Strong electrolytes are completely dissociated into ions in water. The acid or base molecule does not exist in aqueous solution, only ions. Weak electrolytes are incompletely dissociated. Here are definitions and examples of strong and weak acids and strong and weak bases. Strong acids completely dissociate in water, forming H+ and an anion. There are six strong acids. The others are considered to be weak acids. You should commit the strong acids to memory:
If the acid is 100 percent dissociated in solutions of 1.0 M or less, it is called strong. Sulfuric acid is considered strong only in its first dissociation step; 100 percent dissociation isn't true as solutions become more concentrated. H2SO4 → H+ + HSO4- A weak acid only partially dissociates in water to give H+ and the anion. Examples of weak acids include hydrofluoric acid, HF, and acetic acid, CH3COOH. Weak acids include:
Strong bases dissociate 100 percent into the cation and OH- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases.
* These bases completely dissociate in solutions of 0.01 M or less. The other bases make solutions of 1.0 M and are 100 percent dissociated at that concentration. There are other strong bases than those listed, but they are not often encountered. Examples of weak bases include ammonia, NH3, and diethylamine, (CH3CH2)2NH. Like weak acids, weak bases do not completely dissociate in aqueous solution.
Except for their names and formulas, so far we have treated all acids as equals, especially in a chemical reaction. However, acids can be very different in a very important way. Consider HCl(aq). When HCl is dissolved in H2O, it completely dissociates into H+(aq) and Cl−(aq) ions; all the HCl molecules become ions: HCl → H+(aq) + Cl−(aq) (100%) Any acid that dissociates 100% into ions is called a strong acid. If it does not dissociate 100%, it is a weak acid. HC2H3O2 is an example of a weak acid: HC2H3O2→ H+(aq) + C2H3O2−(aq) (~5%) Because this reaction does not go 100% to completion, it is more appropriate to write it as an equilibrium: HC2H3O2 ⇄ H+(aq) + C2H3O2−(aq) As it turns out, there are very few strong acids, which are given in Table 12.1 “Strong Acids and Bases”. If an acid is not listed here, it is a weak acid. It may be 1% ionized or 99% ionized, but it is still classified as a weak acid. The issue is similar with bases: a strong base is a base that is 100% ionized in solution. If it is less than 100% ionized in solution, it is a weak base. There are very few strong bases (see Table 12.1); any base not listed is a weak base. All strong bases are OH– compounds. So a base based on some other mechanism, such as NH3 (which does not contain OH− ions as part of its formula), will be a weak base.
Identify each acid or base as strong or weak. Solution
Test Yourself Answers
Write the balanced chemical equation for the dissociation of Ca(OH)2 and indicate whether it proceeds 100% to products or not. Solution Ca(OH)2 → Ca2+(aq) + 2OH−(aq) Because Ca(OH)2 is listed in Table 12.1, this reaction proceeds 100% to products. Test Yourself Answer HN3 → H+(aq) + N3−(aq) It does not proceed 100% to products because hydrazoic acid is not a strong acid. Certain salts will also affect the acidity or basicity of aqueous solutions because some of the ions will undergo hydrolysis, just like NH3 does to make a basic solution. The general rule is that salts with ions that are part of strong acids or bases will not hydrolyze, while salts with ions that are part of weak acids or bases will hydrolyze. Consider NaCl. When it dissolves in an aqueous solution, it separates into Na+ ions and Cl− ions: NaCl → Na+(aq) + Cl−(aq) Will the Na+(aq) ion hydrolyze? If it does, it will interact with the OH− ion to make NaOH: Na+(aq) + H2O → NaOH + H+(aq) However, NaOH is a strong base, which means that it is 100% ionized in solution: NaOH → Na+(aq) + OH−(aq) The free OH−(aq) ion reacts with the H+(aq) ion to remake a water molecule: H+(aq) + OH−(aq) → H2O The net result? There is no change, so there is no effect on the acidity or basicity of the solution from the Na+(aq) ion. What about the Cl− ion? Will it hydrolyze? If it does, it will take an H+ ion from a water molecule: Cl−(aq) + H2O → HCl + OH− However, HCl is a strong acid, which means that it is 100% ionized in solution: HCl → H+(aq) + Cl−(aq) The free H+(aq) ion reacts with the OH−(aq) ion to remake a water molecule: H+(aq) + OH−(aq) → H2O The net result? There is no change, so there is no effect on the acidity or basicity of the solution from the Cl−(aq) ion. Because neither ion in NaCl affects the acidity or basicity of the solution, NaCl is an example of a neutral salt. Things change, however, when we consider a salt like NaC2H3O2. We already know that the Na+ ion won’t affect the acidity of the solution. What about the acetate ion? If it hydrolyzes, it will take an H+ from a water molecule: C2H3O2−(aq) + H2O → HC2H3O2 + OH−(aq) Does this happen? Yes, it does. Why? Because HC2H3O2 is a weak acid. Any chance a weak acid has to form, it will (the same with a weak base). As some C2H3O2− ions hydrolyze with H2O to make the molecular weak acid, OH− ions are produced. OH− ions make solutions basic. Thus NaC2H3O2 solutions are slightly basic, so such a salt is called a basic salt. There are also salts whose aqueous solutions are slightly acidic. NH4Cl is an example. When NH4Cl is dissolved in H2O, it separates into NH4+ ions and Cl− ions. We have already seen that the Cl− ion does not hydrolyze. However, the NH4+ ion will: NH4+(aq) + H2O → NH3(aq) + H3O+(aq) Recall from the section “Arrhenius Acids and Bases” that H3O+ ion is the hydronium ion, the more chemically proper way to represent the H+ ion. This is the classic acid species in solution, so a solution of NH4+(aq) ions is slightly acidic. NH4Cl is an example of an acid salt. The molecule NH3 is a weak base, and it will form when it can, just like a weak acid will form when it can. So there are two general rules:
Identify each salt as acidic, basic, or neutral. Solution
Test Yourself Answers Some salts are composed of ions that come from both weak acids and weak bases. The overall effect on an aqueous solution depends on which ion exerts more influence on the overall acidity. We will not consider such salts here.
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