1. Arrhenius The resistance of an electrolyte is increased when the dilution is doubled. In very dilute solutions the conductivity is nearly proportional to the concentration. The conductivity of a solution is equal to the sum of conductivities of the salt and the solvent. If these laws are not observed, it must be due to a chemical reaction between the substances including the solvent. The electrical resistance rises with increasing viscosity, complexity of the ion, and the molecular mass of the solvent. (incorrect)

2. Bronsted-Lowry acids and bases. A Bronsted-Lowry (BL) acid is defined as any substance that can donate a hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can accept a hydrogen ion (proton). Thus, according to the BL definition, acids and bases must come in what is called conjugate pairs. For example, consider acetic acid dissolved in water:

3. Lewis acids and bases Lewis extended the theory of acids to cover both non-aquoeus systems and systems that do not involve proton transfers. He defin ed a Lewis acid from the point of view of the electrons rather than from the point of view of hydrogen ions (protons) An electron pair donor becomes a Lewis base and an electron pair receiver is a Lewis acid. To see how this affects Arrhenius acid - base behavious consider the reaction between a hydrogen ion and a hydroxide ion H+ + OH- H2O In this reaction the H+ ion is accepting a lone pair donor electrons from the hydroxide (OH-) ion. According to Lewis' definition the H+ is and acid (as we already know). The hydroxide ion is donating a lone pair of electrons and is defined as a Lewis base Summary Lone pair acceptor - Lewis acid Lone pair donor - Lewis base

4. Strong and Weak Acids and Bases

5. Strong Acids Strong acids are 100% ionized in aqueous solution to form the hydronium ion, H3O+ (also written as H+(aq)) and an anion. For example, HCl in water ionizes completely: HCl + H2O ® H3O+(aq) + Cl–(aq) [goes to completion] (or, equivalently, HCl + water ® H+(aq) + Cl– (aq) [goes to completion]) There are very few strong acids, but they are extremely important in chemistry since they are excellent sources of H+(aq), a highly reactive ion!

6. Weak Acids Most acids are weak. Weak acids are typically less than 5% ionized in water; thus the predominant species is the un-ionized form. Since relatively small amounts of H+(aq) are formed, weak acids are not very reactive. Typical weak acid ionizations in water are HC2H3O2 + H2O H3O+(aq) + C2H3O2–(aq) (or, equivalently, HC2H3O2 + water H+(aq) + C2H3O2–(aq)) SO2(g) + H2O H2SO3 H+(aq) + HSO3–(aq) (or, equivalently, SO2(g) + 2 H2O H3O+(aq) + HSO3–(aq)) In each case above, reaction proceeds only to a very limited extent; typically over 95% of the weak acid remains un-ionized! Since the predominant form is un-ionized, chemists do not split up weak acids into ions when writing an ionic equation.

7. Strong Bases Strong bases are 100% ionized in aqueous solution to form the hydroxide ion, OH–, and a cation. There are very few strong bases, but they are extremely important in chemistry since they are excellent sources of OH–(aq), a highly reactive ion! Typical ionization reactions are NaOH(s) + water ® Na+(aq) + OH–(aq) [goes to completion] Na2O(s) + H2O ® 2 Na+(aq) + 2 OH–(aq) [goes to completion]

8. Weak Bases The vast majority of bases are weak. Much like weak acids, weak bases are typically less than 5% ionized. Since their water solutions contain low concentrations of OH–(aq), they are not very reactive. Examples of weak base ionization reactions include NH3 + H2O NH4+(aq) + OH–(aq) CH3NH2 + H2O CH3NH3+(aq) + OH–(aq) Cu(OH)2(s) + water Cu2+(aq) + 2 OH–(aq) CuO(s) + H2O Cu2+(aq) + 2 OH–(aq) In each case above, reaction proceeds only to a very limited extent; typically over 95% of the weak base remains un-ionized! Weak bases are therefore not split up into ions when writing ionic equations.