AP Chemistry
th
Zumdahl Notes, 9 ed.
A Brief Collection of notes, Chapter 13
Feel free to open these files and annotate as you feel the need…this is
for your success.
Chemical Equilibrium: the Equilibrium Condition
• What is chemical equilibrium? The state where the concentrations of all reactants
and products remain constant…this does not mean that reactants are not forming
products, nor products reverting back to reactants, but, rather, that the rates of
these competing reactions are equal.
• We must realize that as the reaction heads toward this state of equilibrium that
the reactions will also slow down, asymptotically toward a flat line which is the
equilibrium concentration for that component
• Reality check: the equilibrium does not mean equal amounts of reactant and
product present…I think this is intuitive but mention it anyways
• Another reality check: realize that the equilibrium can be shifted by a change in
temperature, pressure (if gases involved), catalysis, etc., as postulated by the
scientist Le Chatelier
Chemical Equilibrium: the equilibrium constant
• The development of the mathematical expression for the equilibrium condition,
resulting in keq, is empirical, that is, based on experimental evidence
• aA + bB ↔ cC + dD
k= [C]c [D]d/ [A]a [B]b
• Yes, the lower case letters, the exponents, correspond to the coefficients
• Realization: this expression is reversible, for when you reverse the reaction.
• Reality check: equilibrium position is the location for an individual set of
concentrations…this will allow for calculations; keq is a constant for a given
reaction…the concentrations will be dependent upon the initial conditions
• This is referred to as the Law of Mass Action, proposed in 1864…wow!
Chemical Equilibriun: Equilibrium expressions involving pressures
• Remembering the ideal gas law (PV=nRT), we can rearrange into P=(n/V)RT, which
translates into P=CRT, where C represents molar concentration of the gas
• This allows us to write a Kp expression using the partial pressures of the gases
present in the reaction
• Exponents are still derived from coefficients from balanced reactions
• Utility of Kp: K=KpRT
Chemical Equilibrium: heterogeneous equilibria
• Realization: the position of a heterogeneous equilibrium does not
depend on the amount of pure solids or liquids present, so their
presence is removed from any mathematical expressions (yes, this
does simplify things somewhat)…this does not include solutions…oh
well!
Chemical Equilibrium: applications of K
• Knowing the equilibrium constant gives us some useful information
• Does the reaction tend to occur?
• Do the condition represent equilibrium?
• What is the equilibrium position for the given set of initial concentrations?
One thing I will say here, and let you digest it: ICE chart!
• The extent of a reaction: how large is K? small number means little product/s,
large number indicates lots of product/s.
• Reaction quotient: expression utilizing initial concentrations, rather than those at
equilibrium
• If at K, system at equilibrium
• If greater than K, reaction will move from products to reactants, to attain equilibrium
• If less than K, more reactants will move to form products, to attain equilibrium
Chemical Equilibrium: applications of K, continued
• Calculating Equilibrium Pressures and Concentrations
• As usual, rearrange expressions so as to solve for unknowns
• Again, think ICE…helps you to see concentrations accurately
• Remember to use properly balanced equations…or your exponents will be wrong!
Chemical Equilibrium: solving equilibrium problems
• Strategic steps (pg. 628):
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Write the balanced equation for the reaction
Write the equilibrium expression using the law of mass action (our usual expression)
List the initial concentrations
Calculate Q (reaction quotient) for these conditions, and determine the direction of the shift
to equilibrium
Define the change needed to reach equilibrium, and define the equilibrium concentrations by
applying the change to the initial concentrations (plugging into our ICE chart)
Substitute the equilibrium concentrations into the equilibrium expression, solving for
unknown/s (pulling out of our ICE charts)
Check your calculated equilibrium concentrations by making sure they give the correct value
of K
Can be used either for concentrations or pressure situations
Chemical Equilibrium: Le Chatelier’s Principle
• Stated: if a change is imposed on a system at equilibrium, the position of the
equilibrium will shift in a direction that tends to reduce that change
• In other words, if a stress is applied, the system will react so as to relieve that stress
• Application of Le Chatelier’s principle
• Changes in concentration from a position of equilibrium…how will system react?
• Changes in pressure from a position of equilibrium…how will system react?
• Add/remove product or reactant
• Add inert gas (one not involved in reaction)…no net effect besides changing total pressure of system, but no
change in position of equilibrium
• Change volume of reaction system…realize that this changes concentration!
• Changes in temperature
• Yes, K is dependent upon temperature, so this is something we cannot ignore!
• Typically, you will have either a series of values given, and the knowledge of whether or not one of the reaction
directions is exothermic/endothermic, so the effects should be something you can deduce