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)

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:

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

Strong and Weak Acids and Bases•

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!

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.

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]

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.