Transcript Reaction Kinetics

Reaction Kinetics and
Equilibrium
How compounds react with each
other
Collision Theory
A reaction is most likely to occur if
reactant particles collide with the
proper energy and orientation
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/
NO+O3singlerxn.html
Reaction Rate Affecting
Factors
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Nature of Reactants
Concentration
Surface Area
Temperature
Presence of a Catalyst
Pressure for Gases
Nature of Reactants
Factors that contribute to reaction
rate:
- Electronegativity
- Ionization energy
- Atomic Radius.
Bonding on Rate of
Reaction
Covalent bonds require more energy during
collisions due to a greater number of bonds
needed to be broken and reformed.
Ionic Bonds are faster to react and require less
energy during a collision.
Concentration
We measure concentration by the number of
moles there are in a L of solution (Molarity).
More moles = more collisions
Ex: Burning paper. When the oxygen concentration
is low the paper burns slowly. Raise the amount of
oxygen the paper burns faster.
http://www.coolschool.ca/lor/CH12/unit1/U01L01.htm
Surface Area
The larger the surface area the easier it
is to react because there are more
chances for collision.
Temperature
• The average kinetic energy of molecules
in a compound.
• The more molecules move the higher
the temperature.
• Higher temperature results in more
collisions
Presence of a Catalyst
A substance whose presence increases
the rate of a chemical reaction.
Catalysts change (decrease) the
activation energy which increases the
rate of reaction
Activation Energy
Activation energy is the minimum
energy required to initiate a reaction.
When particles collide with the right
amount of activation energy it breaks
the existing bond.
Energy
• Kinetic Energy (motion) – this the energy of
work being done
• Potential Energy (static) – the potential for
something to do work
**Remember that there are many types of energy:
- electrical
- thermal
- mechanical
- electromagnetic
- nuclear
Endothermic vs.
Exothermic Reactions
• Exothermic (exit) reactions give off heat
energy during a chemical reaction
• Endothermic (enter) reactions absorb
heat energy during a chemical reaction
http://schools.matter.org.uk/Content/Reactions/BondActivation.html
Potential Energy
Diagram
Used to show the energy released or stored in
endothermic and exothermic reactions
http://www.saskschools.ca/curr_content/chem30/modules/module4/lesson4/p
otentialenergydiagram.htm
Reading Energy
Diagrams
• THE Y AXIS – Potential energy of the
reaction
• THE X AXIS – Reaction as it takes place
over time
• CURVE – Represents the potential
energy at each step of a chemical
reaction
Reading Energy
Diagrams
• EXOTHERMIC REACTION – energy given off
during a reaction
LOOK AT THE CURVE AND SEE IF IT ENDS
AT A LOWER VALUE
The energy of the products is lower than the
energy of the reactants
Reading Energy
Diagrams
• ENDOTHERMIC – energy is absorbed during a
reaction
• The energy of the products is higher than the heat of
the reactants
LOOK AT THE CURVE AND SEE IF IT ENDS AT A
HIGHER VALUE
• ACTIVATION ENERGY – The amount of energy
needed to reach the peak of the curve
SUBSTRACT THE ENERGY AT THE PEAK OF THE
CURVE FROM THE INITIAL ENERGY
Effect of a Catalyst of
Activation Energy
Catalysts Continued
• A catalyst lowers the activation
energy required for the reaction
to occur. By lowering the
activation energy, the chemical
reaction can occur much more
quickly.
Reversible Reactions
• Not all reactions go completely
to completion (all reactants are
used up)
• Instead, some reactions can
occur both forward and reverse
at the same time. A reversible
reaction is symbolized by double
arrows
• 2SO2(g) + O2(g)
2SO3(g)
Equilibrium
The rate at which the products are
formed is at the same rate that the
reactants are formed.
Equilibrium is represented with a double
arrow. Example:
Equilibrium
Entropy
The measure of the randomness or
disorder of a system’s energy.
The greater the randomness the
greater the entropy.
Entropy of Substances
As a substance goes from solid to liquid
to gas, entropy increases.
Systems in nature tend to undergo
changes toward low energy and high
entropy (they want to lose energy and
gain freedom)
Change of State (review)
• A change of state, also called a
phase change, is the conversion
of a substance from one of the 3
states of matter to another. A
change of state always involves
a change in energy.
Heating and Cooling
Curves
Animation
• Shows how a substance
changes states at each
temperature increase over time.
Phase Diagram
Equilibrium Review
• Some chemical reactions are
reversible
• When a reaction is occurring
both forward and reverse at the
same rate, equilibrium is
reached
• Example:
N2 + 3 H2
2 NH3 + energy
Video
Le Chatlier’s Principle
• When a system at equilibrium is
subjected to a stress (a change
in concentration, temperature,
or pressure), the equilibrium will
shift in the direction that tends
to counteract the effect of the
stress.
Le Chatlier’s Principle
• Example:
N2 + 3 H2
2 NH3 + energy
If the concentration of Nitrogen is
increased, the reaction will shift
to the right and favor the
products side, making more
ammonia and giving off more
heat.
Le Chatlier
• Example:
N2 + 3 H2
2 NH3 + energy
If the temperature is increased,
the reaction will shift to the left
and favor the reactants side,
making more N and H
Effect of Pressure on
Equilibrium
• Example:
N2 + 3 H2
(4 moles gas)
2 NH3 + energy
(2 moles gas)
If the pressure is increased, the
reaction will shift to the right,
favoring the side with the lower
number of moles of gas
Heat and Temperature
• Heat (J or calories) – a transfer of energy
from a body of higher temperature to a
body of lower temperature. Thermal energy
is associated with the random motion of
atoms and molecules.
• Ex) steam v water
• Temperature – a measure of the average
kinetic energy of the particles of a
substance. Temperature is not a form of
energy.
• Question: What state of matter has the
most energy? GAS
Heat Calculations
• Specific Heat Capacity – the
amount of energy required to
raise the temperature of a
substance one degree.
• Ex: Takes more energy to heat
a swimming pool than a cup of
water
• Specific heat Capacity of H2O
(l) = 4.18 J/g X C
Calculations con
• Fusion – melting
• Heat of Fusion – the amount of
heat needed to convert unit
mass of a substance from a
solid to a liquid at constant
temperature.
• - when ice is melting the kinetic
energy stays the same
• Heat of Fusion of H2O = 334 J/g
Calculations con
• Heat of Vaporization – the
amount of heat needed to
convert a unit mass of a
substance from a liquid to a gas
at constant temperature.
• - when ice is boiling the kinetic
energy stays the same
• Heat of V of H2O = 2260 J/g
• TABLE B in Reference Table
Calculations con.
• Q = mC∆T
• Q = mHf
• Q = mHv
Vapor Pressure
• In every liquid, some particles are far
enough away from each other to be
considered gas but are pushed down by
atmospheric pressure. When in liquid,
some particles are far enough apart to
escape their neighboring molecules and
enter the gas phase (vapor). As
temperature increases, particles gain
more energy and more particles escape
from the surface. The pressure these
gaseous particles exert is called vapor
pressure. As temp increases, vapor
pressure increases.
Vapor Pressure con.
• Vapor pressure eventually builds
up enough to equal atmospheric
pressure. When it surpasses
atmospheric pressure, the liquid
boils and allows the gaseous
particles to escape.
• Vapor Pressure – pressure
gaseous particles exert upward