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Equilibrium

Chapter 17
by

Alex Pittman

on 26 March 2013

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Transcript of Equilibrium

Equilibrium How Chemical Reactions Occur Chemical Equilibrium: A Dynamic Condition Summary Chemical reactions occur under specific circumstances.
This idea is called the collision model. During Equilibrium it appears that everything has stopped. We can predict the effects of changes in concentration, pressure, and temperature on a system at equilibrium by using Le Chatlier’s principle
It states that when a change is imposed on a system at equilibrium, the position of the equilibrium shifts in a direction that tends to reduce the effect of that change. Chemical reactions are described by the collision model Molecules collide to cause reaction. The energy of the collision must be higher than the activation energy required for a reaction to occur. They must also collide in the same orientation for a reaction to occur. The collision model also explains why reaction speeds vary as temperature changes. Conditions that Affect Reaction Rates The activation energy is the minimum amount of energy required for a reaction to occur. As temperature increases, the speed of molecules increases, causing it to react faster. Catalyst, such as enzymes, cause reactions to occur faster. Catalysts are present in the human body, as well as in car exhaust systems. Le Chatelier's Principle The Equilibrium Condition Equilibrium is the exact balancing of two opposite processes. Chemical equilibrium is a dynamic state where the concentrations of all reactants and products remain constant. Chemical reactions reach equilibrium when they are moving forward and reverse at the same rate. Applications Involving the Equilibrium Constant Knowing the equilibrium constant tells us to do many things
The size of K tells us the tendency of the reaction to occur
A value of K much larger than one means that at equilibrium, the reaction system will consist of mostly products
Small value of K means that the system consists largely of reactants Solubility Equilibria That is not the case, on the molecular level the compounds move back and forth equally. Ksp is the solubility product constant, or simply the solublility product. In other words, the rate of the forward reaction equals the rate of the reverse reaction The Equilibrium Constant:
An Introduction Cato Maximilian Guldberg and Peter Waage proposed the Law of Chemical Equilibrium
aA + bB cC + dD ->
<- A, B, C, & D represent chemical species
a, b, c, & d are their coefficients in a balanced equation The law of chemical equilibrium is represented by the Equilibrium Expression
K = [C]^c[D]^d / [A]^a[B]^b The brackets represent the concentrations of the chemical species at equilibrium.
K is a constant known as the equilibrium constant The Law of Chemical Equilibrium is based on experimental observations. Each set of equilibrium positions is called an equilibrium position. There can only be one equilibrium constant for a certain system at a certain temperature On the other hand, there are an infinite number of equilibrium positions. Heterogeneous Equilibria Homogeneous Equilibria - equilibria in which all substances are in the same state

Heterogeneous Equilibria - equilibria in which the substances have more than one state In heterogeneous equilibria, you do not include solids and liquids in the equilibrium expression. For example:
The Equilibrium Expression for
CaCo3 (s) CaO (s) + CO2 (g)
Would be
K = [CO2] ->
<- Which says that molecules must collide with the same orientation and that Ecollision must be greater than Eactivation Several factors can be changed to affect the speed of a reaction. Catalysts can be added, or temperature and/or pressure can be altered, which affects the rate at which a reaction occurs. Equilibrium is eventually achieved when a reaction occurs in a closed container. The law of chemical equilibrium states that in a reaction where aA + bB cC + dD -> <- The equilibrium expression is K=([C]^c*[D]^d)/([A]^a*[B]^b) Where K is the equilibrium constant. Le Chatelier's Principle predicts the effects of changes in reaction conditions on a reaction.
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