6.1 Brønsted-Lowry and Lewis Acid-Base Theories

Brønsted-Lowry Acids and Bases

The Brønsted-Lowry and Lewis theories provide two complementary frameworks for understanding acid-base behavior. Brønsted-Lowry theory centers on proton transfer, while Lewis theory broadens the picture to electron pair donation and acceptance. Together, they explain reactivity patterns across aqueous solutions, coordination chemistry, and organometallic systems.

Defining Acids and Bases in the Brønsted-Lowry Theory

A Brønsted-Lowry acid is a proton ($H^+$) donor, and a Brønsted-Lowry base is a proton acceptor. Every Brønsted-Lowry acid-base reaction involves a proton being transferred from one species to another.

When an acid donates its proton, it becomes a conjugate base. When a base accepts a proton, it becomes a conjugate acid. These paired species are called conjugate acid-base pairs, and they always differ by exactly one proton.

• $HCl$ donates $H^+$ → its conjugate base is $Cl^-$
• $NH_3$ accepts $H^+$ → its conjugate acid is $NH_4^+$

Amphoteric (or amphiprotic) substances can act as either an acid or a base depending on what they're reacting with. Water is the classic example: it donates a proton to strong bases and accepts a proton from acids. The bicarbonate ion ($HCO_3^-$) is another common one.

Examples and Applications of Brønsted-Lowry Theory

$HCl$ in water (acid behavior):

$HCl + H_2O \rightarrow H_3O^+ + Cl^-$

Here $HCl$ donates a proton to $H_2O$. Two conjugate pairs exist in this reaction: $HCl / Cl^-$ and $H_3O^+ / H_2O$. Notice that water acts as the base here.

$NH_3$ in water (base behavior):

$NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$

Now water acts as the acid, donating a proton to ammonia. The equilibrium arrow ($\rightleftharpoons$) signals this is a weak base reaction; it doesn't go to completion. The conjugate pairs are $NH_4^+ / NH_3$ and $H_2O / OH^-$.

Amino acids demonstrate amphoteric behavior in biological systems. The carboxyl group ($-COOH$) can donate a proton, while the amino group ($-NH_2$) can accept one. This is why amino acids exist as zwitterions at physiological pH.

Lewis Acids and Bases

Fundamentals of Lewis Acid-Base Theory

Lewis theory is broader than Brønsted-Lowry. A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor. The key product of a Lewis acid-base reaction is a new coordinate covalent bond (also called a dative bond), where both electrons in the bond come from the Lewis base.

Every Brønsted-Lowry acid-base reaction is also a Lewis acid-base reaction (the proton is the electron pair acceptor), but Lewis theory captures many reactions that Brønsted-Lowry cannot, particularly in coordination chemistry and reactions involving species with no transferable proton.

The HSAB (Hard-Soft Acid-Base) principle further categorizes Lewis acids and bases:

• Hard acids and bases are small, highly charged, and weakly polarizable (e.g., $Fe^{3+}$, $F^-$, $OH^-$)
• Soft acids and bases are large, low in charge, and highly polarizable (e.g., $Pt^{2+}$, $I^-$, $CO$)

The guiding rule: hard prefers hard, soft prefers soft. This principle predicts which complexes and mineral pairings are most stable.

Examples and Applications of Lewis Acid-Base Theory

$BF_3$ with $NH_3$ (adduct formation):

$BF_3 + NH_3 \rightarrow BF_3 \cdot NH_3$

Boron in $BF_3$ has an empty p-orbital, making it electron-deficient and a strong Lewis acid. The lone pair on nitrogen in $NH_3$ donates into that orbital, forming a coordinate bond.

Metal aqua complexes:

$Cu^{2+} + 6H_2O \rightarrow [Cu(H_2O)_6]^{2+}$

The $Cu^{2+}$ ion accepts electron pairs from six water molecules, each acting as a Lewis base. This is the foundation of coordination chemistry: metal ions are Lewis acids, and ligands are Lewis bases.

Carbocations ($R_3C^+$) in organic reactions are Lewis acids. Their empty p-orbital accepts electron pairs from nucleophiles (Lewis bases) during substitution and addition reactions.

HSAB in coordination chemistry:

• Hard acid $Fe^{3+}$ binds preferentially to hard base $F^-$ (forming stable $FeF_6^{3-}$)
• Soft acid $Pt^{2+}$ binds preferentially to soft base $I^-$
• This explains why certain metal-ligand combinations are far more stable than others, and it's useful for predicting which complexes will form in solution

Acid-Base Strength and pH

Understanding Acid-Base Strength

Acid-base strength describes how completely a species donates or accepts protons in solution.

Strong acids (like $HCl$, $HNO_3$, $H_2SO_4$) and strong bases (like $NaOH$, $KOH$) dissociate completely in water. There's essentially no undissociated acid or base left in solution.

Weak acids (like $CH_3COOH$) and weak bases (like $NH_3$) only partially dissociate, establishing an equilibrium. Their strength is quantified by $K_a$ (for acids) or $K_b$ (for bases); larger values mean stronger acids or bases.

There's an inverse relationship between an acid and its conjugate base:

• A strong acid has a very weak conjugate base (e.g., $HCl$ is strong; $Cl^-$ is an extremely weak base)
• A weak acid has a relatively stronger conjugate base (e.g., $CH_3COOH$ is weak; $CH_3COO^-$ is a moderately strong base)

This relationship is expressed quantitatively as $K_a \times K_b = K_w = 1.0 \times 10^{-14}$ at 25°C.

pH Scale and Its Applications

The pH scale quantifies how acidic or basic a solution is, ranging from 0 to 14 in aqueous solutions at 25°C.

• pH < 7 → acidic
• pH = 7 → neutral
• pH > 7 → basic

pH is calculated from the hydrogen ion concentration:

$pH = -\log[H^+]$

Similarly, pOH measures hydroxide ion concentration:

$pOH = -\log[OH^-]$

These two quantities are linked at 25°C by:

$pH + pOH = 14$

This relationship comes from the autoionization of water, where $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$.

Buffer solutions resist changes in pH when small amounts of acid or base are added. They consist of a weak acid paired with its conjugate base (or a weak base with its conjugate acid). For example, the carbonic acid/bicarbonate buffer system maintains blood pH near 7.4. Buffers work because the weak acid neutralizes added base, and the conjugate base neutralizes added acid, keeping $[H^+]$ relatively stable.