Inorganic Chemistry > Topic 9 - Kinetics I > Flashcards

Topic 9 - Kinetics I Flashcards

(58 cards)

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1
Q

What are the conditions for a successful collision between particles?

A
  • Collision in the right direction

* Collision with at least a certain minimum amount of kinetic energy

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2
Q

What is the activation energy?

A

The minimum amount of energy that must be supplied to particles in order to break bonds and start a reaction.

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3
Q

Hwo does the activation energy of a reaction determine how easily it happens?

A

The lower the activation energy, the more easily the reaction happens.

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4
Q

Do all molecules in a substance have the same kinetic energy?

A

No, there is a distribution.

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5
Q

What is the name for the curve showing the distribution in energies of particles in a substance?

A

Maxwell-Boltzmann Distribution

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6
Q

On a Maxwell-Boltzmann distribution, what is on the x and y axis?

A

x-axis -> Kinetic energy

y-axis -> Number of molecules

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7
Q

Describe the shape of the Maxwell-Boltzmann distribution curve.

A
  • Starts at the origin
  • Steep gradient up to peak, then downwards gradient
  • Gradually plateauing -> Never reaches x-axis

(See diagram pg 112 of revision guide)

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8
Q

Why does the Maxwell-Boltzmann distibution curve go through the origin?

A

No molecules can have zero energy.

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9
Q

How is the activation energy for a reaction shown on a Maxwell-Boltzmann distribution curve?

A

A vertical line at a certain energy up to the curve.

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10
Q

On a Maxwell-Boltzmann distribution curve, how can you tell which particels have sufficient energy to react?

A

All the particles to the right of the vertical ‘activation energy’ line have sufficient energy to react.

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11
Q

Remember to practise drawing out the Maxwell-Boltzmann distribution curve.

A

Pg 112 of revision guide.

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12
Q

Describe how the Maxwell-Boltzmann distribution curve changes when the temperature is increased.

A
  • Shifts to the right -> More particles above activation energy line
  • Peak is more to the right, but lower than before
  • Area under curve is the same
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13
Q

What 5 factors affect the rate of a reaction?

A 
• Temperature
• Concentration
• Pressure
• Catalysts
(• Surface area)
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14
Q

Explain how increasing the temperature affects the rate of reaction.

A

• Increases it
• Because particles have on average more kinetic energy, so:
1) More collisions happen per second
2) More of these collisions happen with sufficient activation energy to react successfully

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15
Q

Explain how increasing the concentration affects the rate of reaction.

A
  • Increases it
  • Because there are more particles per unti volume, so more collisions can occur per second -> More successful collisions per second
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16
Q

Explain how increasing the pressure affects the rate of reaction.

A
  • Increases it (if any reactants are gases!)
  • Because there are more gas particles per unit volume, so more collisions can occur per second -> More successful collisions per second
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17
Q

Explain how a catalyst affects the rate of reaction.

A
  • Increases it
  • Because they lower the activation energy by providing an alternative pathway for the reaction -> So more particles have sufficient energy to react -> More successful collisions per second
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18
Q

Explain how increasing the surface area affects the rate of reaction.

A
  • Increases it
  • Because there is a larger surface area available for collisions, so more can happen per second -> More successful collisions per second
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19
Q

What is collision theory?

A

The way in which reactions of particles can be explained by their movement, etc.

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20
Q

When will increasing the pressure increase the rate of reaction?

A

When at least one of the reactants is a gas.

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21
Q

What is the reaction rate?

A

The change in amount of reactant or product per unit time.

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22
Q

Are there any defined units for rate of reaction?

A

No, it depends on what is being measured.

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23
Q

How can the rate of reaction be worked out from a graph?

A
  • Work out the gradient of the graph at the given time
  • Pick appropriate units (y units / x units)

(This is assuming the graph is of reactant or product (y) against time (x))

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24
Q

How can you work out the gradient from a curved graph?

Study These Flashcards
A

Draw a tangent at the given point.

25
For a graph of concentration against time, what is the unit for the rate of reaction?
mol / dm^3 / min
26
What is the initial rate of a reaction?
The rate at the start of a reaction.
27
How can the initial rate of reaction be worked out from a graph?
* Draw a tangent at t = 0 * Work out the gradient * Pick appropriate units
28
In an experiment which times how long it takes for x amount of reactant to be used up or product to be formed, how can you work out the rate of reaction?
* Rate of reaction = Amount of reactant used or product formed / Time taken * Pick appropriate units
29
If a reaction takes 10 seconds to produce 20cm^3 of a gas, what is the rate of reaction?
2cm^3/s
30
In an experiment which times how long it takes for x amount of reactant to be used up or product to be formed, what is the rate proportional to?
1/Time
31
A student measures the time taken for a colour change to occur in a reaction as he varies the concentration of a reactant, A. His results are shown in the table. Calculate the relative rates of reaction. ``` [A] = 0.10 mol/dm3 -> 124s to colour change [A] = 0.15 mol/dm3 -> 62s to colour change [A] = 0.20 mol/dm3 -> 25s to colour change ```
1) Relative rate of each reaction: When [A] = 0.10 mol/dm3 -> 1 / 124 = 0.00806 s^-1 When [A] = 0.15 mol/dm3 -> 1 / 62 = 0.0161 s^-1 When [A] = 0.20 mol/dm3 -> 1 / 25 = 0.0400 s^-1 2) Divide by the smallest relative rate to get the rates as the smallest whole number: 0.0081 : 0.016 : 0.040 = 1 : 2 : 5
32
Remember to practise calculating reaction rates from graphs and experiments.
Pgs 114-115 of revision guide.
33
Explain how a catalyst increases the rate of reaction.
* A catalyst increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy, so a greater proportion of collisions result in a reaction. * The catalyst is chemically unchanged at the end of the reaction.
34
Is a catalyst used up in a reaction?
No, it remains chemically unchanged.
35
Do catalysts take part in a reaction?
Yes, but they're remade at the end of it.
36
How many reactions can a catalyst catalyse?
Usually only 1.
37
What are the two types of catalyst?
* Heterogeneous | * Homogeneous
38
What is a heterogeneous catalyst?
One that is in a different physical state to the reactants.
39
Give an example of a heterogeneous catalyst.
In the Haber process, gases are passed over a solid iron catalyst.
40
On what part of a heterogeneous catalyst does a reaction take place?
The surface.
41
How can the effectiveness of a solid heterogeneous catalyst be increased?
* Increasing its surface area | * Because this increases the number of molecules that can react at the same time -> Increasing the rate of reaction
42
Describe how a solid heterogeneous catalyst works.
1) Reactant molecules arrive at the surface and bond with the catalyst -> A*d*sorption 2) Bonds been the reactant's atoms are weakened and break up -> Lower activation energy -> This forms radicals 3) Radicals join to make products 4) Products detach from the catalyst -> Desorption
43
What is the name for the reactants bonding to a heterogeneous catalyst?
Adsorption | NOTE: It's adsorption, not absorption
44
What is the name for the reactants detaching from a heterogeneous catalyst?
Desorption
45
Remember to revise how a heterogeneous catalyst works.
Pg 116 of revision guide.
46
What is a homogeneous catalyst?
* One that is in the same physical state as the reactants. | * Usually, it is an aqueous catalyst for a reaction between two aqueous solutions
47
What state is a homogeneous catalyst in?
Usually aqueous (so it catalyses reactions between aqeuous solution).
48
Describe how a homogeneous catalyst works.
1) Reactants combine with the catalyst -> Lower activation energy E1 2) Intermediate species is formed 3) Intermediate species then reacts to form product and reform catalyst -> Lower activation energy E2 (i.e. There are two reactions, each with a lower activation energy)
49
How many stages are there in a reaction catalysed by a homogeneous catalyst?
There are 2 separate reactions, each with a reaction energy lower than the original uncatalysed reaction.
50
How many activation energies are there in a reaction catalysed by a homogeneous catalyst?
2
51
Describe the reaction profile for a reaction catalysed by a heterogeneous catalyst.
* Like the uncatalysed curve, except below it and with a lower peak * Reaches the products line at the same point (See diagram pg 116 of revision guide)
52
Describe the reaction profile for a reaction catalysed by a homogeneous catalyst.
* Two smaller peaks below the origianl curve, each with a separate activation energy * After these peaks, the line is the same as the uncatalysed curve
53
In the reaction profile for a reaction catalysed by a homogeneous catalyst, where is the intermediate species formed?
In the trough between the two small peaks.
54
In the reaction profile for a catalysed reaction, does the curve reach the products line before the normal uncatalysed reaction line?
No, it reaches the products line at the same time.
55
Remember to practise drawing out reaction profiles for an: • Uncatalysed reaction • Homogeneous catalyst reaction • Heterogeneous catalyst reaction
Pgs 116
56
How is a catalyst shown on a Maxwell-Boltzmann distribution?
It shifts the vertical 'activation energy' line to the left, so a greater proportion of the particles are to the right of it.
57
What are the benefits of using catalysts in industry?
1) Lowers cost due to lower needed temperature and pressure 2) Faster production 3) May improve properties of product
58
Give an example of a case where using a catalyst improves the properties of the product.
``` POLYETHENE WITHOUT CATALYST: • Less dense • Less rigid POLYETHENE WITH CATALYST: • More dense • More rigid • Higher melting point ```
Inorganic Chemistry
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