7. Electrolytes and Solubility Rules for Salts
7.1 Electrolytes.
Electrolytes are chemical substances that ionize at least to some extent in solutions, i.e. they release ions into solutions. For our purposes, the most important solutions are those of involving water as the solvent. Many ionic solids dissolve in water with complete or very substantial dissociation into ions, e.g. NaCl, MgCl2, CuSO4 etc., and can form quite concentrated solutions. These are called strong electrolytes. Another example is HCl. It is a covalently bonded gas at room temperature, but in water it dissolves to high concentrations and completely dissociates into H+ and Cl-. Sulfuric acid when pure is present mainly as H2SO4 molecules, although there is a modest degree of self-ionization. It is also a strong electrolyte, dissociating mainly into H+ and HSO4-. The HSO4- ion dissociates to a lesser extent.
Some ionic solids dissolve in water to only a slight extent, but largely or fully dissociate in solution. An example is PbSO4. This is soluble in water to only about
8 x 10-4 M and forms mainly Pb+2 and SO42- ions in solution, rather than dissolving as PbSO4. However, such weakly soluble salts, while comprised of cations and anions, are classified as weak electrolytes since their solutions are poorly conducting. Hence to be a strong electrolyte, a compound (ionic or covalent) must form concentrated solutions and dissociate into ions to a high degree.
Many covalently bonded neutral molecules dissolve in water and ionize (form ions) to only a limited extent. Examples are NH3, HCN, HgCl2 etc. Even though they might be highly soluble, they ionize to only limited degrees, e.g.
NH3 + H2O = NH4+ + OH- K = 1.8 ´ 10-5 at 25°C (136)
(The small value of the equilibrium constant indicates that the forward reaction proceeds to only a small extent.) Such compounds are called weak electrolytes. Mercury(II) chloride is an interesting example. It is quite soluble in water and exists in solution largely as HgCl2 molecules. Water itself dissociates very slightly according to:
2H2O = H3O+ + OH- Kw = 1.00 ´ 10-14 at 25°C (137)
Other covalently bonded molecules exhibit virtually no dissociation into ions in aqueous solution. These are termed non-electrolytes. Examples include O2, H2, N2, sugar, alcohols, and many others.
Finally, some highly soluble strong electrolytes exhibit ion pairing at high concentration, e.g.
Cu+2aq + SO42-aq = CuSO4 aq (138)
There is still a substantial fraction of the copper sulfate present as the ions, but a weak complex (which is what a soluble ion pair is) between the ions also begins to form. This lowers the concentration of the ions in solution relative to the concentration of the salt and can affect chemical reaction rates and electrical conductivity. Conductivity is important in electrolysis processes for metal recovery (electrowinning). Electric current in solution is carried by the migration of ions. Solutions used for electrowinning always involve strong electrolytes.
7.2 Solubility Rules for Salts.
Below is a list of the generalized solubility rules for salts in aqueous solution. These enable qualitative prediction of whether a salt may be expected to be soluble in water or insoluble. Insoluble is a relative term. It means only very slightly soluble. Very few compounds have genuinely zero solubility in a solvent. The rules are generalizations that do have exceptions. They need to be applied in the order given. A preceding rule takes precedence. For example, whereas most sulfides are insoluble in water, Na2S is very water soluble. The preceding rule that Na+ salts are soluble overrides the later rule. Most silver salts are insoluble, but AgNO3 is soluble. The preceding rule that nitrates are generally soluble takes precedence.
1. Most Na+, K+ and NH4+ salts are soluble.
2. Most NO3-, CH3CO2-, and ClO4- salts are soluble. (KClO4 is only moderately soluble.)
3. Most Ag+, Pb+2, Pb(IV) and Hg2+2 salts are insoluble. (The nitrates are soluble; there is no Pb(IV) nitrate.)
4. Most Cl-, Br- and I- salts are soluble. (Exceptions are AgCl, PbCl2 and Hg2Cl2, which are insoluble.)
5. Most S2-, CO32-, O2- and OH- salts are insoluble (also most SO32- and CrO42- salts). An important exception is NaOH which is readily soluble.
Note: Oxide ions, O2-, does not survive in water! Upon dissolution (whether soluble or insoluble) all metal oxides react with water to form OH-, e.g.
CaO s + H2O = Ca+2aq + 2OH-aq (139)
Na2O s + H2O l → 2Na+aq + 2OH-aq (140)
6. Most sulfates are soluble with the notable exceptions CaSO4, BaSO4 and PbSO4. Calcium sulfate generally precipitates from aqueous solution as CaSO4·2H2O (gypsum).
7.3 The Effect of Temperature on Solubility.
Most salts and neutral compounds become more soluble in water as temperature is raised. Most gases become less soluble with increasing temperature. Most gases are weakly soluble in water to begin with. Increasing temperature increases the kinetic energy of the gas solute molecules, increasing their tendency to be in the gas phase.
However, there are plenty of exceptions to these general rules. Dissolution is a simple chemical reaction; a change of state, e.g.
NaCl s = Na+aq + Cl-aq (141)
A simple rationale for temperature effects has to do with the enthalpy of the reaction. If the enthalpy change for the reaction is positive (H), then heat is absorbed by the solution upon dissolution (if it is to remain at the same temperature). Thinking of heat as a type of reactant then, adding more of it (increasing the temperature) will drive the reaction to the right; solubility will increase. But if heat is released upon dissolution (H < 0, e.g. as when NaOH dissolves), then adding more of it will lower the solubility. Solubility then would decrease with increasing temperature. The same applies to gases and molecular solutes.
8. Representations of Organic Compound Structures and Formulas
For organic chemical formulas it is customary to list the carbon, then hydrogen content, then other atoms in alphabetical order. Sometimes formulas are written so that they also give some sense of the structure, e.g. C12H25NaO4S indicates only atomic composition, whereas C12H25OSO3Na indicates both composition and something about the structure; that there are four oxygens attached to the sulfur. (This is an organic derivative of sulfate.)
A shorthand for organic chemical structures is common. This is outlined below.
Group 7, Grouped objectCH3-CH-CH3 may be represented as:
CH3
The single C-C bonds are reduced to straight lines. Hydrogen atoms are usually omitted. Since the valence of carbon is almost always 4, the hydrogen atoms are understood to be present to satisfy this requirement. Benzene has the structure:
This may be represented by:
The arrangement of carbon atoms is represented by a hexagon. The alternating single and double bonds is represented by a circle. Hydrogens are omitted. This simplifies depiction of structures. Hence,
is the representation for styrene, C6H5-CH=CH2. Discrete double bonds are represented by double lines. The complexing agent EDTA, mentioned previously, can be represented as,
An important class of organic compounds in copper solution purification is the hydroxyoximes. A representative compound in this class is 5-alkyl-2-hydroxyacetophenone oxime,
The alkyl group is indicated as R. It may be nonyl (C9H19-) or dodecyl (C12H25-). This and related compounds form stable complexes quite selectively with Cu+2, which enables it to be extracted from aqueous solution and thus purified of impurities.
