Chapter 14

Chemical Equilibruim

Objectives

• Describe chemical equilibrium

• Write an equilibrium constant expression

• Calculate the equilibrium constant for a reaction

• Relate the equilibrium constant to the position of a reaction

Chemical Equilibrium

• Point in a chemical reaction where the forward and reverse reactions happen at the same rate.

• There does not need to be equal amounts of reactants and products

– There can be a large amount of product or a large amount of reactant at equilibrium

• The rates just need to be equal

Chemical Equilibrium

• Equilibrium is indicted by double arrows in reactions– Forward reaction points right

• Products– Reverse reaction points left

• Reactants• The right is still products, the left is still

reactants

Chemical Equilibrium

• Equilibrium is a dynamic process

– Continually happening

• Reaction will continue even though it will appear to have stopped

• Similar to a person on a treadmill

Example

• Consider the reaction

A + B C• Initially only A & B are present

Example

• Consider the reaction

A + B C• A & B begin to disappear, C begins to appear

Example

• Consider the reaction

A + B C• A & B begin to disappear, C begins to appear

Example

• Consider the reaction

A + B C• A & B begin to disappear, C begins to appear

Example

• Consider the reaction

A + B C• A & B begin to disappear, C begins to appear

Example

• Consider the reaction

A + B C• Eventually equilibrium is reached

Equilibrium Constant Expressions

• A mathematical equation that relates the concentration of products to the concentration of reactants for a reaction at equilibrium

• The equilibrium constant is Keq• K for Konstant, eq for equilibrium• General form of the expression is

concentration of products over concentration of reactants– Pressure can also be used

The Equilibrium Constant Expression

• For the reaction

jA + kB lC + mD• Where j, k, l, & m are coefficients and • A, B, C, and D are chemicals• The equilibrium constant expression is

kj

ml

BA

DCKeq

][][

][][

Equilibrium Constant Expressions

• A species in brackets indicates concentration

• Only solutions and gases are included in equilibrium expressions

• If you have a solid or liquid leave it out

– Just think of it as a 1 in the expression

Example• Write the equilibrium expression for the

following reaction. HF(aq) H+ + F-

Example• Write the equilibrium expression for the

following reaction. 2NO(g) N2(g) + O2(g)

Example• Write the equilibrium expression for the

following reaction. 2C(s) + O2(g) 2CO(g)

The Equilibrium Constant

• A value that relates the position of a reaction at equilibrium

• Calculated by inserting values into the equilibrium constant expression

• Keq does NOT have units• Keq > 1

– Product favored• Keq <1

– Reactant favored

The Equilibrium Constant & Temp

• Equilibrium constants are constant for a given temperature

– If the temp. changes so does the constant

– Change depends on whether the reaction is exothermic or endothermic

• More on that later

Example• Calculate the equilibrium constant for the

following reaction and equilibrium concentrations [HF]= 0.12M [H+]= 9.3x10-3M [F-]= 9.3x10-3M

HF(aq) H+ + F-

Example• Calculate the equilibrium concentration of

NO given the reaction, Keq, and equilibrium concentrations. [N2] = 0.37M [O2] = 0.45M Keq = 2.3x10-9 2NO(g) N2(g) + O2(g)

Homework

• p.573 #26,32,35,39,52,55,61

Objectives

• Explain LeChatliers Principle

• Predict equilibrium shifts

• Describe solubility equilibrium

• Write Ksp expressions

• Perform calculations with Ksp

My Equilibrium is Stressed

• A reaction at equilibrium can be stressed out by changing reaction conditions– There are three stresses for reactions

• Concentration• Pressure • Temperature

• Q) If you are stressed what do you do?• A) Work to relieve the stress!

– Same goes for reactions

LeChatlier’s Principle

• When a stress is applied to a reaction at equilibrium the reaction will shift to relieve the stress placed upon it.

Concentration

• A reaction is stressed when the concentration of a reactant or product is increased OR decreased

– Add a chemical to increase

– Remove a chemical to decrease

• Reactions shift away from increased concentrations

• Reactions shift toward decreased concentrations

Example

• Consider the reaction

2SO2(g) + O2(g) 2SO3(g)

• How will the reaction shift with the following– Increase the amount of SO2

• Right– Increase the amount of SO3

• Left– Decrease the amount of O2

• Left

Pressure

• Pressure can increased

• By decreasing volume

• Pressure can be decreased

• By increasing volume

• Increased pressure shifts reactions away from side with more gas molecules

• Decreased pressure shifts reactions toward side with more gas molecules

Example

• Consider the reaction

2SO2(g) + O2(g) 2SO3(g)

• How will the reaction shift with the following

– Increase the pressure in the container

• Right

– Decrease the pressure in the container

• Left

Temperature

• Direction of shift depends of whether the reaction is exothermic or endothermic

– How do you know

• Exothermic reactions release heat

–It is a product

• Endothermic reactions absorb heat

–It is a reactant

What Type of Thermic Is It?

• Exothermic

A + B C + 49kJ

A + B C ΔH = -49kJ

• ΔH means enthalpy (energy change)

• Endothermic

A + B + 49kJ C

A + B C ΔH = 49kJ

Temperature Change

• Think of the temp as a concentration of heat

– Increasing the temp. will shift the reaction away from the heat

– Decreasing the temp. will shift the reaction toward the heat.

• Same as concentration shift

• Change in temp. is the only stress that changes the equilibrium constant.

Example

• Consider the reaction

2SO2(g) + O2(g) 2SO3(g) + 197kJ

• How will the reaction shift with the following

– Increase the temperature

• Left

– Decrease the temperature

• Right

Common Ion Effect

• The addition of an ion that is present in the equilibrium

HF H+ + F-

Adding a H+ or F- shifts the reaction left

• How do you add them?

– H+ can come from any acid

– F- could be from NaF, KF, CaF2

Addition of Other Stuff

• Adding species not related to the equilibrium reaction have no effect on the reaction

– NO SHIFT

• For Example

–Catalysts

–Inert Gases

Solubility Equilibrium

• Dissolving is an equilibrium process

• Chemicals that we say are insoluble are actually very slightly soluble

– Meaning only a small amount of the solute dissolves

– Concentrations are very small

• 10-5M to 10-20M are common

Solubility Equilibrium

• Iron (II) Sulfide is “insoluble” by solubility rules

• However, it still dissolved to a small extent

• Consider the reaction

FeS(s) Fe+2 + S-2

• Since dissolving is an equilibrium process we can write an equilibrium expression

• Keq = [Fe+2] [S-2]

• FeS is omitted because it is a solid

Solubility Product Constant

• The previous equilibriums expression is a “special” type of equilibrium expression

• Ksp – solubility product constant

• Ksp = [Fe+2] [S-2]

– Only used for slightly soluble salts

– Never includes the reactants

– Ksp’s are very small

– Check out the table 14.1 page 568

Ksp Values

• Lead (II) Chloride 2.4X10-4

• Strontium Carbonate 3.8X10-9

• Nickel (II) Hydroxide 5.5X10-16

• Copper (II) Hydroxide 2.2X10-20

• Cadmium Sulfide 8.0X10-28

• Silver Sulfide 6.0X10-51

• And my personal favorite

• Bismuth Sulfide 1.6x10-72

Ksp Values

• Q) What does a small Ksp value mean?

• A) Low concentration of ions

• Small solubility

• Q) Is the compound with the smallest Ksp the least soluble?

• A) Not necessarily

– There are different numbers of ions that changes the expression

Solubility of Salts

• We can easily compare solubility for salts that have the same # of ions

• When there are the same # of ions the salt with the smallest Ksp is least soluble

• Which salt is least soluble, most soluble AgCl, FeS,

– FeS 6.0x10-19 Least soluble

– AgCl 1.6x10-10 Most soluble

– BaSO4 1.1x10-10 Middle

Calculating Solubility

• We can calculate solubility of salts and the concentration of the ions in solutions from Ksp

• Deal with saturated solutions

– The salt has dissolved as much as it can

• Ksp has been reached

• In saturated solutions the concentrations of the ions are related to the mole ratio

Example• Calculate the solubility of silver chloride

Example• Calculate the solubility of

mercury (II) sulfide

Example• Calculate the solubility of calcium phosphate

Example• A saturated solution of Iron (III) Hydroxide

has a concentration of Fe+3 of 1.8x10-8. What is the Ksp of Iron (III) Hydroxide?

Common Ion Revisited

• Anyone ever have a barium shake for a scan?

• Contains barium sulfate.

– Barium ions are very toxic to the body

– Treated as calcium

• How can you have fewer barium ions in solution?

• Add Sodium Sulfate

Homework

• p.574 #71,77,89,95,97,102