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Kinetics and Equilibrium
A chemical reaction requires an effective
collision.
Effective collision requires:
Proper orientation (angle)
Sufficient energy
The more effective collisions, the faster
the rate of reaction.
Rate=the amount of product/time
Factors Affecting
Reaction Rate
- The rate of a reaction can be modified
- Collision theory explains why rate changes occur
6 Factors that affect the
Reaction Rate
1) Temperature
A measure of the average kinetic energy.
- raising temperature=increase in reaction rate
- If you add energy to raise the average kinetic energy
of a substance the particles have more energy
and collide more.
*more energy + more effective collisions = faster
reactions
(collision theory)
Temperature and particle collisions
August 2010
2) Concentration
1) Adding more reactants increases the rate of
reaction [more collisions].
2) In gases: concentration is increased by
increasing pressure [squeezes particles closer].
More stuff = more collisions
3) Surface Area of Reactants
1) Increase of Surface Area = Increase Rate of
Reaction = More Collisions
2) Make the particle size smaller
Ex. What melts quicker, an ice cube or a crushed ice cube? Any reaction involving a solid can only take place at the surface of the solid. If the solid is split into several pieces, the surface area increases. What effect will this have on rate of reaction? The smaller the pieces, the larger the surface area. This means more collisions and a greater chance of reaction. This means that there is an increased area for the reactant particles to collide with. **low surface area high surface area Surface area and particle collisions
- Pressure**
- Little or no effect on solids and liquids.
- Has a huge effect on gases.
Increased Pressure = More Collisions = Increased
Rate
Potential energy diagram: endo June 2011 January 2012
June 2011 ^ Sometimes reactions can go in both forward
and reverse reactions.
A double arrow indicates a reaction is
reversible and that both reactions are
occurring at once.
In lab, you may have had some poor yield
results. Sometimes this is because the
reaction may have been reversible and all of
your product was not formed…the reaction
did not go to completion.
EQUILIBRIUM Equilibrium is when forward and reverse reactions occur at the same rate. A + B C
- A state of balance between the rates of two processes. *Must occur in a closed container. No reactant or product can leave the container. CO2 (g) CO2 (aq) Solution equilibrium is physical equilibrium that occurs in a can of soda.
August 2011
KINETICS
1)Explains how fast reactions occur
2)Identifies the factors that cause or have
an influence on a reaction.
Rate = the number of moles formed (or consumed) in a unit of time. Net Reaction = simplified reaction that sums up all the reactions that occur.
January 2012
1)Energy is released when bonds are formed. 2)Systems of low energy are more stable than systems of high energy. 3)The more energy released = the stronger the bond. REVIEW
June 2011
COLLISION THEORY
- Atom ions and molecules can form a chemical bond when they collide.
- You must have an effective collision to form a bond. Le CHATELIER’S PRINCIPLE
- Any change in temperature, concentration or pressure on an equilibrium system is called a STRESS.
- Le CHATLIER’s explains how a system at equilibrium responds to relieve stress.
- A system moves away or “shifts” from the added stress.
1)Changes rates of both forward and reverse reactions equally. 2)Equilibrium may be established quicker. 3)Does not change any equilibrium concentrations. EFFECT OF A CATALYST
TWO FACTORS THAT CAUSE CHEMICAL
AND PHYSICAL CHANGES TO OCCUR:
- Enthalpy: tendency to change to a state of lower energy.
- exothermic reactions move to lower energy
- products have less PE than reactants.
- the drive towards lower energy is called a drive to lower enthalpy.
- Entropy: tendancy to change to a state of greater randomness or disorder.
- systems in nature tend to favor disorder! SOLID LIQUID GAS Low entropy Great order Maximum Randomness High entropy
LOW ENTHALPY=
GREATER ENTROPY
January 2012
