# 6: Acid-Base Chemistry

Learning Objectives

• Understand the Brønsted and Lewis definitions of acids and bases.
• Identify conjugate acids and bases, and rules for strong & weak acids/bases, in both Brønsted and Lewis acid-base systems.
• Use Pauling’s rules to predict the pKas of oxoacids.
• Understand the periodic trends of acidic, basic, and amphoteric compounds
• Predict, describe, and rationalize acid/base chemistry in non-aqueous systems, including acidic and basic solvents, aprotic solvents, and molten salts.
• Apply the principles of acid-base chemistry to the design of molecules and Lewis acids with target functions.
• Understand the connection between acid-base chemistry and the stabilization of oxidation states.
• Predict favorable and stable compounds using hard-soft acid-base (HSAB) theory.
• Understand the applications of the ECW model.

Acid-base reactions form the basis of the most common kinds of equilibrium problems which you will encounter in almost any application of chemistry. There are three major classifications of acids and bases: (1) The Arrhenius definition states that an acid produces H+ in solution and a base produces OH- and the (2) Brønsted-Lowry and (3) Lewis definitions of acids and bases. Of particular importance in inorganic chemistry is the "hard and soft (Lewis) acids and bases" (HSAB) theory that is widely used for explaining stability of compounds, reaction mechanisms, and reaction pathways.

• 6.1: Prelude to Acid-Base Chemistry
Acids and bases are important for a number reasons in inorganic chemistry.
• 6.2: Brønsted and Lewis Acids and Bases
There are three major classifications of substances known as acids or bases. The Arrhenius definition states that an acid produces H+ in solution and a base produces OH-. This theory was developed by Svante Arrhenius in 1883. Later, two more sophisticated and general theories were proposed. These are the Brønsted-Lowry and the Lewis definitions of acids and bases.
• 6.3: Molecular Structure and Acid-Base Behavior
Inductive effects and charge delocalization significantly influence the acidity or basicity of a compound. The acid–base strength of a molecule depends strongly on its structure. The weaker the A–H or B–H+ bond, the more likely it is to dissociate to form an $$H^+$$ ion. In addition, any factor that stabilizes the lone pair on the conjugate base favors the dissociation of $$H^+$$, making the conjugate acid a stronger acid.
• 6.4: Oxides
Oxides are chemical compounds with one or more oxygen atoms combined with another element.
• 6.5: Lewis Acids and Bases
The Brønsted-Lowry proton donor-acceptor concept has been one of the most successful theories of Chemistry. But as with any such theory, it is fair to ask if this is not just a special case of a more general theory that could encompass an even broader range of chemical science. In 1916, G.N. Lewis of the University of California proposed that the electron pair is the dominant actor in acid-base chemistry.
• 6.6: Hard and Soft Acids and Bases
Lewis acids and bases can be classified by designating them as hard or soft. "Hard" acids and bases have a high charge (positive for acids, negative for bases) to ionic radius ratio along with higher oxidation states. Hard acids are not very polarizable and have high charge densities. In contact, "Soft" acids or bases have a low charge to radius ratio, with low oxidation states. They are normally larger ions that are polarizable.
• 6.7: References