Chemistry Ii Ch 13/14

Front 1

What is kinetics?

Back 1

1. Study the factors that control rate of reaction.

Front 2

What things control rate of reaction? (5)

Back 2

1. Temperature
2. Nature of reactants
3. Concentration
4. Mechanism
5. Catalysis (can be faster/slower)

Front 3

How do you define reaction rate and what is one thing you should always remember?

Back 3

1. Rate of reaction = ΔConcentration/ΔTime
(concentration in molarity)
2. ΔConcentration ALWAYS POSITIVE!!!!!! put negative sign in front of reactant to make positive.

Front 4

What do you need to put in front of the equation for a rate of reaction in respect to a reactant?

Back 4

1. Negative sign (reaction rate needs to always be positive)

Front 5

What do you do to reaction rates such as 3H2 + N2 --> 2NH3?

Back 5

1. Multiply 3H2 by 1/3 and 2NH3 by 1/2 to get all to equal one mole so you can calculate reaction rate.

Front 6

If something has two moles per each mole of product, what does that mean about reaction rate?

Back 6

1. Its reaction rate is twice as fast because twice as much disappears for each mole of product.

Front 7

How do you find instantaneous rate?

Back 7

1. Draw a tangent to the curve at the point in time and then find the slope (still concentration/time)
2. We are using tangent line instead of avg line this time, so that is why instantaneous.

Front 8

What are two ways you can measure reaction rate?

Back 8

1. Concentration (disappearance of product or appearance of product)
2. Speed of reaction (continuous monitoring for fast and sampling for slow)

Front 9

What are the three methods of continuous monitoring and how do they work?

Back 9

1. Polarimetry - measuring how plane polarized light is rotated. Handedness/chiro = rotation of light shows how concentrated it is based on how far it rotates past perpendicular (used in organic chem)
2. Spectrophotometry - measuring how sample interacts with light.
3. Total pressure - monitor how pressure goes up/down

Front 10

What is the key method of sampling and how does it work?

Back 10

1. Gas chromatography - separate a sample based on vaporize passing through a column.
2. You can see how fast particles move based on polarity, size, or ion exchange.
3. Area under the curve represents the concentration

Front 11

Explain how the nature of reactants affects the reaction rate?

Back 11

Physical Condition
1. Smaller pieces (powder = fast, chunks = slow)
2. In solution - get an intimate mixture (better mixed = faster reaction)
3. Ions tend to react more quickly than neutral molecules

Front 12

Explain how temperature affects the reaction rate?

Back 12

1. Need to have 2 species collide with sufficient energy and orientation in order for reaction to go forward.
2. Chemical RXN: collisions affect chemical change - if you increase temperature, you increase probability of successful collisions.

Front 13

What is key in controlling how fast/slow a reaction goes?

Back 13

1. Activation energy
2. Big (hill) activation energy required = slow reaction
3. Small (hill) activation energy required = fast reaction

Front 14

Why does hydrogen peroxide not react at room temperature in the bottle?

Back 14

1. Because it requires more activation energy to react and this occurs due to blood enzymes that help lower the activation energy when they come in contact with each other.

Front 15

Explain how reactant concentration affects the reaction rate?

Back 15

1. Any reactant has the potential to change how quickly/slowly a reaction goes by changing concentration.
2. Many times, if bonds have to be broken for reaction to occur, this will be the biggest barrier and thus concentration can alter this. More particles = more bonds CrNH3 to be broken in same amount of time for Cl- to attach.

Front 16

What is the rate law?

Back 16

1. Rate of Reaction = k [A]^x [B]^y

Front 17

What is the reaction order in terms of the rate law (ie: x and y exponents)

Back 17

0 order: changing concentration has no effect on rate
1st order: rate changes by same factor as concentration.
2nd order: the rate changes by square of the factor the concentration changes.

Front 18

Explain zero order reactions and how it can be shown graphically and mathematically.

Back 18

1. Rate doesn't depend on concentration - [ ]^0 = 1, so depends on constant value (k)
2. [A] = [A]o - (k)(t)
3. k unit is M/s (k will be negative because of negative slope)
4. Half Life equation: T1/2 = [A]. / 2(k)

Front 19

Explain first order reactions and how it can be shown graphically and mathematically.

Back 19

1. Concentration directly affects rate of reaction. (double [] you double reaction rate.
2. k unit is 1/s (negative slope so it is negative)
3. Half life is constant and independent of concentration (t = .693/k)
4. [A] = [A]o * (e^-kt)

Front 20

Explain second order reactions and how it can be shown graphically and mathematically.

Back 20

1. 1/[A] = kt + 1/[A]o
2. t1/2 = 1/k[A].

Front 21

- If a graph in [A] vs time form is straight line, then exponent on A is ________. Rate constant = ________.
- If a graph in Ln[A] vs time form is straight line, then exponent on A is ________. Rate constant = ________.
- If a graph in 1/[A] vs time form is straight line, then exponent on A is ________. Rate constant = ________.

Back 21

1. [A] straight line - 0 - (-slope)
2. Ln[A] straight line - 1 - (-slope)
3. 1/[A] straight line - 2 - (+slope)

Front 22

What does the order tell us?

Back 22

1. Tells us how many molecules have to collide in order for reaction to start, which is the slow step.

Front 23

If a concentration has a rate order of zero, when does that reactant play a role in the reaction?

Back 23

1. AFTER the slow step. It would take place during the fast step.
2. If a concentration had an order of two, it would require two molecules to collide at that slow, first step.

Front 24

What is temperature's effect on rate? Why?

Back 24

1. As you increase temperature, you increase rate of reaction.
2. This occurs because more temp = more KE = more collisions = more turnover. ***

Front 25

What happens to rate if the collisions to start a reaction have to be very specifically oriented?

Back 25

1. Slower reaction because takes longer/harder to get into that position.

Front 26

What happens at the top of the "hill" in activation energy?

Back 26

1. The bond that needs to be broken for reaction to occur is stretched and almost broken. Once broken, it tumbles down the hill!
2. It may rotate one way or another (tipping point)
CH3-N---C ------> CH3 - C---N

Front 27

What are the two types of activated complexes at the top of the hill in activation energy?

Back 27

1. Intermediate - can detect species, at top there is a small notch that it sits so you can detect.
2. Transition state - can't detect species exists at top because once it hits top it goes down.

Front 28

What is the activation energy of reversal? more or less energy overcome?

Back 28

1. Harder to overcome because have to overcome heat of reaction and activation energy to go backwards.

Front 29

What does a larger Ea mean in terms of rate?
Larger T?

Back 29

1. Larger Ea = small k, slow rate
2. Larger T = large k, fast rate

Front 30

How are collisions involved in the theory of kinetics?

Back 30

1. Collisions of molecules in system have to have enough energy to overcome activation energy.
2. Collisions of molecules in system have to have correct orientation to create activated complex.

Front 31

What does a larger and smaller "A" mean? (remember A is the measure of the effect of collisions in k=Ae^-ea/RT

Back 31

1. Larger A - higher probability of collisions being successful.
2. Smaller A - need more specific orientation to be effective.

Front 32

What is the rate determining step? What do you use to create the rate law?

Back 32

1. The slowest step!
2. Use reactants in the slow step to create rate law.
3. If equilibrium sign, the slow step is most likely second and you will have to manipulate to get the rate law.

Front 33

What things do you need to remember in terms of reaction mechanisms and figuring out the rate law?

Back 33

1. Mechanism has to support Rate Law
2. Mechanism has to add steps to get overall chemical RXN.
3. If rate law has zero order, that reactant appears in fast step.

Front 34

Define reaction mechanism...

Back 34

1. Series of reaction/collisions proposed to support the rate law.

Front 35

What are mechanism intermediates?

Back 35

1. Chemical species produced in one step of mechanism and then used in later step.
2. Does not show up in the overall chemical RXN.

Front 36

Define molecularity? What is typical?

Back 36

1. How many chemical species are on the reactant side of any step in the mechanism?
2. What has to collide?
3. 1 particle - unimolecular, 2 particle - bimolecular (3 is rare because orientation tough to get, usually ends up being two steps)

Front 37

Explain how to find rate law with mechanisms with a fast initial step.

Back 37

1. Look for intermediate and then trace back to previous steps and put in reactants that make up intermediate.
2. You need to have rate law describe reactants that are in OVERALL reaction.

Front 38

Define catalysts.

Back 38

1. Species that change the rate of RXN without being consumed. (ie: not in the overall chemical RXN)
2. They will show up as a reactant on one side of equation and then a product on the other side. Brought in, spit out!

Front 39

How does the activation energy (Ea) change with a catalyst?

Back 39

1. Ea = smaller with catalyst (this will show on the energy profile graph)

Front 40

What is the difference between a heterogenous and homogeneous catalyst?

Back 40

1. Homogeneous: same physical state as reactants (ie: gas phase catalyst = gas phase reactants)
2. Heterogenous: different physical state than reactants (usually a solid).

Front 41

What is chemical equilibrium in terms of reactions? Dynamic equilibrium?

Back 41

1. Rate of forward reaction = rate of reverse reaction
2. Dynamic: Rate = continuous exchange of reactions to products and products to reactants.

Front 42

How do you "know" when you have reached equilibrium in terms of concentration in a reaction?

Back 42

1. Change in [ ] will equal ZERO
2. Appears fixed, although reactants and products moving back forth at same rate.

Front 43

What occurs when you start with same amount of reactants and products?

Back 43

1. Equilibrium ratio will be the same no matter if you start with products and go reverse or start with reactants and go forwards.

Front 44

What is the equilibrium constant (Keq)?

Back 44

Keq = Π [products] / Π [reactants]
Π = multiply

Front 45

How would you write the Keq for 2N2O --><-- 4NO2 + O2

Back 45

Keq = [NO2]^4 [O2] / [N2O] ^2

Front 46

What does the value of Keq imply?

Back 46

Keq > 1 shows more product involved, so "the position of equilibrium favors the products"
Keq < 1 shows more reactant involved, so "the position of equilibrium favors the reactants"

Front 47

What happens to the Keq when you flip a chemical reaction?

Back 47

1. It becomes the inverse of the regular Keq.
2So2 + O2 --> 2SO3 K = [SO3]^2 / [SO2]^2 [O2]
2SO3 --> 2SO2 + O2 K = [SO2]^2 [O2] / [SO3]^2

Front 48

What happens to Keq when you multiply the reaction by some number? Ie: 3(2SO2 + O2 --> 2SO3)

Back 48

1. The Keq of the original reaction is taken to the power of the coefficient out front (in this case a 3)
Keq ^3

Front 49

When adding equilibrium reactions, how can you denote this?

Back 49

1. You can add the equations or you can multiply the Keq together. Certain elements cancel and then you are left with final Keq.
2. K3 = K1 * K2

Front 50

If all reactants/products are in the gas phase, what K can you use?

Back 50

Kp - pressure
Kp = Pc^c * Pd ^d / Pa^a * Pb^b

Front 51

How are Kp and Kc related?

Back 51

Kp = Kc * (RT)^delta N
(delta N = change in number of gas moles) products minus reactants.

Front 52

Using the equation, N2(g) + 3H2 (g) --> 2NH3 (g), a temperature of 400 K and Kc of 6.0 X 10 ^ -2, what is Kp?

Back 52

Kp = 5.6 X 10 ^ -5

Use Kp = Kc * (RT)^ delta N
delta N in this case is (2-4) = -2

Front 53

What is heterogeneous equilibria?

Back 53

1. Reactants and products not necessarily in same physical state (phase)

Front 54

If amount changes, [ ] or pressure will CHANGE. What phases does this include?

Back 54

1. Aqueous
2. Gas

Front 55

If amount changes, [ ] or pressure will stay SAME. What phases does this include?

Back 55

1. Solid
2. Liquid
Pure compounds

Front 56

What does the [ ] or pressure stay the same in pure compounds, such as liquids/solids?

Back 56

1. Since mass directly related to volume (ie: fixed density), if you change volume by X, the mass will also change by X.
2. Concentration stays same
3. Disregard these in Keq, Kc, or Kp equations.

Front 57

What would the Keq be for this reaction? NH3 (aq) + H20 (l) --> NH4 (aq) + OH (aq)

Back 57

Kc = [NH4] [OH] / [NH3]

H20 is left out because it is pure compound (liquid) and therefore concentration does not change since mass directly related to volume (ie: fixed density).

Front 58

What is the reaction quotient? What does it tell us?

Back 58

1. Qc = mass action expression with INITIAL concentrations - predicts direction of change
2. Qc > Kc ..... Too much product - shift reverse
3. Qc < Kc ..... Too much reactant - shift forward
4. Qc = Kc ..... at equilibrium

Front 59

What is a good method to remember when calculating equilibrium concentrations?

Back 59

ICE box
Initial concentration
Change in concentration
Equilibrium molarity
I + C = E
Use this to find X

Front 60

How do you find equilibrium concentrations when given the equilibrium constant and initial concentrations/pressures?

Back 60

1. Set up ICE box
2. Calculate Qc or Qp
3. Use expression in "E" line to put into mass action
4. Solve for X
5. Place value for X into "C" line and solve for "E" line.

Front 61

How do concentration changes affect equilibrium?

Back 61

1. Add more reactant - product gets more concentrated
2. Remove reactants - product less concentrated

Front 62

How does adding or removing reactant (any phase) change equilibrium - which way does it shift?

Back 62

1. Remove reactant - shifts reverse to make product concentration go down to re-establish equilib (Q > K)
2. Adding reactant - shifts forward because more reactant, need even out product side (Q < K)

Front 63

What is a common rule to remember for shifting reactions due to changes in equilibrium?

Back 63

1. Add in - shift away
2. Take out - shift towards

Front 64

What effect does Volume change have on equilibrium?

Back 64

1. Decrease volume - pressure all gases increases - equilibrium shifts to side with FEWER moles of gas.
2. Increase volume - pressure all gases decreases - equilibrium shifts to side with MORE moles of gas.

Front 65

What happens when a non-reactive gas is added to the system of equilibrium?

Back 65

1. Equilibrium stays the same - didn't change the reactants/products, so no change

Front 66

Effect of temperature changes on equilibrium?

Back 66

1. Exothermic: product (releases heat)
2. Endothermic: reactant (needs heat)
- If reaction exothermic: increasing temp = adding product, so shifts reverse.
- If reaction endothermic: increasing temp = adding reactant, so shifts forward.

Front 67

Does a catalyst change position of equilibrium?

Back 67

1. NO, catalyst only makes the establishment of equilibrium come faster.

Front 68

What is the solubility product (Ksp)?

Back 68

1. equilibrium of saturated solution for sparingly soluble salts (very little dissolves)
2. REMEMBER NOT TO INCLUDE Liquids or Solids in Ksp

Front 69

What is the molar solubility?

Back 69

1. S = x
2. Moles of solid salt that can be dissolved

Front 70

Which is more insoluble? S= 3.32 X 10 ^-4 M or S= 7.07 X 10 ^ -5 M?

Back 70

7.07 X 10 ^-5 M (smaller number of moles can be dissolved)

Front 71

If the H30 + increases when something is added to water, then it is an _________. If the OH- increases when something is added to water, then it is a __________.

Back 71

1. Acid
2. Base

Front 72

What are properties of acids/bases?

Back 72

1. Acids:
- sour taste
- react with metal to produce Hydrogen gas
- Litmus test = blue --> red
2. Bases:
- bitter (alkali)
- slippery feel
- metal reacting with water
- Litmus test = red --> blue