Chapter 4: Acids and Bases

Practice flashcards covering definitions, chemical equations, and stability factors for Brønsted-Lowry and Lewis acid-base chemistry based on lecture notes.

Brønsted-Lowry acid

A substance that can donate a proton ($H^+$).

Brønsted-Lowry base

A substance that can accept a proton ($H^+$).

Lewis acid

A substance that can accept a pair of electrons.

Lewis base

A substance that can donate a pair of electrons.

$$K_w$$

The equilibrium constant for water, where $K_w = [H^+][OH^-] = 1 \times 10^{-14}$.

Neutral solution conditions

A solution where $[H^+] = [OH^-] = 1 \times 10^{-7}$ and $pH = 7$.

Acidic solution conditions

A solution where $[H^+] > 1 \times 10^{-7}$ and the $pH$ range is $0$ to $7$.

Basic solution conditions

A solution where $[H^+] < 1 \times 10^{-7}$ and the $pH$ range is $7$ to $14$.

Henderson-Hasselbalch equation (for acids)

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

Relationship between $pK_a$ and $pK_b$

The relationship between conjugate acid-base pairs expressed as $pK_a + pK_b = 14$.

Strong Acid (properties)

Characterized by a high $[H^+]$, high $K_a$, and a low $pK_a$.

Strong Base (properties)

Characterized by a high $[OH^-]$, high $K_b$, and a low $pK_b$.

Equilibrium Preference in Acid-Base Reactions

The equilibrium will favor the formation of the weaker acid (the one with the higher $pK_a$ value) or the weaker base pair.

Electronegativity (in Acid-Base context)

The measure of an element’s affinity for an electron or its ability to accept an electron; acidity increases as the electronegativity of the charge-bearing atom increases within the same row.

Atomic Size (in Acid-Base context)

A factor where larger anions disperse negative charge over a larger volume, increasing conjugate base stability; size is more important than electronegativity when comparing atoms in the same column.

Inductive effects

The stabilization of partial or full charges by electron-donating or electron-withdrawing groups transmitted through $\sigma$ bonds.

Nitro Group

A neutral electron-withdrawing group with a positively charged nitrogen and a negative charge that resonates onto two oxygens, significantly increasing acidity (e.g., nitroethanol).

Solvation Effects

The stabilization of a negatively charged conjugate base by a solvent like water, which helps disperse the charge; steric hindrance can decrease acidity by making a compound harder to solvate.

s character (hybridization)

A property where electrons closer to the nucleus are lower in energy and more stable; as $s$ character increases ($sp > sp^2 > sp^3$), the anion becomes more stable and less basic.

Amine Base Strength Factors

Basicity depends on the stability of the amine ($R_3N$) relative to its conjugate acid ($R_3NH^+$) and the availability of the nitrogen lone pair.

Acetonitrile ($pK_b$)

A substance with $sp$ hybridization and a $pK_b$ of $24.1$, making it a very weak base.

Pyridine ($pK_b$)

A substance with $sp^2$ hybridization and a $pK_b$ of $8.75$.

Piperidine ($pK_b$)

A substance with $sp^3$ hybridization and a $pK_b$ of $2.9$, making it more basic than pyridine.

Amides (basicity)

Very weak bases due to the resonance of the lone pair onto the electronegative oxygen atom.