Book: UTC: 19 Flashcards
How the presence of a common ion suppresses a reaction that forms it.
When a common ion is present in the partial dissociation of an acid (or base) and something is added to the mix that contains this ion, it shifts the weak dissociation away from forming more of that common ion, thereby “fixing” the concentration of the ion and its corresponding acid / base.
Why the concentrations of buffer components must be high to minimize the change in pH due to addition of small amounts of H3O+ or OH-.
A buffer works because a large amount of the acidic component (HA) of the buffer consumes small amounts of added OH- and a large amount of the basic component consumes small amounts of added H3O+. This can be seen by rearranging the K_a expression for H3O+:
[H3O+] = K_a [HA] / [A-]
If [HA] and [A-] are very large, addition of H3O+ will not significantly change as a single buffer component goes up or down.
How buffer capacity depends on buffer concentration and on the pK_a of the acid component; why buffer range is within ±1 pH unit of the pK_a.
Buffer capacity is a measure of how much a buffer can withstand before it fails to maintain the pH of the solution. Failure occurs when one of the buffer components runs out, so since components are “used up” when H3O+ or OH- are added, the more concentrated the components are in the mix, the higher the buffer capacity.
If relative buffer concentrations are off from each other by a factor of 10, then pH of the buffer ends up at pK_a ±1. The further buffer concentrations are from each other, the poorer the buffer performs because of the Henderson-Hasselbalch equation.
Why the shapes of strong acid-strong base, weak acid-strong base, and weak base-strong acid titration curves differ.
For a strong acid-strong base curve, the pH increases slowly before and after the equivalence point and very quickly jumps from acidic to basic. Slow increase before is due to the acid being neutralized by the base; this doesn’t affect pH for a strong acid. After the equivalence point, pH rises asymptotically to reflect the pH of the added base as all of the acid has been neutralized and the solution as a whole is just a less and less dilute version of the base.
For the former group (weak acid-strong base and vice versa), there exists a buffer region before the equivalence point where the weak acid acts as a buffer, pushing pH towards the pKa of the weak acid until the equivalence point is reached, at which point the “buffer” has been used up.
How the pH at the equivalence point is determined by the species present; why the pH at the midpoint of the buffer region equals the pK_a of the acid.
For the titration of an acid, the number of moles of OH- added has to equal the number of moles of H3O+ originally present in the solution.
At the midpoint, [HA] = [A-], so pH = pK_a + log( 1 ) or just pH = pK_a.
The nature of an acid-base indicator as a conjugate acid-base pair with differently colored acidic and basic forms.
An indicate works because it is a conjugate acid-base pair that has different colored components, so we have:
pH = pK_a - 1, [In-] / [HIn] = 0.1 and HIn color is seen, but
pH = pK_a, then HIn and In- colors merge into an intermediate hue, and finally,
pH = pK_a + 1, [In-] / [HIn] = 10 and the base (In-) color is seen.
The distinction between equivalence point and end point in an acid-base titration.
The equivalence point of a titration occurs when the acid (or base) has been titrated / neutralized completely by the base (or acid), but the end point is just the point at which the indicator changes color. In theory, these are supposed to be the same point, but the indicator color change is visible, whereas the equivalence point is not a visible change inherently.
How the titration curve of a polyprotic acid has a buffer region and equivalence point for each ionizable proton.
Polyprotic acids are all weak, and therefore generate buffer regions during titration. They generate a number of buffer regions equal to the number of H’s or ionizable protons in the acid, so H2SO3, for example, has two buffer regions because its titration looks like this:
H2SO3 -> HSO3- -> SO3(2-).
How a slightly soluble ionic compound reaches equilibrium in water, expressed by an equilibrium (solubility-product) constant K_sp.
Slightly soluble compounds dissociate via the following:
AB ⇌ A- + B+, so it reaches an equilibrium when Q_sp = K_sp = [A-] [B+]. Everything here works like a normal problem with K and Q.
Why incomplete dissociation of an ionic compound means that calculated values for K_sp and solubility are approximations.
Because all compounds technically “slightly” dissociate, assuming either complete dissociation or no dissociation at all is always incorrect.
Why a common ion in a solution decreases the solubility of its compounds.
Since dissociation of a compound is essentially an equilibrium problem, we have this:
AB ⇌ A- + B+
If a common ion is added, then [A-] or [B+] increases, which means that overall the expression must shift left and therefore AB isn’t dissociating as much.
How pH affects the solubility of a compound that contains a weak-acid anion.
When a compound dissociates into ions and there are H3O+ or OH- present, the dissociated ions will bond with these ions and cause a rightward shift in the equation of compound ion dissociation: AB ⇌ A- + B+ leading to an effective increase in solubility for AB.
How precipitate formation depends on the relative values of Q_sp and K_sp.
If Q_sp = K_sp, then the solution is saturated and no change will occur. If Q_sp > K_sp, a precipitate will form until the remaining solution is saturated. If Q_sp < K_sp, no precipitate will form because the solution is unsaturated.
How selective precipitation and simultaneous equilibria are used to separate ions.
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How complex-ion formation occurs in steps and is characterized by an overall equilibrium (formation) constant, K_f.
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