Acids & Bases

Brønsted-Lowry Acid: A chemical species that acts as a proton donor, releasing a hydrogen ion ($H^+$) during a reaction. Since a hydrogen atom consists of one proton and one electron, the loss of that electron leaves only the nucleus, which is why $H^+$ is synonymous with a 'proton'.

Brønsted-Lowry Base: A chemical species that acts as a proton acceptor by using a lone pair of electrons to form a dative covalent bond with an incoming $H^+$ ion. This definition expands the scope of bases beyond just hydroxide-containing compounds to include species like ammonia ($NH_3$).

Amphoteric Substances: Molecules or ions that can behave as either an acid or a base depending on the reaction environment. Water ($H_2O$) is the most common example, as it can donate a proton to become $OH^-$ or accept one to become $H_3O^+$.

Strong Acids and Bases: These species undergo complete dissociation in aqueous solution. For a strong acid $HA$, the reaction $HA \rightarrow H^+ + A^-$ goes to completion, meaning the concentration of $H^+$ ions is essentially equal to the initial concentration of the acid.

Weak Acids and Bases: These species only partially dissociate, establishing a dynamic equilibrium between the molecular form and the ionized forms. The position of equilibrium lies heavily to the left, resulting in a much lower concentration of $H^+$ or $OH^-$ ions compared to a strong acid of the same molarity.

Degree of Dissociation: This is a measure of the fraction of molecules that have ionized. It is influenced by the chemical nature of the substance and is independent of the total concentration of the solution.

The pH Scale: A logarithmic scale used to specify the acidity or basicity of an aqueous solution. It is defined by the negative base-10 logarithm of the hydrogen ion concentration: $pH = -\log_{10}[H^+]$. Because it is logarithmic, a change of one pH unit represents a tenfold change in $H^+$ concentration.

Ionic Product of Water ($K_w$): Water undergoes self-ionization to a very small extent: $H_2O \rightleftharpoons H^+ + OH^-$. At $298 K$, the equilibrium constant $K_w = [H^+][OH^-] = 1.0 \times 10^{-14} mol^2 dm^{-6}$.

Neutrality Condition: A solution is neutral when $[H^+] = [OH^-]$. At $298 K$, this occurs at $pH = 7$. If $pH < 7$, the solution is acidic ($[H^+] > [OH^-]$); if $pH > 7$, it is alkaline ($[OH^-] > [H^+]$).

Identify Conjugate Pairs: Always look for species that differ by exactly one $H^+$ ion. The species with the extra $H^+$ is the acid, and the one without it is the base. In an exam, you may be asked to label these in a given equation.

Diprotic Acids: Be careful with acids like $H_2SO_4$. The first dissociation is usually strong, but the second ($HSO_4^- \rightleftharpoons H^+ + SO_4^{2-}$) is weak. If an exam asks for the pH of a strong diprotic acid, check if you should assume both protons dissociate fully or just the first.

Significant Figures: When converting $[H^+]$ to $pH$, the number of decimal places in the $pH$ value should match the number of significant figures in the concentration value. For example, $[H^+] = 1.0 \times 10^{-3}$ (2 sig figs) results in $pH = 3.00$ (2 decimal places).

The 'Strong = Concentrated' Myth: Students often assume that a strong acid is always more dangerous or 'stronger' in effect than a weak one. However, a highly concentrated weak acid (like glacial ethanoic acid) can be more corrosive than a very dilute strong acid (like $0.0001 M$ $HCl$).

pH Limits: While the standard scale is $0-14$, $pH$ can actually be negative for very concentrated strong acids or greater than $14$ for very concentrated strong bases. Do not assume a calculation is wrong just because it falls slightly outside the $0-14$ range.

Temperature Dependence: $K_w$ changes with temperature because dissociation is an endothermic process. As temperature increases, $K_w$ increases, which means the $pH$ of pure water decreases even though it remains neutral ($[H^+]$ still equals $[OH^-]$).