heterogenous equilibria

Chemical equilibrium

Dynamic equilibrium where forward and reverse reactions occur at equal rates

Dynamic equilibrium

State where rate of forward reaction equals rate of reverse reaction

Law of mass action

At constant temperature the equilibrium constant equals the ratio of product concentrations to reactant concentrations each raised to stoichiometric coefficients

Equilibrium constant K

Experimentally determined constant that depends on temperature and equation form

Unit of equilibrium constant

Given as mol per dm3

Dependence of K on equation

Value of K changes if the balanced equation is altered

Extent of reaction

Tendency of a reaction to occur

K greater than 1

Equilibrium favors formation of products

K less than 1

Equilibrium favors reactants

Reaction speed vs K

Equilibrium constant gives no information about reaction rate

Homogeneous equilibrium

Equilibrium where reactants and products are in the same phase

Heterogeneous equilibrium

Equilibrium where reactants and products are in different phases

Pure solids in heterogeneous equilibrium

Their concentrations are constant and taken as 1

Pure liquids in heterogeneous equilibrium

Their concentrations do not affect equilibrium constant

Equilibrium constant in heterogeneous systems

Does not depend on amounts of pure solids or liquids

Solubility equilibrium

Heterogeneous equilibrium between solid and dissolved ions

Solubility

Concentration of saturated solution at a given temperature

Units of solubility

Gram per dm3 or mol per dm3

Molar solubility

Solubility expressed in mol per dm3

Solubility of chlorides

Mostly soluble except silver mercury(I) and lead chlorides

Solubility of bromides

Mostly soluble except silver mercury and lead bromides

Solubility of iodides

Mostly soluble except silver mercury and lead iodides

Solubility of sulfates

Mostly soluble except calcium strontium barium lead mercury and silver sulfates

Nitrates chlorates acetates

Always soluble in water

Solubility of sulfides

Mostly insoluble except alkali alkaline earth and ammonium sulfides

Solubility of hydroxides

Mostly insoluble except alkali hydroxides and barium strontium calcium hydroxides

Solubility of carbonates

Mostly insoluble except alkali metal and ammonium carbonates

Solubility of sulphites

Mostly insoluble except alkali metal and ammonium sulphites

Solubility of phosphates

Mostly insoluble except alkali metal and ammonium phosphates

Factors affecting solubility

Temperature common ion pH and complex formation

Temperature effect on solubility

Solubility usually increases with temperature

Endothermic dissolution

Solubility increases with temperature

Exothermic dissolution

Solubility decreases with temperature

Constant solubility

Solubility nearly independent of temperature

Recrystallization suitability

Compounds with steep solubility curves

Common ion effect

Presence of common ion reduces solubility

Le Chatelier principle application

System shifts to oppose added common ion

Common ion example lead chloride

Added chloride ions cause precipitation of lead chloride

pH effect on solubility basic anions

Solubility increases as pH decreases

pH effect on solubility neutral anions

Solubility unaffected by pH

Complex formation effect

Formation of complex can increase or decrease solubility

Iodine potassium iodide example

Formation of soluble triiodide complex increases solubility

Tetracycline calcium example

Complex formation decreases solubility

Solubility product Ksp

Equilibrium constant for slightly soluble ionic compounds

Definition of Ksp

Product of ion concentrations raised to stoichiometric powers

Ksp and solids

Concentration of solid omitted from expression

Meaning of high Ksp

Higher Ksp means greater solubility

Precipitation condition

Ksp exceeded leads to precipitate formation

Comparison of lead halides solubility

PbCl2 most soluble PbBr2 less PbI2 least soluble

Relation between solubility and Ksp

Depends on stoichiometry of salt

AX type salt solubility

Ksp equals x squared

AX3 type salt solubility

Ksp equals 27 x to the fourth power

Ion product IP

Product of initial ion concentrations

IP greater than Ksp

Precipitation occurs

IP less than Ksp

Solution is unsaturated

IP equals Ksp

Solution is saturated and at equilibrium

Precipitation prediction method

Compare IP with Ksp

Dilution calculation

Use c1V1 equals c2V2

BaSO4 precipitation example

Q greater than Ksp indicates precipitate formation