6 Kinetics
Year 11 DP Chemistry
Particle Collisions
For a reaction to proceed to products, the
reactants must collide with one another.
Rate of reaction
The rate of a reaction is the rate that reactants disappear and
products form. As the reaction proceeds and reactants are used up,
the rate decreases. Since the reacting particles must collide to
react, increasing the rate of collisions increases the rate of reaction.
Rate of reaction: the change in concentration of reactants or products over time
Rate Equations
This graph shows the decrease
in concentration of the reactants
in a reaction.
The instantaneous rate can be
determined at any point along
the graph by measuring the
gradient which is change in
concentration over time
Rate =
Δ[R]
Δt
The average rate of reaction can
be calculated by using the
concentration change over the
entire reaction.
Some Factors Affecting Reaction Rates
• Temperature
• Higher temperature means particles have more kinetic energy,
leading to an increase in reaction rate.
• Concentration
• Increasing the concentration of the reactants increases the
chance of collisions, which increases the rate of reaction.
• Surface area
• Fine powders have more surface area than large pieces leading
to more collisions. For this reason, fine powders create an
explosion hazard when mixed with air
• Catalysts
• Catalysts lower the activation energy, therefore increasing the
rate of reaction
Kinetic Theory of Matter
Kinetic Theory describes the movement of
particles in matter. There are three basic
movements:
• Vibrational (vibrate)
• Rotational (spin)
• Translational (move from here to there)
Solids vibrate around fixed positions as the
atoms are strongly attracted holding a
definite shape.
Liquids have particles that are still
close together, but they can move
around. This is why they flow.
Gases have much greater movement
and all three motions are possible.
Source: http://www.saburchill.com/physics/chapters/0098.html
Collision Theory
Collision Theory states that in order for reactions to occur,
successful collisions must take place between reacting
particles. To collide, they must be in motion...
Kinetic Theory of Gases describes
properties that result from the motion of gas
particles.
Some assumptions of the Kinetic Theory of Gases:
1.
2.
3.
4.
5.
A gas consists of a collection of small particles travelling in straight-line motion.
The molecules in a gas occupy negligible volume (that is, they are points).
Collisions between molecules are perfectly elastic (that is, no energy is gained or lost during the
collision, only transferred from one particle to the other).
There are no attractive or repulsive forces between the molecules.
The average kinetic energy of a molecule is proportional to the absolute temperature ( in Kelvin).
(Specifically, 3kT/2, where T is the absolute temperature and k is the Boltzmann constant.)
Kinetic Theory - Gases
Source: http://isite.lps.org/sputnam/chem_notes/UnitVI_Gases.htm
Collision Theory
For a reaction to
occur, there must
be successful
collisions. This
means they must
have:
• Enough energy
(>Ea)
• Correct orientation
http://boomeria.org/chemlectures/rates/rates.html
Collision Theory – particle size effect
How will the rate of reaction be affected by the change in particle size?
Collision Theory – Temperature effect
Remember, temperature is a
measure of the average kinetic
energy of the particles.
These particles need energy to
get over the Activation Energy
barrier in order to go to
products.
How will the rate of reaction be affected by the increase in temperature?
Go to: http://www.ltscotland.org.uk/highersciences/chemistry/animations/collision_theory.asp
to see an animation about the effect of temperature
Collision Theory – Concentration effect
How will the rate of reaction be affected by the increase in concentration?
Collision Theory – Pressure effect
How will the rate of reaction be affected by the increase in pressure?
Energy Distribution Curves
Particles in a sample of gas do not all have the same
kinetic energy at a given temperature.
The curve on the right shows that particles can have a
wide range of energy values. The particles over the
threshold energy, are able to react
It also shows that if we increase the temperature, the
number of particles over the threshold also increases.
Note that the area under the two curves is the
same and represents the total number of
particles.
The diagram on the right shows how the energy
distribution curves relate to the enthalpy
diagram.
Notice that the threshold energy is the same as
the Ea.
Maxwell-Boltzmann Energy
Distribution Curves
Gas particles in a container will increase their kinetic energy and their velocity with an increase in temperature.
(K.E. = ½ mv2)
Maxwell-Boltzmann distribution curves show this relationship.
Maxwell-Boltzmann Energy
Distribution Curves
M-B curves can be thought of in terms of probability:
•The highest point in the curve represents the most probable kinetic energy value
•The average kinetic energy is slightly higher due to the asymmetric shape of the curve
Catalysis and energy distribution
Recall that catalysts lower the Activation energy of a reaction.
This means that more particles are energetic enough to react at the same temperature.
See the enthalpy diagram and M-B distribution curve below.
6 Kinetics – 2011 Rob Slider
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