EXPERIMENT 15: AQUEOUS SOLUBILITIES OF IONIC COMPOUNDS PURPOSE To determine which cation-anion combinations form water-soluble or water-insoluble salts. To formulate general solubility rules from the results of this determination.

BACKGROUND A crystal of a salt is a solid that does not have discrete molecules, but rather consists of a very large number of cations and anions packed closely together. This array of positive and negative ions is called a crystal lattice. The attractions between these positive and negative charges are called ionic bonds and are the forces which hold the crystal together. These attractive forces operate only when the ions are very close together. When a salt crystal dissolves, its ionic bonds are broken: the cation and anion become separated from one another and each becomes surrounded by water molecules. These hydrated ions move about independently in the solution. The solubility of an ionic compound, therefore, depends on the strength of its ionic bonds: the stronger the bonds, the lower the solubility. The strength of an ionic bond depends on the charge density of the cation and the anion. Just as density is a ratio of mass to volume, charge density is a ratio of charge to volume. An ion has a high charge density if its charge is large and its diameter small, and a low charge density if its charge is small and its diameter large. In each of the pairs below, the charge on the two ions is the same. In each case, the first ion is smaller and therefore has a higher charge density: –

–

F (fluoride ion) and Cl (chloride ion) 2–

2–

O (oxide ion) and S (sulfide ion) +

+

Na (sodium ion) and Rb (rubidium ion) Ca

2+

(calcium ion) and Sr

2+

(strontium ion).

In the series below, the charge density increases from left to right for two reasons. +

Na (sodium ion), Mg

2+

(magnesium ion), Al

3+

(aluminum ion)

The charge increases and the ionic size decreases along the series.

An ionic bond is a coulombic force. It is proportional to the charges on the species involved and inversely proportional to the distance between them. An ion that has a low charge density will form weaker ionic bonds than an ion that has a high charge density. A compound with weak ionic bonds will therefore be more soluble in water than one with strong ionic bonds. Ions with low charge densities will form more soluble compounds than those with high charge density. –

–

+

Solubility increases in the series below from NaF to RbCl since Cl is larger than F and Rb is + larger than Na NaF (sodium fluoride) NaCl (sodium chloride) RbCl (rubidium chloride) +

–

2–

Polyatomic ions such as NH4 (ammonium ion), NO3 (nitrate ion) or SO4 (sulfate ion) are larger and therefore have lower charge densities than monatomic ions with the same charge. In each of the following pairs, the first compound has weaker ionic bonds and is more soluble than the second compound. NH4Cl (ammonium chloride) and LiCl (lithium chloride) LiNO3 (lithium nitrate) and LiF (lithium fluoride) Al2(SO4)3 (aluminum sulfate) and Al2S3 (aluminum sulfide). Using just the charge density, both RbCl (rubidium chloride) and AgCl (silver chloride) are expected to be very soluble in water. Rubidium and silver are in the same period on the periodic table, have the same charge, are approximately the same size and, therefore, have nearly the same charge density. Experiment shows that AgCl is much less soluble than RbCl, however. The lower solubility of AgCl compared to RbCl is explained by partial covalent nature of the bonding in AgCl. This partial covalent bonding is stronger than the ionic bonding in RbCl. Thus, when determining solubility, the degree of covalency must be considered. The degree to which a bond is ionic or covalent depends on the difference in electronegativities between the bonded atoms: the greater the difference in electronegativity, the more ionic the bond. Electronegativities increase upward in a group and to the right across a period. The farther apart two elements are on the periodic table, the more ionic the bond between them will be. In general, when a compound is more ionic, it is more likely to be soluble in water. Compounds with predominantly covalent bonds are usually insoluble in water. All salts are at least slightly soluble in water, but in many cases this solubility is extremely small. It is common practice to define an arbitrary division between salts that are very soluble and those that are slightly soluble, and to refer to the former as soluble and the latter as insoluble.

MATERIALS 7 test tubes Test tube rack

0.1 M Cation solutions: magnesium calcium barium zinc copper silver lead

0.1M Anion solutions: chloride bromide sulfide hydroxide carbonate phosphate sulfate

(n.b. cation solutions may be chloride or nitrate salts, anion solutions should be sodium salts.)

PROCEDURE CAUTION: Salts containing silver, lead(II) and barium ions are potentially toxic. Avoid skin contact. Construct a data table similar to the one shown on the final page of this experiment before you start any experimental work. 1.

In each of seven numbered test tubes place three drops of a different anion solution in the order shown in the data table. Be sure to replace the droppers into the correct bottles.

2.

Add three drops of Mg

3.

Observe each tube to see if a precipitate forms. If there is no precipitate, write “soluble” in the appropriate square in the data table. If there is a precipitate, carefully describe its color and appearance as light or heavy, and write its formula in the appropriate square.

4.

Rinse the tubes with de-ionized water and drain them thoroughly; they do not need to be dry.

5.

Repeat steps 1-4 for each remaining cation in the data table. Dispose of Ba , Pb , and Ag precipitates and solutions in the labeled waste bottles in the hood.

2+

solution to each tube and mix well.

2+

2+

+

Name: _______________________________

Date: ___________

PRE-LAB QUESTIONS 1. When you combine a cation and an anion, how can you tell if the resulting compound is soluble or insoluble?

2. What makes some ionic compounds more water-soluble than others? Give three reasons.

3. Put these compounds in order of their expected aqueous solubility, starting with the most soluble: CuCl2

NaCl

AgCl

AlCl3

MgCl2

4. Choose the least soluble salt: KNO3

KBr

K2SO4

KF

Name: _____________________________________

Date: ___________

AQUEOUS SOLUBILITY RESULTS & POST-LAB QUESTIONS 1. Write the net ionic equation for any mixture that formed a precipitate (anything insoluble).

2. Based on your observations, write a set of solubility rules. Are there any rules that do not agree with the rules given in your book?

3. From the rules in your book and the results of this experiment, predict whether these salts will be soluble or insoluble. Write the formula for each salt. You have not tested these compounds, so you must use information about similar compounds and by analogy predict their solubility. a. sodium sulfide

f.

tin(II) chloride

b. potassium arsenate

g.

sodium sulfate

c. ammonium oxalate

h.

cobalt(II) acetate

d. nickel(II) chloride

i.

strontium bromide

e. chromium(III) hydroxide

j.

aluminum phosphate

AQUEOUS SOLUBILITY DATA TABLE anions cations Mg

2+

Ca

2+

Ba

2+

2+

Zn

2+

Cu

Ag

Pb

+

2+

Cl

–

Br

–

2–

S

–

OH

2–

CO3

3–

PO4

2–

SO4