Aqueous Ionic Equilibrium, Acids, and Bases

Comprehensive practice flashcards covering Chapters 17 and 18: acid-base theories, aqueous ionic equilibria, buffers, titrations, solubility products, precipitation, and complex ions.

What two combinations of components can be used to prepare a buffer solution?

A buffer can be formed by combining either significant amounts of a weak acid and its conjugate base, or significant amounts of a weak base and its conjugate acid.

What specific buffer system is responsible for maintaining human blood pH within the range of $7.36$ to $7.42$?

A mixture of carbonic acid ($H_2CO_3$) and bicarbonate ion ($HCO_3^-$).

How does Le Châtelier's principle explain the common ion effect in an acetic acid and sodium acetate buffer?

The added acetate ion ($C_2H_3O_2^-$) acts as a common ion that shifts the ionization equilibrium of acetic acid ($HC_2H_3O_2$) to the left, causing it to ionize even less than in pure water and resulting in a less acidic solution.

What is the Henderson-Hasselbalch equation used to calculate buffer pH?

$$pH = pK_a + \text{log}\frac{[\text{base}]}{[\text{acid}]}$$

Under what condition is the "small x" approximation considered valid when applying the Henderson-Hasselbalch equation?

It is valid if the initial concentrations of the acid and conjugate base are not too dilute (typically $100$ to $1000$ times greater than the equilibrium constant $K_a$) and the equilibrium constant is relatively small.

What criteria define the optimal effectiveness and effective pH range of a buffer?

A buffer is most effective when the concentrations of acid and conjugate base are high and equal; its effective buffering range extends one pH unit on either side of the $pK_a$ (pH = pK_a \buildrel\textstyle.\textstyle.\belowbaselineskip 1).

What is buffer capacity?

The amount of acid or base that can be added to a buffer solution without causing a large change in pH.

What is the difference between the equivalence point and the end point in an acid-base titration?

The equivalence point is the stoichiometric point where the number of moles of base equals the number of moles of acid, whereas the end point is the point where the indicator dye changes color.

What type of titration is shown in this pH curve, and what is significant about the pH at the half-equivalence point?

It represents the titration of a weak base ($0.100\,M\,NH_3$) with a strong acid ($HCl$); at the half-equivalence point, $[NH_3] = [NH_4^+]$ and $pH = pK_a = 9.25$.

What characteristic feature appears on the titration curve when a diprotic acid is titrated with a strong base?

The curve exhibits two separate equivalence points, provided that $K_{a1}$ and $K_{a2}$ differ sufficiently from each other.

What is the solubility product constant ($K_{sp}$)?

The equilibrium constant for the chemical equation representing the dissolution of a sparingly soluble ionic compound in water.

When can $K_{sp}$ values be compared directly to determine the relative solubilities of two compounds?

Only when the compounds dissociate into the same number of ions upon dissolving.

How does the common ion effect influence the solubility of an ionic compound?

The presence of a common ion drives the dissolution equilibrium back to the left (toward solid formation), decreasing the molar solubility compared to its solubility in pure water.

Why does the solubility of an ionic compound containing a basic anion increase in an acidic solution?

Added hydronium ions ($H_3O^+$) react with and neutralize the basic anion (such as $OH^-$ or $F^-$), reducing its concentration and driving the dissolution equilibrium to the right.

How is the reaction quotient ($Q$) used to determine whether precipitation will occur in a solution?

If $Q < K_{sp}$, the solution is unsaturated and no precipitate forms; if $Q = K_{sp}$, the solution is saturated; if $Q > K_{sp}$, the solution is supersaturated and precipitate will form.

According to the qualitative analysis scheme shown, which reagent precipitates Group 1 cations, and which ions are precipitated?

Dilute $6\,M\,HCl$ is added to precipitate Group 1 insoluble chlorides: $Ag^+$, $Pb^{2+}$, and $Hg_2^{2+}$ as $AgCl$, $PbCl_2$, and $Hg_2Cl_2$.

How are Group 5 cations identified in qualitative analysis, and what distinct flame colors are observed for sodium and potassium?

They are identified by flame tests; sodium produces a persistent bright yellow/orange flame, while potassium produces a pale violet/blue flame.

What is a complex ion, and how is its stability represented quantitatively?

A complex ion contains a central metal cation bonded to one or more ligands (Lewis bases); its stability is quantified by the formation constant ($K_f$), the equilibrium constant for its formation from the hydrated metal ion.

Which metal cations form amphoteric hydroxides that dissolve in both acidic and basic solutions?

$Al^{3+}$, $Cr^{3+}$, $Zn^{2+}$, $Pb^{2+}$, and $Sn^{2+}$.

How do Arrhenius and Brønsted-Lowry define acids and bases?

Arrhenius defines acids as producing $H^+$ and bases as producing $OH^-$ in water; Brønsted-Lowry defines acids as proton ($H^+$) donors and bases as proton ($H^+$) acceptors.

What is a conjugate acid-base pair?

Two substances related to each other by the transfer of a single proton ($H^+$).

What are the six common strong acids?

Hydrochloric acid ($HCl$), hydrobromic acid ($HBr$), hydroiodic acid ($HI$), nitric acid ($HNO_3$), perchloric acid ($HClO_4$), and sulfuric acid ($H_2SO_4$).

What are the six common strong bases listed in the lecture?

Lithium hydroxide ($LiOH$), sodium hydroxide ($NaOH$), potassium hydroxide ($KOH$), strontium hydroxide ($Sr(OH)_2$), calcium hydroxide ($Ca(OH)_2$), and barium hydroxide ($Ba(OH)_2$).

What is the autoionization reaction of water, and what is the value of $K_w$ at $25\,^{\circ}\text{C}$?

$H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$ with $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$.

What rule governs the number of significant figures reported for a logarithmic value such as pH?

Only the digits to the right of the decimal point are significant; the number of decimal places in the logarithm must equal the number of significant digits in the original concentration value.

How does increasing the initial concentration of a weak acid affect equilibrium $[H_3O^+]$ versus percent ionization?

Increasing initial acid concentration increases the equilibrium concentration of $[H_3O^+]$, but decreases the percent ionization of the acid.

What mathematical relationship links $K_a$ and $K_b$ for a conjugate acid-base pair in aqueous solution at $25\,^{\circ}\text{C}$?

$K_a \times K_b = K_w = 1.0 \times 10^{-14}$, or $pK_a + pK_b = 14.00$.

Which two categories of cations behave as weak acids in aqueous solutions?

Cations that are the conjugate acids of weak bases (e.g., $NH_4^+$) and small, highly charged metal cations (such as $Al^{3+}$ and $Fe^{3+}$).

How do the periodic trends of electronegativity and bond strength determine the acidity of binary acids ($H\text{—}Y$)?

Across a period (left to right), acidity increases as electronegativity increases and the bond becomes more polarized; down a group (top to bottom), acidity increases as bond strength decreases.

What two factors determine the acid strength of oxyacids ($H\text{—}O\text{—}Y$)?

The electronegativity of the central element $Y$ (greater electronegativity pulls electron density away and weakens the $H\text{—}O$ bond) and the number of attached oxygen atoms (more electronegative oxygen atoms further weaken the $H\text{—}O$ bond).

How does the Lewis model define acids and bases?

A Lewis acid is an electron pair acceptor (must have an empty orbital or be capable of rearranging electrons), and a Lewis base is an electron pair donor.