Topic 5- Kinetics Flashcards Preview

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Flashcards in Topic 5- Kinetics Deck (8)
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What is the definition of rate of reaction?

What is this usually given as?

What are the units for rate of reaction?

What three things does the collision theory state?

What is the steric effect?

Describe the following ways to increase the rate of reaction according to collision theory...

More frequent collisions

More successful collisions

Describe what increasing the surface area does in a reaction?

Change in concentration of a substance per unit time

Change in concentration of a substance per second

mol dm-3 s-1

-Particles must collide before a reaction can take place
-Not all collisions lead to a reaction
-Reactants must possess at least a minimum amount of energy (activation energy)

Particles must approach each other in a certain relative way

Increase particle speed or have more particles present

Give particles more energy or lower the activation energy

Increases chances of a collision- more particles are exposed.


What are three consequences of increasing the temperature in a reaction?

What are two consequences of increasing the pressure in a reaction?

What does it mean the more particles there are in a given volume?

What two things does this then result in?

What are two consequences of increasing concentration in a reaction?

What does low concentration mean?

What does higher concentration mean?

-Increasing the temperature increases the rate of a reaction
-Particles get more energy so can overcome the energy barrier
-Particles speeds also increase so collisions are more frequent

-Increasing the pressure forces gas particles closer together
-This increases the frequency of collisions so the reaction rate increases

The greater the pressure

The greater the pressure, the more frequent the collisions. The more frequent the collisions, the greater the chance of a reaction

Increasing concentration = more frequent collisions = increased rate of reaction

Fewer collisions

More collisions.


What is the definition of activation energy?

The minimum amount of kinetic energy particles need to react.


How do you get a Maxwell-Boltzmann distribution?

What is the area under a Maxwell-Boltzmann distribution curve equal to?

Why does the curve start at (0,0)?

What are some molecules doing?

What are most molecules doing and so what are their energies?

What do some molecules have more than and what can only these molecules do?

What does the peak of the curve represent?

Where is the mean (average) energy of all the molecules?

If you plot the graph of the number of molecules in a gas with different kinetic energies

The total number of molecules

Because no molecules have zero energy

Some molecules are moving slowly

Most molecules are moving at a moderate speed- their energies are somewhere in the middle

Some molecules have more than the activation energy. These are the only ones that can react

The peak of the curve represents the most likely energy of any single molecule

A bit to the right of the peak.


What are two things that particles have in a reaction?

What do no particles have?

In what two ways do particles exchange energy?

What two things can this result in the particles doing?

What two things happen if you increase the temperature of a reaction?

What two things will this result in a greater proportion of molecules having?

What does this do to a Maxwell-Boltzmann distribution curve?

What two things happen at higher temperatures?

Why can quite small increases in temperature lead to quite large increases in reaction rate?

-Some have very low and some have very high energies/velocities
-Most have intermediate velocities

No particles have zero energy/velocity

Through collisions between themselves and within the walls of the containers

Losing or gaining energy

The particles will on average have more kinetic energy and will move faster

At least the activation energy and be able to react

This changes the shape of the Maxwell-Boltzmann distribution curve- it pushes it over to the right

Because the molecule are flying about faster, they'll collide more often

Because you get both of these effects happening at once (more collisions and more energetic collisions).


What are all reactions accompanied by?

Why does the enthalpy rise as the reaction starts?

What two things then happen to the enthalpy?

What does the area under the curve beyond the activation energy correspond to on a Maxwell-Boltzmann distribution curve?

What happens two things happen if a catalyst is added?

Changes in enthalpy

Because energy is being put in to break bonds

It reaches a maximum then starts to fall as bonds are formed and energy is released

The number of molecules with sufficient energy to overcome the energy barrier and react

The activation energy (Ea) is lowered- Ea will also move to the left.


What is the definition of a catalyst?

How do catalysts work by?

What do catalysts not affect in a reaction?

Why are catalysts safer and cheaper?

What happens to catalysts at the end of the reaction?

What do catalysts no affect the position of?

Catalysts alter the rate of a chemical reaction without being chemically changed themselves

Catalysts work by providing an alternative reaction pathway which has a lower activation energy

Catalysts do not affect the enthalpy change of the reaction

Because using catalysts avoids the need to supply extra heat

Catalysts remain chemically unchanged at the end of the reaction

Catalysts do not affect the position of any equilibrium.


Give examples of the following advantages of why catalysts are widely used in industry:

Allow reactions to take place at lower temperatures

Enable different reactions to be used

Are often enzymes

Have great economic importance in the industrial production of substances

Can reduce pollution

-Save energy (by lowering Ea)
-Reduces CO2 output

-Better atom economy
-Reduces waste

-Generate specific products
-Operate effectively at room temperatures

Eg Poly(ethene); sulphuric acid; ammonia; ethanol

Eg Catalytic converters.