Study Guides
Big Picture
Many reactions are reversible; that is, the products can react to form the reactants. When this happens, the two
reactions (reactants
products and products
reactants) will occur at the same time and will ultimately happen
at the same rate. This is called equilibrium. Equilibrium can be manipulated by various factors and can help us affect
the outcome of a reaction.
Key Terms
Chemistry
Equilibrium
Equilibrium: When the forward reaction rate and reverse reaction rate are equal. This does not mean that no reactions
are happening but instead that the two reactions are occurring at the same rate, so the concentrations of reactants
and products are not changing.
Forward Reaction: The reaction from reactants to products.
Reverse Reaction: The reaction from products to reactants.
Reaction Rate: The change in concentration of reactants over time; the speed at which a reaction occurs.
Equilibrium Constant: Describes the concentration of reactants and products when equilibrium is reached.
Le Châtelier’s Principle: A reaction at equilibrium will move in the direction to relieve stress.
Ions
Equilibrium
State
A reaction that is at equilibrium is shown by drawing double arrows ( ) between the products and the reactants.
This means that the reaction rate of the forward reaction is equal to the reaction rate of the reverse reaction.
The reaction is reversible.
• The forward and reverse reactions are taking place, but there is no change for the system as a whole.
• Almost all reactions are reversible to some degree under the right conditions.
• Although the reaction is at equilibrium, usually one of reactions (forward or reverse) is favored (produces a greater
concentration of either the reactants or products).
and
are not the same - one shows that a reaction is in equilibrium and the other shows resonance.
Equilibrium Constant
Kc is the symbol for the swequilibrium constant. For the general reaction
,
The general rule is to write the coefficient as an exponent. The coefficients must be from the balanced equation.
• EXCEPTION: Solids and liquids that are not dissolved in a solution are not part of the equilibrium constant because
increasing or decreasing the concentration of these substances do not change, so they will not affect the reaction
rate.
The concentrations used to calculate the equilibrium constant must be the equilibrium concentrations!
Let’s say we have a reaction
.
• The
coefficients become exponents. D is not included because it is a solid. A and B are included because they are
dissolved in water.
A reaction will typically favor one side - either the reactants or the products.
• A
reaction that favors the products means that more products are present than reactants and the equilibrium
constant is greater than 1.
• A reaction that favors the reactants means that more reactants are present and the constant is less than one.
• This must be determined through experimental trials.
The units for equilibrium constants depend on the particular equilibrium equation. Remember that [] means
concentration.
This guide was created by Steven Lai, Rory Runser, and Jin Yu. To learn more
about the student authors, visit http://www.ck12.org/about/about-us/team/
interns.
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v1.11.4.2011
Disclaimer: this study guide was not created to replace
your textbook and is for classroom or individual use only.
• Kc for this reaction would be
Chemistry
Equilibrium
cont .
Le Châtelier’s Principle
Le Châtelier’s principle describes what happens when you alter various conditions for a reaction at equilibrium.
Equilibrium is dependent on temperature, pressure, volume, concentration, and other factors. One reaction, either
the forward or reverse, will occur at a faster reaction rate when these conditions are changed, and more products
or reactants will be produced and be present at equilibrium. The reaction will either “shift forward” (produce more
products) or “shift backward” (produce more reactants).
• Adding
reactants: shifts forward. Having a greater [reactants] makes it more likely that a collision will occur and
that the reaction will occur.
• Removing reactants: shifts backward.
• Increasing pressure: shifts toward side with fewer moles of gas. At greater pressure, gases are more likely to collide
and react. Thus, the side with more gases will react more to produce the compounds on the side with less gas.
• Decreasing pressure: shifts toward side with more moles of gas.
• Changing
volume: volume doesn’t directly change equilibrium, but in reactions with gases, increasing volume
decreases pressure and decreasing volume increases pressure.
• Raising
temperature: shifts toward products if reaction is exothermic and toward reactants if it is endothermic.
You can think of the heat consumed as a reactant (endothermic) and released as a product (exothermic). Thus,
increasing temperature (adding heat) favors the side with less heat.
• Decreasing temperature: shifts toward reactants if reaction is exothermic and toward products if it is endothermic.
• Catalyst: affects the reaction rate, but it does not change equilibrium constant.
• Adding an inert gas, such as a noble gas: does not change equilibrium constant.
Notes
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Find the Kc of a Reaction
Often, you will be given a chemical reaction and the concentrations of reactants and products at equilibrium. The
problem will ask you to find a numerical value for the equilibrium constant.
Example: The equation
is reversible and at equilibrium in a 2.00-liter container. In this
container, there are 2.0 moles of oxygen, 1.5 moles of sulfur dioxide, and 0.25 moles of sulfur trioxide. Find the
equilibrium constant.
1.Write the equilibrium constant!
- remember exponents!
2.Find the concentrations of each substance.
[O2] = 2.0 mol/2.00 L = 1.0 M
[SO2] = 1.5 mol/2.00 L = 0.75 M
[SO3] = 0.25 mol/2.00 L = 0.13 M
3.Plug and chug.
Applying Le Châtelier’s Principle
A very common problem gives an equation and asks to determine how changes to the setup will affect the equilibrium constant.
Example:
The Haber Process is the formation of ammonia from hydrogen and nitrogen, given
, where
∆H = 92.22 kJ/mol. Determine how the following will affect its equilibrium rate if it is initially in equilibrium in a closed container.
Before you do anything else, write the equilibrium constant.
• Increasing temperature
• Since this reaction is exothermic (∆H < 0) increasing the temperature favors the reactants, so the reaction shifts
backward.
• Decreasing volume
• Decreasing volume increases pressure. The left side of the equation has 4 moles of gas and the right has 2, so
increasing the pressure will favor the right side and shift the reaction forward.
• Adding 1 mole of argon to the mix.
• Inert gases do not change equilibrium constants. Thus, no change occurs.
• Increasing the concentration of hydrogen gas.
• Increasing [reactants] favors the products, so the reaction will shift forward.
ICE Chart
Given a equilibrium constant and an initial molarity, there will be questions asking you about the equilibrium
concentrations. To solve this problem, you should use an ICE (Initial, Change, Equilibrium) chart. This chart lets you
figure out the ending molarities.
Example:
A typical ICE chart
Initial
Change
Equilibrium
A
[Initial]
-x
[Initial - x]
B
0
+x
x
C
0
+x
x
Once you have the end, plug it into your expression.
You can solve for x, but you need to solve a quadratic equation. Because sometimes the change is quite small, chemists
came up with the 5% rule. If the change is less than 5% of the initial molarity, you can just disregard the change since
it is negligible. So the simplified equation is
. Now, you can easily solve for x. Once you have x, you can
solve for whatever you needed to find.
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Chemistry
Equilibrium Problem Guide