Classification and Models of Acids and Bases

The Lewis Model: Electron Pair Transfer

The Brønsted-Lowry model is limited because it requires the presence of a proton. To broaden the definition, chemists use the Lewis model, which describes acid-base behavior in terms of electron pair transfer.

• Lewis Base: A substance that can donate a pair of electrons. Every Brønsted-Lowry base is also a Lewis base.
• Lewis Acid: A substance that can accept a pair of electrons.

Lewis acids do not necessarily need a proton; they simply need a vacant orbital to accommodate an incoming electron pair. This category includes cations (like $Mg^{2+}$), species with incomplete octets (like Boron), and certain compounds with multiple bonds (like $CO_2$).

Acid and Base Strength

The "strength" of an acid or base refers to how completely it dissociates (breaks apart) in an aqueous solution.

Strong Acids and Bases

Strong acids and bases are considered strong electrolytes. They undergo nearly 100% dissociation in water. When writing their reactions, we use a single arrow ($\rightarrow$) rather than an equilibrium arrow ($\rightleftharpoons$) because the reactants are almost entirely consumed.

• Common Strong Acids: $HCl$, $HBr$, $HI$, $HNO_3$, $H_2SO_4$, $HClO_4$.
• Common Strong Bases: Group 1 hydroxides (like $NaOH$, $KOH$) and Group 2 hydroxides (like $Ca(OH)_2$, $Ba(OH)_2$).

Weak Acids and Bases

Weak acids and bases are weak electrolytes. They do not dissociate completely; instead, they exist in a state of equilibrium where the undissociated molecule predominates. These reactions are represented with equilibrium arrows ($\rightleftharpoons$), and their strength is measured by equilibrium constants, $K_a$ (for acids) and $K_b$ (for bases).

Pro-Tip: In any acid-base pairing, the equilibrium will always favor the side with the weaker acid and the weaker base.

Protonation States and Biological Significance

The protonation state of a molecule refers to whether it has gained or lost a proton, which directly determines its electrical charge.

For example, ammonia ($NH_3$) can exist as the neutral molecule $NH_3$ or the positively charged ammonium ion $NH_4^+$. The specific ratio of these two forms in a solution is determined by the pH.

This is critically important in medicine. Highly charged molecules pass through cell membranes slowly, while neutral molecules pass more quickly. For instance, aspirin (acetylsalicylic acid) is mostly protonated (neutral) in the low pH of the stomach, allowing it to be absorbed more readily into the bloodstream.

Acid-Base Properties of Salts

When an acid and a base react, they form a salt and water. However, not all salts are neutral. The pH of a salt solution depends on whether the constituent ions undergo hydrolysis—the reaction of an ion with water to form $H\_3O^+$ or $OH^-$.

1. Neutral Salts: Formed from a strong acid and a strong base (e.g., $NaCl$).
2. Basic Salts: Formed from a strong base and a weak acid. The anion from the weak acid reacts with water to produce $OH^-$ (e.g., $NaF$).
3. Acidic Salts: Formed from a strong acid and a weak base. The cation from the weak base reacts with water to produce $H\_3O^+$ (e.g., $NH_4Cl$).
4. Highly Acidic Salts: Certain small, highly charged metal cations (like $Al^{3+}$ or $Fe^{3+}$) can act as acids by reacting with water.