{"id":1181,"date":"2025-11-14T03:28:47","date_gmt":"2025-11-14T08:28:47","guid":{"rendered":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/?post_type=chapter&p=1181"},"modified":"2026-03-20T16:19:19","modified_gmt":"2026-03-20T20:19:19","slug":"7-electrolytes-and-solubility-rules-for-salts","status":"publish","type":"chapter","link":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/chapter\/7-electrolytes-and-solubility-rules-for-salts\/","title":{"raw":"7. Electrolytes and Solubility Rules for Salts","rendered":"7. Electrolytes and Solubility Rules for Salts"},"content":{"raw":"

7.1 Electrolytes.<\/h2>\r\nElectrolytes 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<\/sub>, CuSO4<\/sub> 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 H2<\/sub>SO4<\/sub> molecules, although there is a modest degree of self-ionization. It is also a strong electrolyte, dissociating mainly into H+ and HSO4<\/sub>-<\/sup>. The HSO4<\/sub>-<\/sup> ion dissociates to a lesser extent.\r\n\r\nSome ionic solids dissolve in water to only a slight extent, but largely or fully dissociate in solution. An example is PbSO4<\/sub>. This is soluble in water to only about 8 x 10-4<\/sup> M and forms mainly Pb+2<\/sup> and SO4<\/sub>2-<\/sup> ions in solution, rather than dissolving as PbSO4<\/sub>. 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.\r\n\r\nMany covalently bonded neutral molecules dissolve in water and ionize (form ions) to only a limited extent. Examples are NH3<\/sub>, HCN, HgCl2<\/sub> etc. Even though they might be highly soluble, they ionize to only limited degrees, e.g.\r\n\r\n\\[\\ce{NH3 + H2O <=> NH4+ + OH-},\\quad K_{25^\\circ C} = 1.8 \\times 10^{-5}\\tag{163}\\]\r\n\r\n(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<\/sub> molecules. Water itself dissociates very slightly according to:\r\n\r\n\\[\\ce{2H2O <=> H3O+ + OH-},\\quad K_\\mathrm{w} = 1.00 \\times 10^{-14}\\ \\text{at } 25^\\circ C \\tag{164}\\]\r\n\r\nOther covalently bonded molecules exhibit virtually no dissociation into ions in aqueous solution. These are termed non-electrolytes. Examples include O2<\/sub>, H2<\/sub>, N2<\/sub>, sugar, alcohols, and many others.\r\n\r\nFinally, some highly soluble strong electrolytes exhibit ion pairing at high concentration, e.g.\r\n\r\n\\[\\ce{Cu^{+2}_{(aq)} + SO_4^2-_{(aq)} = CuSO4_{(aq)}} \\tag{165}\\]\r\n\r\nThere 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.\r\n

7.2 Solubility Rules for Salts.<\/h2>\r\nBelow 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, Na2<\/sub>S is very water soluble. The preceding rule that Na+ salts are soluble overrides the later rule. Most silver salts are insoluble, but AgNO3<\/sub> is soluble. The preceding rule that nitrates are generally soluble takes precedence.\r\n
    \r\n \t
  1. Most Na+<\/sup>, K+<\/sup> and NH4<\/sub>+<\/sup> salts are soluble.<\/li>\r\n \t
  2. Most NO3<\/sub>-<\/sup>, CH3<\/sub>CO2<\/sub>-<\/sup>, and ClO4<\/sub>-<\/sup> salts are soluble. (KClO4<\/sub> is only moderately soluble.)<\/li>\r\n \t
  3. Most Ag+<\/sup>, Pb+2<\/sup>, Pb(IV) and Hg2<\/sub>+2<\/sup> salts are insoluble. (The nitrates are soluble; there is no Pb(IV) nitrate.)<\/li>\r\n \t
  4. Most Cl-<\/sup>, Br-<\/sup> and I-<\/sup> salts are soluble. (Exceptions are AgCl, PbCl2<\/sub> and Hg2<\/sub>Cl2<\/sub>, which are insoluble.)<\/li>\r\n \t
  5. Most S2<\/sub>-<\/sup>, CO3<\/sub>2-<\/sup>, O2<\/sub>-<\/sup> and OH-<\/sup> salts are insoluble (also most SO3<\/sub>2-<\/sup> and CrO4<\/sub>2-<\/sup> salts). An important exception is NaOH which is readily soluble.<\/li>\r\n<\/ol>\r\n
    Note:<\/strong> Oxide ions, O2-<\/sup>, does not survive in water! Upon dissolution (whether soluble or insoluble) all metal oxides react with water to form OH-<\/sup>, e.g.<\/blockquote>\r\n\\[\\ce{CaO_{(s)} + H2O_{(l)} = Ca+^2_{(aq)} + 2OH-_{(aq)}} \\tag{166}\\]\r\n\r\n\\[\\ce{Na2O_{(s)} + H2O_{(l)} \u2192 2Na+_{(aq)} + 2OH-_{(aq)}} \\tag{167}\\]\r\n\r\n6. Most sulfates are soluble with the notable exceptions CaSO4<\/sub>, BaSO4<\/sub> and PbSO4<\/sub>. Calcium sulfate generally precipitates from aqueous solution as CaSO4<\/sub>\u00b72H2<\/sub>O (gypsum).\r\n

    7.3 The Effect of Temperature on Solubility.<\/h2>\r\nMost 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.\r\n\r\nHowever, there are plenty of exceptions to these general rules. Dissolution is a simple chemical reaction; a change of state, e.g.\r\n\r\n\\[\\ce{NaCl_{(s)} = Na+_{(aq)} + Cl-_{(aq)}} \\tag{168}\\]\r\n\r\nA simple rationale for temperature effects has to do with the enthalpy of the reaction. If the enthalpy change for the reaction is positive (\u0394H), 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 (\u0394H < 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.","rendered":"

    7.1 Electrolytes.<\/h2>\n

    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<\/sub>, CuSO4<\/sub> 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 H2<\/sub>SO4<\/sub> molecules, although there is a modest degree of self-ionization. It is also a strong electrolyte, dissociating mainly into H+ and HSO4<\/sub>–<\/sup>. The HSO4<\/sub>–<\/sup> ion dissociates to a lesser extent.<\/p>\n

    Some ionic solids dissolve in water to only a slight extent, but largely or fully dissociate in solution. An example is PbSO4<\/sub>. This is soluble in water to only about 8 x 10-4<\/sup> M and forms mainly Pb+2<\/sup> and SO4<\/sub>2-<\/sup> ions in solution, rather than dissolving as PbSO4<\/sub>. 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.<\/p>\n

    Many covalently bonded neutral molecules dissolve in water and ionize (form ions) to only a limited extent. Examples are NH3<\/sub>, HCN, HgCl2<\/sub> etc. Even though they might be highly soluble, they ionize to only limited degrees, e.g.<\/p>\n

    \\[\\ce{NH3 + H2O <=> NH4+ + OH-},\\quad K_{25^\\circ C} = 1.8 \\times 10^{-5}\\tag{163}\\]<\/p>\n

    (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<\/sub> molecules. Water itself dissociates very slightly according to:<\/p>\n

    \\[\\ce{2H2O <=> H3O+ + OH-},\\quad K_\\mathrm{w} = 1.00 \\times 10^{-14}\\ \\text{at } 25^\\circ C \\tag{164}\\]<\/p>\n

    Other covalently bonded molecules exhibit virtually no dissociation into ions in aqueous solution. These are termed non-electrolytes. Examples include O2<\/sub>, H2<\/sub>, N2<\/sub>, sugar, alcohols, and many others.<\/p>\n

    Finally, some highly soluble strong electrolytes exhibit ion pairing at high concentration, e.g.<\/p>\n

    \\[\\ce{Cu^{+2}_{(aq)} + SO_4^2-_{(aq)} = CuSO4_{(aq)}} \\tag{165}\\]<\/p>\n

    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.<\/p>\n

    7.2 Solubility Rules for Salts.<\/h2>\n

    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, Na2<\/sub>S is very water soluble. The preceding rule that Na+ salts are soluble overrides the later rule. Most silver salts are insoluble, but AgNO3<\/sub> is soluble. The preceding rule that nitrates are generally soluble takes precedence.<\/p>\n

      \n
    1. Most Na+<\/sup>, K+<\/sup> and NH4<\/sub>+<\/sup> salts are soluble.<\/li>\n
    2. Most NO3<\/sub>–<\/sup>, CH3<\/sub>CO2<\/sub>–<\/sup>, and ClO4<\/sub>–<\/sup> salts are soluble. (KClO4<\/sub> is only moderately soluble.)<\/li>\n
    3. Most Ag+<\/sup>, Pb+2<\/sup>, Pb(IV) and Hg2<\/sub>+2<\/sup> salts are insoluble. (The nitrates are soluble; there is no Pb(IV) nitrate.)<\/li>\n
    4. Most Cl–<\/sup>, Br–<\/sup> and I–<\/sup> salts are soluble. (Exceptions are AgCl, PbCl2<\/sub> and Hg2<\/sub>Cl2<\/sub>, which are insoluble.)<\/li>\n
    5. Most S2<\/sub>–<\/sup>, CO3<\/sub>2-<\/sup>, O2<\/sub>–<\/sup> and OH–<\/sup> salts are insoluble (also most SO3<\/sub>2-<\/sup> and CrO4<\/sub>2-<\/sup> salts). An important exception is NaOH which is readily soluble.<\/li>\n<\/ol>\n

      Note:<\/strong> Oxide ions, O2-<\/sup>, does not survive in water! Upon dissolution (whether soluble or insoluble) all metal oxides react with water to form OH–<\/sup>, e.g.<\/p><\/blockquote>\n

      \\[\\ce{CaO_{(s)} + H2O_{(l)} = Ca+^2_{(aq)} + 2OH-_{(aq)}} \\tag{166}\\]<\/p>\n

      \\[\\ce{Na2O_{(s)} + H2O_{(l)} \u2192 2Na+_{(aq)} + 2OH-_{(aq)}} \\tag{167}\\]<\/p>\n

      6. Most sulfates are soluble with the notable exceptions CaSO4<\/sub>, BaSO4<\/sub> and PbSO4<\/sub>. Calcium sulfate generally precipitates from aqueous solution as CaSO4<\/sub>\u00b72H2<\/sub>O (gypsum).<\/p>\n

      7.3 The Effect of Temperature on Solubility.<\/h2>\n

      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.<\/p>\n

      However, there are plenty of exceptions to these general rules. Dissolution is a simple chemical reaction; a change of state, e.g.<\/p>\n

      \\[\\ce{NaCl_{(s)} = Na+_{(aq)} + Cl-_{(aq)}} \\tag{168}\\]<\/p>\n

      A simple rationale for temperature effects has to do with the enthalpy of the reaction. If the enthalpy change for the reaction is positive (\u0394H), 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 (\u0394H < 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.<\/p>\n","protected":false},"author":2529,"menu_order":7,"template":"","meta":{"pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":[],"pb_section_license":""},"chapter-type":[],"contributor":[],"license":[],"class_list":["post-1181","chapter","type-chapter","status-publish","hentry"],"part":1126,"_links":{"self":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1181","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/users\/2529"}],"version-history":[{"count":27,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1181\/revisions"}],"predecessor-version":[{"id":3778,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1181\/revisions\/3778"}],"part":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/parts\/1126"}],"metadata":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1181\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/media?parent=1181"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapter-type?post=1181"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/contributor?post=1181"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/license?post=1181"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}