Ionic Equilibrium (11th) flash cards

1. 1.State the Arrhenius definition of an acid and a base.

   An acid is a substance that gives $\text{H}^+$ (protons) in aqueous solution; a base gives $\text{OH}^-$ ions in aqueous solution. E.g. $\text{HCl} \to \text{H}^+ + \text{Cl}^-$ and $\text{NaOH} \to \text{Na}^+ + \text{OH}^-$.

   Hint: Think in terms of $\text{H}^+$ and $\text{OH}^-$ in water only.

2. 2.What is the main limitation of the Arrhenius concept?

   It is restricted to aqueous solutions and cannot explain acidic/basic behaviour in non-aqueous or gaseous media (e.g. $\text{NH}_3 + \text{HCl}$ in gas phase), nor basicity of substances like $\text{NH}_3$ that contain no $\text{OH}^-$.

   Hint: Water-bound; can't handle $\text{NH}_3$ as a base without $\text{OH}^-$.

3. 3.State the Bronsted-Lowry definition of acids and bases.

   An acid is a proton donor and a base is a proton acceptor. Acid-base reactions are proton-transfer reactions and are not limited to water.

   Hint: Donor vs acceptor of $\text{H}^+$.

4. 4.What is a conjugate acid-base pair?

   Two species differing by a single proton. When an acid $\text{HA}$ loses $\text{H}^+$ it forms its conjugate base $\text{A}^-$; when a base $\text{B}$ gains $\text{H}^+$ it forms its conjugate acid $\text{BH}^+$.

   Hint: Differ by exactly one $\text{H}^+$.

5. 5.In $\text{HCl} + \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{Cl}^-$, identify the conjugate pairs.

   $\text{HCl}/\text{Cl}^-$ is one conjugate acid-base pair (HCl acid, $\text{Cl}^-$ its conjugate base). $\text{H}_3\text{O}^+/\text{H}_2\text{O}$ is the other (water is the base, $\text{H}_3\text{O}^+$ its conjugate acid).

   Hint: Water accepts the proton here.

6. 6.What does it mean that water is amphoteric (amphiprotic)?

   Water can act as both an acid and a base. With $\text{HCl}$ it accepts a proton (base); with $\text{NH}_3$ it donates a proton (acid): $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$.

   Hint: It can both donate and accept $\text{H}^+$.

7. 7.State the Lewis definition of acids and bases.

   A Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. This is the most general definition and needs no proton or solvent.

   Hint: Electron-pair acceptor vs donor.

8. 8.Give typical examples of Lewis acids.

   Electron-deficient species: $\text{BF}_3$, $\text{AlCl}_3$, $\text{FeCl}_3$; cations like $\text{H}^+$, $\text{Ag}^+$, $\text{Cu}^{2+}$; molecules with vacant $d$-orbitals like $\text{SiF}_4$.

   Hint: Incomplete octet or empty orbital, seeks electrons.

9. 9.Give typical examples of Lewis bases.

   Electron-rich species with lone pairs: $\text{NH}_3$, $\text{H}_2\text{O}$, $\text{OH}^-$, $\text{F}^-$, $\text{CN}^-$, and molecules like ethers and amines.

   Hint: Has a lone pair to donate.

10. 10.Relate the strength of an acid to the strength of its conjugate base.

    A strong acid has a weak conjugate base, and a weak acid has a relatively strong conjugate base. Strength is inversely related: $K_a \times K_b = K_w$ for a conjugate pair.

    Hint: Stronger acid $\Rightarrow$ weaker its conjugate base.

11. 11.Why is $\text{Cl}^-$ a very weak base while $\text{CH}_3\text{COO}^-$ is a stronger base?

    $\text{HCl}$ is a strong acid, so its conjugate base $\text{Cl}^-$ has essentially no tendency to accept protons. $\text{CH}_3\text{COOH}$ is weak, so its conjugate base $\text{CH}_3\text{COO}^-$ readily accepts protons and hydrolyses.

    Hint: Conjugate of strong acid is a spectator; conjugate of weak acid is basic.

12. 12.What distinguishes a strong electrolyte from a weak electrolyte?

    A strong electrolyte ionizes almost completely in water (e.g. $\text{HCl}$, $\text{NaOH}$, $\text{NaCl}$); a weak electrolyte ionizes only partially and sets up an equilibrium (e.g. $\text{CH}_3\text{COOH}$, $\text{NH}_4\text{OH}$).

    Hint: Complete vs partial ionization.

13. 13.Write the ionization equilibrium and $K_a$ expression for a weak monoprotic acid HA.

    $\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$, with $K_a = \dfrac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$. Larger $K_a$ means a stronger acid.

    Hint: Products over reactant, water omitted.

14. 14.Write the ionization equilibrium and $K_b$ expression for a weak base B.

    $\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^-$, with $K_b = \dfrac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}$. Larger $K_b$ means a stronger base.

    Hint: Water is solvent, left out of $K_b$.

15. 15.Define $pK_a$ and $pK_b$.

    $pK_a = -\log K_a$ and $pK_b = -\log K_b$. A smaller $pK_a$ means a stronger acid; a smaller $pK_b$ means a stronger base.

    Hint: Negative log of the constant; small $p$ = strong.

16. 16.For a conjugate acid-base pair, how are $K_a$ and $K_b$ related?

    $K_a \times K_b = K_w = 10^{-14}$ at $25^\circ\text{C}$, which gives $pK_a + pK_b = 14$.

    Hint: Multiply constants to get $K_w$.

17. 17.Define the degree of ionization (dissociation) $\alpha$.

    $\alpha = \dfrac{\text{number of molecules ionized}}{\text{total number of molecules dissolved}}$. It ranges from 0 to 1 (or as a percentage) and increases on dilution for weak electrolytes.

    Hint: Fraction that actually splits into ions.

18. 18.State Ostwald's dilution law for a weak acid.

    For $\text{HA}$ of concentration $C$ and degree of ionization $\alpha$: $K_a = \dfrac{C\alpha^2}{1-\alpha}$. If $\alpha \ll 1$, then $K_a \approx C\alpha^2$ so $\alpha \approx \sqrt{K_a/C}$.

    Hint: $K_a = C\alpha^2/(1-\alpha)$; approximate for small $\alpha$.

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