A

Equilibrium Constant;
[products]^coefficients/
[reactants]^coefficients

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4
Q

What changes the value of K

A

temperature ONLY

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5
Q

Homogeneous Equlibreum

A

all reactants and products are present in the same phase

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6
Q

Heterogeneous Equilibrium

A

reactants and products present in 2 or more different phases

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7
Q

Product Favored Reaction

A

spontaneous in the forward direction

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8
Q

Reactant Favored Reaction

A

non spontaneous/ spontaneous in the reverse direction

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9
Q

K>1
K=1
K<1

A

K>1: equilibrium favors products
K=1: equal product and reactant conc at equilibrium
K<1: equilibrium favors reactants

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10
Q

Standard State

A

commonly accepted state of conditions used as a reference point
-Gas phase: 1 bar/ 1 atm
-Liquid/solid Phase: pure substance
-Solution (aq) phase: 1M

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11
Q

ICE tables

A

I= initial conc
C= change in conc
E= equilibrium conc
given 2 pieces of info -> find third

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12
Q

What is Q

A

Reaction Quotient:
[products]^coefficients/
[reactants]^coefficients
-measures the relative amounts of products and reactants present during a reaction at a particular point in time

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13
Q

Q=K
Q<K>K</K>

A

Q=K: at equilibrium
Q<K: spontaneous in forward direction
Q>K: spontaneous in reverse direction

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14
Q

Le Chatelier’s Principle

A

when a chemical system is at equilibrium and experiences a disturbance, the system reacts to achieve a new equilibrium

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15
Q

Concentration Impact on Equilibrium

A

Increase reactant/ decrease product conc -> shift right and visa vera

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16
Q

Pressure and Volume Relationship

A

inverse relationship: decrease in volume -> increase in total pressure AND increase in partial pressure/ concentration

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17
Q

Pressure and Volume Equilibrium Shift

A

Equilibrium Shifts to the side with fewer molecules/ more space

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18
Q

∆rH

A

enthalpy change for gas phase chemical reaction
equ: ∆rH= Σ E bonds broken - Σ E bonds formed

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19
Q

Hess’s Law

A

if a process can be written as the sum of several steps, the total enthalpy change is the sum of all of the steps

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20
Q

∆fH

A

standard heat of formation

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21
Q

Entropy (S)

A

measure of how spread out the energy in a system is

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22
Q

Microstate

A

one possibility/way to distribute energy in a system
-more microstates -> increased modes of motion -> greater S

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23
Q

Modes of Motion (3)

A

translation, rotation, vibration
-motions of each particle in a system

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24
Q

W

A

of available microstates

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25
Standard Molar Entropy (S°)
S°= k ln(w) k= constant; gas constant/ avogadro's number
26
3rd Law of Thermodynamics
the entropy of a perfect crystal (undisturbed lattice) at absolute 0 (0 k) is equal to 0 -only 1 way to allocate 0 kinetic energy/ absence of kinetic energy -> 1 microstate -S°= k ln(1) =0
27
change in standard molar entropy equation
∆rS° = ΣS°(products) - ΣS°(reactants)
28
Entropy analogy
-different types of motion= different items on sale -energy= money -total KE= total money one has to spend -microstates= different shopping carts; assortment of goods
29
What happens to S when temperature is increased?
there is an increase in energy in the system which leads to more microstates/ options for energy distribution
30
How do phase changes change S?
change the number of translations and rotation available; more spread out (gas) have more microstates and therefore higher S than solids.
31
What happens to S when you change molar mass?
greater/heavier mass -> closer spacing so more available energy levels, therefore more microstates and greater S
32
What happens to S when the complexity of the molecule is increased (more atoms and IMF)
additional types of motions become available; ex. rotations and vibrations
33
What different ways/ modes of motion can allocate or distribute energy differently (higher S)
-more ways to rotate -larger molar mass more complex bonds/ structures -more/ higher occupied every levels (bc greater spread of energy across electron cloud)
34
2nd Law of Thermodynamics
energy spreads out across as many modes of motion in the universe as possible; as it spreads out more (more microstates) S increases *entropy of the universe is always increasing
35
2nd Law of Thermodynamics S equation
∆S (universe)= ∆S-∆H/T
36
Gibbs Free Energy
all the energy that can be used to do work
37
Equation to find ∆G or ∆G°
∆G°= ∆H°- T∆S° or ∆G= ∆G° + RTln(Q°)
38
Equation to find ∆G if system at 298K/25°C
∆rG°= Σ∆G°(products)- Σ∆G°(reactants) *use free energy of formation values
39
Spontaneous process
entropy increase -∆G ∆S (universe)> 0 K>Q
40
non spontaneous process
entropy decrease +∆G ∆S (universe)<0 Q>K
41
products favored process
-∆G ∆S (universe) >0 K>1
42
reactant favored process
+∆G ∆S (universe)< 0 K<1
43
what happens when H and S have different predictions regarding thermodynamics of reaction?
the reaction is temperature dependent!! -if temp is increased then H has a more significant impact
44
Transition State/ activated complex
state in the midst of bonds breaking and forming where the system is at its highest energy; energetic awkwardness
45
Steric Factor
probability of molecules colliding with the right orientation
46
Rate Law
k [A]^x[B]^y *x and y NOT coefficients; rxn orders
47
What factors impact the rate constant (k)
-temperature (higher temp/faster rxn -> higher rate constant) -phase dependent -Ea -collision frequency -orientation statistics
48
Arrhenius Equ.
k=Ae^(-Ea/RT) -A: frequency of correct orientation -e^(-Ea/RT): fraction of collisions with sufficient energy
49
What does a larger A do to k (rate constant)
increases k
50
What does a higher temperature do to k (rate constant)
increases k
51
What does a lower Ea do to k (rate constant)
increases k because larger fraction of particles have energy > Ea/ sufficient energy for the reaction to proceed (smaller energy barrier)
52
what is K (equilibrium constant) and k (rate constant) relationship?
K= k(forward reaction)/ k(reverse reaction)
53
∆G: -∆S and +∆H
reactant favored at all temps (non spontaneous)
54
∆G: +∆S and -∆H
product favored at all temps (spontaneous)
55
∆G: both ∆S and ∆H positive
entropy driven rxn; product favored at high temp
56
∆G: both ∆S and ∆H negative
enthalpy driven rxn; product favored at low temp
57
V'ant Hoff plot eqution
lnK= (-∆H/R) x (1/T) + (∆S/R) y. =. m x +. b (linear graph)
58
What happens to an endothermic reaction when heat is increased
reaction shifts right (larger K)
59
What happens to an exothermic reaction when heat is increased
reaction shift left (smaller K)
60
Energy Barrier
highest energy of any point of the reaction (top of the curve) *minimum energy needed to rotate around a double bond
61
Thermodynamically stable
reactants are at lower energy than products (so no rxn occurs)
62
thermodynamically unstable
products at lower energy than reactants so reaction will occur
63
kinetically metastable
when a product favored reaction takes place so slowly it is considered to not be occurring (very high energy barrier) **relative comparison ex. diamond ∆G=2.9 kj/mol graphite ∆G=0 diamond is thermodynamically unstable/ kinetically metastable compared to graphite
64
What happens when 2 simultaneous reactions are running?
the reaction with the more negative gibbs free energy (more
65
why do some temperatures only occur at high temps?
because they have more stable products but a larger energy barrier so if sufficient energy not present then other reaction will occur.
66
Work
a process that transfers KE to or from a macroscopic object -∆G= w; w= the max work that can be done to the surroundings

Chem 109 (6 decks)

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