Acid–Base Equilibria (5A) - Chem & Physics

1. What is the Lewis definition of an acid?

   An electron pair acceptor. Lewis theory defines acids as species that accept electron pairs to form coordinate bonds.

2. What is the relationship between conjugates: $K_a(\text{HA})K_b(\text{A}^-)$ at $25\,^{\circ}\text{C}$?

   $K_aK_b=K_w$. For conjugate pairs, the product of Ka and Kb equals Kw at 25°C.

3. What is the formula for $K_b$ for $\text{B} + \text{H}_2\text{O} \rightleftharpoons \text{BH}^+ + \text{OH}^-$?

   $K_b=\frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}$. Kb expresses the equilibrium constant for weak base protonation, excluding water as a pure liquid.

4. What is the Henderson–Hasselbalch equation for a buffer of $\text{HA}/\text{A}^-$?

   $pH=pK_a+\log\!\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$. The equation relates buffer pH to pKa and the ratio of conjugate base to acid.

5. Identify the major species at $pH=pK_a-2$ for a monoprotic acid system $\text{HA}/\text{A}^-$.

   $\text{HA}$ predominates. At pH = pKa - 2, [A-]/[HA] = 0.01, so the protonated form is major.

6. What is the approximate buffer range (in pH units) around $pK_a$ for effective buffering?

   $pK_a\pm^1$. Effective buffering occurs when [A-]/[HA] is between 0.1 and 10, spanning pKa ±1.

7. Which side is favored for $\text{HA}+\text{B}^-\rightleftharpoons \text{A}^-+\text{HB}$ when $pK_a(\text{HB})>pK_a(\text{HA})$?

   Products (equilibrium favors the weaker acid). When pKa(HB) > pKa(HA), HB is weaker, so equilibrium shifts to the weaker acid side.

8. What is the conjugate base of the acid $\text{HA}$?

   $\text{A}^-$. The conjugate base forms by removing a proton from the acid HA.

9. In a buffer, what is $pH$ when $[\text{A}^-]=[\text{HA}]$?

   $pH=pK_a$. Equal concentrations give log(1) = 0, so pH equals pKa at half-equivalence.

10. What is the formula for percent ionization of a weak acid $\text{HA}$ in terms of $[\text{H}^+]$ and $[\text{HA}]_0$?

    $\%\text{ionization}=\frac{[\text{H}^+]}{[\text{HA}]_0}\times100\%$. Percent ionization quantifies dissociation extent using equilibrium [H+] over initial [HA].

11. What is the relationship between $pOH$ and $[\text{OH}^-]$?

    $pOH=-\log([\text{OH}^-])$. pOH measures basicity as the negative logarithm of hydroxide ion concentration.

12. At $25\,^{\circ}\text{C}$, what is the relationship between $pH$ and $pOH$?

    $pH+pOH=14$. At 25°C, pH + pOH equals pKw = 14 from water's ion product.

13. Identify the major species at $pH=pK_a+2$ for a monoprotic acid system $\text{HA}/\text{A}^-$.

    $\text{A}^-$ predominates. At pH = pKa + 2, [A-]/[HA] = 100, so the deprotonated form is major.

14. What is the Lewis definition of a base?

    An electron pair donor. Lewis theory defines bases as species that donate electron pairs to form coordinate bonds.

15. At $25\,^{\circ}\text{C}$, what is $[\text{H}^+]$ in pure water?

    $[\text{H}^+]=1.0\times10^{-7}\,\text{M}$. In pure water at 25°C, [H+] = [OH-] from autoionization, yielding neutral pH 7.

16. What is the relationship between $pK_a$ and $K_a$?

    $pK_a=-\log(K_a)$. pKa is the negative logarithm of Ka, indicating acid strength inversely.

17. What is the relationship between $pK_b$ and $K_b$?

    $pK_b=-\log(K_b)$. pKb is the negative logarithm of Kb, indicating base strength inversely.

18. What is the Brønsted–Lowry definition of a base?

    A proton ($\text{H}^+$) acceptor. Brønsted–Lowry theory defines bases as species that accept protons from acids.

19. What is the value of $K_w$ at $25\,^{\circ}\text{C}$?

    $K_w=1.0\times10^{-14}$. Kw is the ion product of water, constant at 25°C due to autoionization equilibrium.

20. What is the formula for $K_a$ for $\text{HA} + \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{A}^-$?

    $K_a=\frac{[\text{H}_3\text{O}^+][\text{A}^-]}{[\text{HA}]}$. Ka expresses the equilibrium constant for weak acid dissociation, excluding water as a pure liquid.

21. What is the net ionic equation for neutralization of a strong acid by a strong base in water?

    $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$. Strong acid-base neutralization produces water by complete reaction of H+ and OH- ions.

22. What is the relationship between $pH$ and $[\text{H}^+]$?

    $pH=-\log([\text{H}^+])$. pH measures acidity as the negative logarithm of hydrogen ion concentration.

23. What is the Brønsted–Lowry definition of an acid?

    A proton ($\text{H}^+$) donor. Brønsted–Lowry theory defines acids as species that donate protons in reactions with bases.

24. What is the relationship between conjugates: $pK_a(\text{HA})+pK_b(\text{A}^-)$ at $25\,^{\circ}\text{C}$?

    $pK_a+pK_b=14$. For conjugate pairs, pKa + pKb equals pKw = 14 at 25°C.

25. What is the conjugate acid of the base $\text{B}$?

    $\text{BH}^+$. The conjugate acid forms by adding a proton to the base B.

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