Rates of Reaction & Equilibrium
Part 1:Rates of Reaction
Collision Theory
• Atoms/molecules must collide in order to react.
• Increasing the rate of collision will increase how fast a reaction takes place
Effective collisions depend upon:
• Nature of Reactants • Concentration • Temperature • Surface Area • Catalysts
Activation Energy
• The energy put in to a reaction to get it started.
• Small activation energies can be as simple as heat from the classroom, or the spark from a striker.
• Large activation energies can be a lot of heat (like baking a cake).
• Creates an “activated complex”.
Activation Energy is B
Activated Complex
Catalysts
• lower the activation energy, thus change the rate of rxn.
• change the mechanism of a rxn, involving less activation energy
• do not change the overall process • Do not get used up in the process. • Do not start a chemical rxn • ex. Enzymes
Catalysts lower activation energy
Inhibitor
• a substance that interferes with the action of a catalyst.
• reduces the amount of catalyst available and therefore lowers the reaction rate.
Enthalpy
• Heat of reaction• The difference in potential energy
(PE) between products and reactants represents Enthalpy
• ΔH = Hproducts - Hreactants
• table I
• In an Exothermic reaction, energy is released, products have a lower P.E. than the reactants, and the sign of ΔH is negative.
Exothermic graph. ΔH = C = amount of heat given off during rxn
• In an Endothermic reaction, energy is absorbed, products have a higher P.E. than the reactants, and the sign of ΔH is positive.
Endothermic graph. ΔH = (3) = amount of heat absorbed by rxn
Speeding up reactions:
• Most reactions, especially Endothermic reactions, will go faster with higher temperatures
• Exothermic reactions will be inhibited by very high temperatures
• Increasing surface area will increase rate of reaction
• Increasing concentrations will increase rate of reaction
Reversible Reactions
• A reversible reaction is one in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously.
• Example: • Forward reaction: 2SO2(g) + O2(g) →
2SO3(g)
• Reverse reaction: 2SO2(g) + O2(g) ← 2SO3(g)
Part 2:Equilibrium
Chemical Equilibrium
• When the rates of the forward and reverse reactions are equal, the reaction has reached a state of balance
• no net change occurs in the actual amounts of the components of the system.
• (Concentrations of Products and Reactants do not change)
Types of equilibrium
• Solution Equilibrium– dissolving is occurring at the same rate
at precipitation
• Phase Equilibrium– Vaporization occurring at the same rate
as condensation
• Reaction Equilibrium– Products are forming equal to the rate of
the reverse reaction re-forming reactants
Equilibrium Graphs…
adding catalyst to a system that was already at equilibrium?
• A catalyst will bring a system to equilibrium sooner
• The rates of the forward and reverse reactions would increase but the overall net reactions would not change.
Equilibrium Continued
Le Châtelier’s Principle Concentration
• If additional reactants (or products) are added to a reaction system at equilibrium, the eq point (point of equilibrium) will shift favoring the reaction that would relieve the stress.
Le Châtelier’s Principle Temperature
• If additional heat were added to a reaction system at equilibrium (raise the temperature of the system), the eq point will shift favoring the endothermic reaction to relieve the stress.
• An increase in temperature favors all reactions, but endothermic reactions benefit more.
Le Châtelier’s Principle Pressure
• Changing the pressure of a system only affects reactions that have components in the gaseous phase.
• If additional pressure were added to a system at equilibrium, the eq point will shift favoring the reaction that makes less gas molecules to relieve the stress.
The Haber Process• http://mail.kenton.k12.ny.us/~Bob_Ventola/chemistry/habermovie.swf
Law Of Chemical Equilibriumonly write down if you’re taking AP next year:
• ►When a reversible reaction reaches equilibrium at a given temp. the following mathematical relationship occurs
• Ex. aA + bB cC + dD • ►lower case = coefficient • ►upper case = formula • § Keq = [products] = [C]c[D]d
[reactants] [A]a[B]b
Spontaneous Reactions
• occur naturally and favors the formation of products at the specified conditions.
• produce substantial amounts of products at equilibrium and release free energy.
• fireworks
Entropy
• is a measure of the disorder of a system. (randomness) (S)
• §Recall that heat (Enthalpy) changes accompany most chemical and physical processes.
“E” words
• Enthalpy is a measure of heat energy + value = endothermic
- value = exothermic • Entropy is a measure of the disorder,
randomness, or lack of organization of a system. – ex. solid (less random) - liquid - gas
(more random) – High temp. = High entropy
Forces of the Universe
• Systems move naturally toward – a decrease in Enthalpy ( - ΔH) – an increase in Entropy ( + ΔS)
• The universe naturally makes things go to lower energy and more disorder.
Enthalpy, Entropy and Free Energy
• every chemical reaction, heat is either released or absorbed and entropy either increases or decreases.
• size and direction of enthalpy changes and entropy changes together determine whether a reaction is spontaneous
Gibb’s Free Energy Change
• The difference between energy change (DH ) and entropy change (DS ) was studied by Willard Gibb
• Gibb formula: ΔG = ΔH - TΔS • ΔG = Free energy change • ΔH = Total Heat • T = Temp in Kelvin • ΔS = Entropy
According to Gibb
• ΔG = ΔH - TΔS • If ΔG = negative value, then reaction
is spontaneous • If ΔG = positive value, then reaction
is Non-spontaneous • Zero ΔG means reactions are at
equilibrium.
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