Solubility Rules

Salts are all generally soluble to some extent. However, the solubility of some salts is so small that they are described as being insoluble. Other salts are described as being soluble or partially soluble. A general guide for describing the solubility of a salt is outlined below:

  • Soluble salts dissolve at a rate of 10g/L or more
  • Insoluble salts dissolve at a rate of 1g/L or less
  • Partially soluble salts dissolve at a rate between 1-10g/L

There is a value that describes the solubility of partially and insoluble salts, Ksp. This concept is explored further in a later section.

When analysing the solubility of salts, recall that a solid ionic compound will dissociate in water into a solution of its ions:

NaCl(s) → Na+(aq) + Cl–(aq)

Precipitation reactions occur when two soluble solutions are mixed and there are ions present that combine to form an insoluble product, the precipitate. In the diagram below, the first two beakers contain soluble ionic compounds which have formed solutions in water. The third beaker illustrates how the yellow and green ions have formed a precipitate and the blue and pink ions are free to move around the solution.

Solubility rules were developed experimentally and can be used to help predict the formation of precipitates and the identity of spectator ions.

Understanding how solubility rules apply to the chemistry of ionic solutions and reactions is not an overly complex idea. However, the scope and number of precipitates and rules that could be understood is very large. We will focus on some of the main ions that may be encountered in this course and some general tips that can be used to help remember these.

Solubility Rules

Ideas for Remembering the Solubility Rules

Below are some ideas that you may be able to use to help you identify precipitates and spectator ions in reactions. Whilst it is not a complete summary of all possible rules, it will assist you in understanding most of the reactions you may encounter:

NAGSAG LMS CASTROBAR

NAGSAG:

  • nitrates (NO3–) are always soluble
  • acetates (CH3COO–) are always soluble
  • group 1 are always soluble
  • sulfates (SO42-) are always soluble ◊ ©
  • ammonium (NH4+) is always soluble
  • group VII are always soluble ◊

Exceptions – LMS and CASTROBAR form a list of exceptions to the rules above:

LMS (noted as ◊ above)

  • Lead (Pb2+), Mercury (Hg2+), Silver (Ag+)

CASTROBAR (noted as © above)

  • Calcium (Ca2+), Strontium (Sr2+), Barium (Ba2+)

Ionic Equations

Another fundamental skill when identifying precipitates and spectator ions is writing ionic and net ionic equations.

  • An ionic equation will write any solutions as the constituent ions and precipitates as solids
  • A net ionic equation will omit the spectator ions

Potassium chloride and silver nitrate

KCl(aq) + AgNO3(aq) → KNO3(aq) + AgCl(s)

K+(aq) + Cl–(aq) + Ag+(aq) + NO3–(aq) → K+(aq) + NO3–(aq) + AgCl(s)

Cl–(aq) + Ag+(aq) →  AgCl(s)

Potassium iodide and lead nitrate

2KI(aq) + Pb(NO3)2(aq) → 2KNO3(aq) + PbI2(s)

2K+(aq) + 2I–(aq) + Pb2+(aq) + 2NO3–(aq) → 2K+(aq) + 2NO3–(aq) + PbI2(s)

2I–(aq) + Pb2+(aq) → PbI2(s)

Sodium sulfate and barium nitrate

Na2SO4(aq) + Ba(NO3)2(aq) → 2NaNO3(aq) + BaSO4(s)

2Na+(aq) + SO42-(aq) + Ba2+(aq) + 2NO3–(aq)  →  2Na+(aq) + 2NO3–(aq) + BaSO4(s)

SO42-(aq) + Ba2+(aq)   →  BaSO4(s)

Example 1:

Use the table of solubility rules to determine if the following ionic compounds are soluble or insoluble?

a) CaSO4

b) KOH

c) Magnesium phosphate

Answers:

a) insoluble  b)  soluble  c) insoluble

Example 2.

Name the precipitate that forms in the following reaction:

Na2CO3(aq) + Mg(NO3)2(aq)

Answer:

MgCO3(s)