🧶Inorganic Chemistry I Unit 6 Review

Brønsted-Lowry theory focuses on proton transfer. A Brønsted-Lowry acid donates a proton ( $H^+$ ), while a Brønsted-Lowry base accepts one. A classic example:

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

Here, $HCl$ donates a proton to water, making $HCl$ the acid and $H_2O$ the base.

Lewis theory is broader. A Lewis acid accepts an electron pair, and a Lewis base donates one, forming a coordinate covalent bond. This framework captures reactions that don't involve protons at all:

$BF_3 + NH_3 \rightarrow F_3B\text{:}NH_3$

$BF_3$ is the Lewis acid (electron-pair acceptor) and $NH_3$ is the Lewis base (electron-pair donor). In inorganic chemistry, Lewis theory is especially useful for understanding metal-ligand interactions and reactions of oxides that don't neatly fit the Brønsted-Lowry model.

Conjugate Acid-Base Pairs and Amphoteric Behavior

A conjugate acid-base pair consists of two species that differ by exactly one proton. When an acid donates a proton, what remains is its conjugate base. When a base accepts a proton, the product is its conjugate acid.

For example, in the reaction $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$ , the pairs are $NH_3 / NH_4^+$ and $H_2O / OH^-$ .

A key relationship: the stronger an acid, the weaker its conjugate base, and vice versa. This inverse relationship helps you predict the direction of proton-transfer equilibria.

Amphoteric species can act as either an acid or a base depending on what they react with. Aluminum oxide ( $Al_2O_3$ ) is a textbook example:

With a strong acid: $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$ (behaves as a base)
With a strong base: $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + H_2O$ (behaves as an acid)

Water itself is amphoteric, acting as a base with $HCl$ and as an acid with $NH_3$ .

Oxides and Their Properties

Brønsted-Lowry and Lewis Acid-Base Concepts, Brønsted-Lowry Acids and Bases | Chemistry for Majors

Acidic and Basic Oxides

The acid-base character of an oxide depends largely on whether the element bonded to oxygen is a metal or a nonmetal, and on its oxidation state.

Acidic oxides (also called acid anhydrides) are typically nonmetal oxides. They dissolve in water to produce acidic solutions:

$CO_2$ , $SO_2$ , $SO_3$ , $N_2O_5$ , $P_4O_{10}$

Basic oxides are typically metal oxides, especially those of the alkali and alkaline earth metals. They dissolve in water to produce basic solutions:

$Na_2O$ , $CaO$ , $MgO$ , $BaO$

Two trends to remember:

Oxide acidity increases as the oxidation state of the central atom increases. For example, $CrO_3$ (Cr in +6) is acidic, while $CrO$ (Cr in +2) is basic.
Oxide basicity increases with increasing ionic character of the metal-oxygen bond. This is why Group 1 oxides (highly ionic) are strongly basic, while oxides of metals with high oxidation states and more covalent bonding tend toward acidic behavior.

These trends connect directly to position on the periodic table: moving left and down gives more basic oxides; moving right and up gives more acidic ones.

Hydrolysis and Neutralization Reactions

Hydrolysis is the reaction of an oxide with water to form an acid or a base.

Acidic oxide hydrolysis:

$SO_3 + H_2O \rightarrow H_2SO_4$

Basic oxide hydrolysis:

$CaO + H_2O \rightarrow Ca(OH)_2$

Not all oxides dissolve readily in water. Many transition metal oxides and amphoteric oxides (like $Al_2O_3$ ) are insoluble, so their acid-base character is demonstrated through direct reactions with acids or bases rather than hydrolysis.

Neutralization can also occur directly between an acidic oxide and a basic oxide, producing a salt without water as a solvent:

$CaO + SO_3 \rightarrow CaSO_4$

This is conceptually the same as neutralizing an acid with a base, just skipping the aqueous step.

Aqueous Solutions and pH

Brønsted-Lowry and Lewis Acid-Base Concepts, Proton donors and acceptors

pH Scale and Buffer Solutions

The pH scale quantifies how acidic or basic an aqueous solution is, based on the concentration of hydrogen ions:

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

At 25°C, pH ranges from 0 (strongly acidic) to 14 (strongly basic), with 7 being neutral. Each whole-number change represents a tenfold change in $[H^+]$ , so a solution at pH 3 is ten times more acidic than one at pH 4.

Buffer solutions resist changes in pH when small amounts of acid or base are added. A buffer is made from a weak acid and its conjugate base (or a weak base and its conjugate acid). For example, an acetic acid/acetate buffer: $CH_3COOH / CH_3COO^-$ .

How it works: if you add acid, the conjugate base ( $CH_3COO^-$ ) neutralizes the added $H^+$ . If you add base, the weak acid ( $CH_3COOH$ ) neutralizes the added $OH^-$ . The pH shifts only slightly instead of dramatically.

Buffer capacity depends on two things: the total concentration of the acid-base pair (higher concentration means more buffering power) and how close the ratio of acid to conjugate base is to 1:1. A buffer works best when $pH \approx pK_a$ .

Acid-Base Indicators and Titrations

An acid-base indicator is a weak acid or base that changes color over a specific pH range. You choose an indicator whose color change falls near the expected equivalence point of your titration.

Phenolphthalein: colorless below ~pH 8.2, pink above. Good for strong acid/strong base titrations.
Methyl orange: red below ~pH 3.1, yellow above ~pH 4.4. Useful for titrations with a low-pH equivalence point.

In an acid-base titration, you add a solution of known concentration (the titrant) to a solution of unknown concentration (the analyte) until the reaction is complete.

The equivalence point is where moles of acid exactly equal moles of base.
The endpoint is where the indicator changes color. A well-chosen indicator makes the endpoint and equivalence point nearly coincide.
A titration curve (pH vs. volume of titrant added) shows a characteristic steep rise or drop near the equivalence point. The shape of this curve differs for strong acid/strong base, weak acid/strong base, and other combinations.

Acid-Base Equilibrium Constants

Acid and Base Dissociation Constants

The acid dissociation constant ( $K_a$ ) quantifies how completely an acid ionizes in water:

$HA \rightleftharpoons H^+ + A^-$

$K_a = \frac{[H^+][A^-]}{[HA]}$

A large $K_a$ means extensive ionization (strong acid behavior). A small $K_a$ means the acid mostly stays intact (weak acid).

Similarly, the base dissociation constant ( $K_b$ ) measures base strength:

$B + H_2O \rightleftharpoons BH^+ + OH^-$

$K_b = \frac{[BH^+][OH^-]}{[B]}$

Water does not appear in these expressions because it's the solvent and its concentration is treated as constant.

pKa, pKb, and Relationship to pH

Taking the negative log converts these constants to a more convenient scale:

$pK_a = -\log K_a \qquad pK_b = -\log K_b$

Lower $pK_a$ = stronger acid. Higher $pK_a$ = weaker acid. The same logic applies to $pK_b$ and base strength.

For any conjugate acid-base pair in water at 25°C:

$pK_a + pK_b = 14$

This comes from the relationship $K_a \times K_b = K_w = 1.0 \times 10^{-14}$ at 25°C. If you know $pK_a$ for an acid, you immediately know $pK_b$ for its conjugate base.

The Henderson-Hasselbalch equation connects pH to $pK_a$ in buffer solutions:

$pH = pK_a + \log\frac{[A^-]}{[HA]}$

When $[A^-] = [HA]$ , the log term is zero and $pH = pK_a$ . This is the point of maximum buffer capacity, and it's a useful shortcut for estimating buffer pH.

🧶Inorganic Chemistry I Unit 6 Review