Chapter 14: Chemical Equilibrium

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Some Basics

  • A + B => C + D (forward reaction)
  • C + D =>A + B (reverse reaction)
  • Initially there is only A and B so only the forward reaction is possible
  • As C and D build up, the reveres reactions speeds up while the forward reaction slows down.
  • Eventually the rates become equal.

The Concept of Equilibrium and the Equilibrium Constant

  • Most chemical reactions are reversible to some extent.
  • As soon as some product molecules are formed, the reverse process begins to take place and reactant molecules are formed from product molecules
  • Chemical equilibrium- achieved when the rate of the forward and reverse reactions are equal and he concentrations for the reactants and products remain constant.
  • Physical equilibrium- the changes that occur are physical processes
  • Double arrows (<=> means reaction is reversible)

The Equilibrium Constant

  • Gas concentrations are expressed in morality, which can be calculated from the number of moles of the gases present initially and at equilibrium and volume of the flask in liters
  • Where K is a constant for the equilibrium reaction 3.jpg, then4.jpg which is also known as the equilibrium constant.
  • Law of mass action- for a reversible reaction at equilibrium and a constant temperature, a certain ratio of reactant and product concentrations as a constant value, K (equilibrium constant)
  • The equilibrium constant is defined by a quotient, the numerator is obtained by multiplying together the equilibrium concentrations of the products and then each is raised to a power equal to the stoichiometric coefficient in the balanced equation
  • The denominator is obtained by multiplying the equilibrium concentrations of the reactant and raising each to the power that is equal to the coefficient.
  • Magnitude of equilibrium constant tells whether an equilibrium reactions favors the products or the reactants
  • If K is much greater than 1, the equilibrium will favor the right, also known as the products
  • If K is much smaller than 1, the equilibrium will favor the left, also known as the reactants.

Writing Equilibrium Constant Expressions

  • There are many ways to express the concentrations of reactants and products because they are not always in the same phase
  • It is also a general rule of thumb to not use units for the equilibrium constant

Homogeneous Equilibria

  • Homogeneous equilibrium- reactions in which all reacting species are in the same phase
homogenous.JPG (
  • Kp tells that equilibrium concentrations are expressed in terms of pressure
  • Usually Kp is not equal to Kc because partial pressures of reactions and products are not equal to their concentrations expressed in moles per liter.
1.JPG (Zumdahl, 4th Edition.)

Heterogeneous Equilibria

  • Heterogeneous equilibrium- reversible reaction involving reactants and products that are in different phases.
  • Basically, the products and reactants are in different states of matter.
  • “Concentration” of a solid is an intensive property and does not depend on how much of the substance is present so when writing your equilibrium equation it is not necessary to include solids.
  • Liquids also are not included in equilibrium equations.
hetero.JPG (

Multiple Equilibria

  • Product molecules in one equilibrium system are involved in a second equilibrium process
  • If a reaction can be expressed as the sum of two or more reactions, the Equilibrium constant for the overall reaction is given by the product of the equilibrium constants of the individual reactions
  • An example would be the ionization of diprotic acids in aqueous solution.
1.gif(Chang, 8th edition)

The Form of K and the Equilibrium Equation

  • There are two important rules for writing equilibrium constants
  • 1. When the equation for a reversible reaction is written in the opposite direction, the equilibrium constant becomes the reciprocal of the original equilibrium constant.
  • 2. The value of K also depends on how the equilibrium equation is balanced.
2.gif (Chang, 8th edition)

The Relationship between Chemical Kinetics and Chemical Equilibrium

  • The equilibrium constant helps us to predict the direction in which a reaction mixture will proceed to achieve equilibrium and to calculate the concentrations of reactant and products once equilibrium has been reached

Predicting the Direction of a Reaction

  • Reaction quotient (Qc)- for reactions that have not reached equilibrium, instead of the equilibrium constant by substituting he initial concentrations into the equilibrium constant expression
  • We compare the values of Qc and Kc to determine the direction in which the net reaction will achieve equilibrium
  • Qc < Kc : ratio of products to reactant is too small. To reach equilibrium reactants must be converted to products. System proceeds from left to right to reach equilibrium
  • Qc = Kc : The concentrations are equilibrium concentrations. The system is at equilibrium
  • Qc > Kc : Ratio of concentration is too large. To reach equilibrium, products must be converted to reactants. System proceeds from right to left to reach equilibrium.

Factors That Affect Chemical Equilibrium

  • Changes in experimental conditions may disturb the balance and shift the equilibrium positions so that more or less of the desired products is formed.
  • If it shifts to the right, the net reaction is now from left to right
  • Variables that can be controlled experimentally are concentration, pressure, volume, and temperature

Le Chatelier’s Principle

  • Helps predict the direction in which an equilibrium reaction will move when a change in concentration, pressure, volume, or temperature occurs
  • Le Chatelier’s principle- if an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially offset as the system reaches a new equilibrium position
  • “stress” means a change in concentrations, pressure, volume, or temperature that removes the system from the equilibrium state.
  • We use Le Chatelier’s principle to assess the effects of the changes

Changes in Concentration

  • All reactants and products are present in the reacting system at equilibrium
  • Increasing the concentrations of the prod cuts shifts the equilibrium to the left, and decreasing the concentration of the products shifts the equilibrium to the right
  • The change in equilibrium is predicted by Le Chatelier's principle

Changes in Volume and Pressure

  • Usually, changes in pressure do not affect the concentrations of reacting species in condensed phases because gas and liquids are pretty much impressible
  • Gases are greatly affected by changes in pressure though
  • PV=nRT and P= (n/V) RT
  • P and V are related inversely (greater the pressure, smaller the volume, and vice versa)
  • Term (n/v) is the concentration of the gas in mol/L (varies directly)
  • An increase in pressure (decrease in volume) favors the net reaction that decreases the total number of moles of gases (the reverse reaction)
  • decrease in pressure (increase in volume) favors the net reaction that increases the total number of moles of gases (the forward reaction)
  • For reactions where there is no change in the number of moles of gases, a pressure (or volume) change has no effect on the position of equilibrium
  • It’s possible to change the pressure of a system without changing its volume

Changes in Temperature

  • A change in concentration, pressure, or volume does not change the value of the equilibrium constant
  • Only a change in temperature can alter the equilibrium constant
  • At equilibrium at a certain temperature, the heat effect is zero because there is no net reaction
  • A rise in temperature “adds” heat to the system and a drop in temperature ‘removes” heat from the system
  • Like any change in concentration, pressure, or volume, the system shifts to reduce the effect of the change
  • Temperature increase= endothermic direction (left to right of the equilibrium equation)
  • Temperature decrease = exothermic direction ( from right to left of the equilibrium equation
  • The equilibrium constant increases when the system is heated and decreases when the system Is cooled

The Effect of Catalyst

  • Catalysts enhance the rate of a reaction by lowering the reaction’s activation energy
  • Catalysts lowers activation of the forward and reverse reactions to the same extent
  • Adding a catalyst not at equilibrium will simply cause the mixture to reach equilibrium sooner
  • The same equilibrium mixture could be obtained without the catalyst, but might take longer

Summary of the Factors That May Affect the Equilibrium Positions

  • Total of four ways to affect a reacting system at equilibrium
  • Only a change in temperature changes the value of the equilibrium constant
  • Changes in concentration, pressure, and volume can alter the equilibrium concentrations of the reacting mixture, but cannot change the equilibrium constant
  • A catalyst can speed up the process but has no effect on the equilibrium constant or on the equilibrium concentrations of the reacting species


Here are problems of all different types that you have encountered in this chapter... (problems from Chang Chemistry, 8ed.)

Now that you know all the background information, let's see it in action!!!!

Some bonus info
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Majority of the information came from Chang Chemistry, 8ed