Problem 2 - Chemistry and Chemical Reactivity

Question

Does the pH of the solution increase, decrease, or stay the same when you (a) Add solid sodium oxalate, \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4},\) to \(50.0 \mathrm{mL}\) of \(0.015 \mathrm{M}\) oxalic acid, \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4} ?\) (b) Add solid ammonium chloride to 75 mL of \(0.016 \mathrm{M}\) HCl? (c) Add \(20.0 \mathrm{g}\) of \(\mathrm{NaCl}\) to \(1.0 \mathrm{L}\) of \(0.10 \mathrm{M}\) sodium acetate, \(\mathrm{NaCH}_{3} \mathrm{CO}_{2} ?\)

Step-by-Step Solution

Answer

(a) Increase (b) Stay the same (c) Stay the same

Step 1: Understand the Chemistry Concepts

To answer these questions, we first need to know that oxalic acid ( H_2C_2O_4 ) is a diprotic acid, meaning it can donate two protons (H+ ions), while sodium oxalate ( Na_2C_2O_4 ) is its conjugate base. Ammonium chloride ( NH_4Cl ) in solution provides chloride ions and ammonium ions, the latter of which can release protons and affect the pH. Sodium chloride ( NaCl ) is a neutral salt with no effect on pH, although it can affect ionic strength when added to solutions involving other ions.

Step 1: Calculate the concentration changes: A

Adding sodium oxalate to oxalic acid introduces its conjugate base, shifting the equilibrium:\[ H_2C_2O_4 \rightleftharpoons H^+ + HC_2O_4^- \rightleftharpoons 2H^+ + C_2O_4^{2-} \]Adding Na_2C_2O_4 provides C_2O_4^{2-} ions, lowering the H^+ concentration by shifting equilibrium left, thus increasing the pH.

Step 2: Consider the equilibrium shift: B

Adding solid NH_4Cl to HCl doesn't change pH significantly, as HCl is a strong acid that dissociates completely. Ammonium ion ( NH_4^+ ) can contribute slightly to the H+ concentration, but it is negligible compared to the HCl.

Step 3: Look at ionic strength and common ion: C

NaCl added to sodium acetate does not contribute extra H^+ or OH^- ions because it is a neutral salt. The common ion effect from acetate might be insignificant due to salt's neutrality and the large volume of water, so the pH stays approximately constant.

Key Concepts

Oxalic Acid

Oxalic acid, \(H_{2}C_{2}O_{4}\), is a naturally occurring organic acid found in various plants, including spinach and rhubarb. It is classified as a diprotic acid because it can donate two protons, or hydrogen ions, per molecule. This proton donation capability allows it to participate in various chemical equilibria.
The dissociation of oxalic acid occurs in two steps:

• In the first dissociation, oxalic acid loses a hydrogen ion to form the hydrogen oxalate ion \( HC_{2}O_{4}^{-} \).
• In the second step, the hydrogen oxalate ion can further lose a hydrogen ion to form the oxalate ion \( C_2O_{4}^{2-} \).

The overall reactions are as follows:

• Step 1: \(H_{2}C_{2}O_{4} ightleftharpoons H^{+} + HC_{2}O_{4}^{-}\)
• Step 2: \(HC_{2}O_{4}^{-} ightleftharpoons H^{+} + C_{2}O_{4}^{2-}\)

By understanding these equilibria, we can analyze how the addition of the conjugate base, sodium oxalate, will influence the solution's pH.

Conjugate Base

The concept of conjugate base is essential when dealing with acids like oxalic acid. A conjugate base is formed when an acid donates a proton. For oxalic acid, the conjugate bases are hydrogen oxalate, \( HC_{2}O_{4}^{-} \), and oxalate, \( C_{2}O_{4}^{2-} \).
Upon addition of a conjugate base into a solution, the acid-base equilibrium shifts. This shift can be understood by applying Le Chatelier's principle:

• Adding the conjugate base, sodium oxalate (which provides \( C_{2}O_{4}^{2-} \)), to a solution of oxalic acid shifts the equilibrium to the left.
• This shift occurs because the system attempts to decrease the concentration of \( C_{2}O_{4}^{2-} \) by converting some of it into \( HC_{2}O_{4}^{-} \) and \( H_{2}C_{2}O_{4} \).
• As a result, the overall hydrogen ion concentration \( [H^{+}] \) decreases, causing an increase in the pH of the solution.

This concept is at the heart of buffer solutions, where a balance between the weak acid and its conjugate base is maintained to resist changes in pH.

Ammonium Chloride

Ammonium chloride, \( NH_{4}Cl \), is a salt derived from the reaction of ammonia (a weak base) and hydrochloric acid (a strong acid). When dissolved in water, it dissociates into ammonium ions \( NH_{4}^{+} \) and chloride ions \( Cl^{-} \).
The presence of ammonium ions can slightly affect the pH of a solution. However, when mixed with a strong acid like HCl, this influence is negligible due to the complete dissociation of HCl:

• A strong acid fully dissociates in solution, providing a large concentration of \( H^{+} \) ions.
• While the \( NH_{4}^{+} \) ion can release a proton to slightly increase the hydrogen ion concentration, its effect is overshadowed by HCl's dominant contribution.

Therefore, the addition of ammonium chloride to hydrochloric acid solution does not significantly alter the pH due to the preexisting overwhelming presence of \( H^{+} \) ions.

Sodium Chloride Effect

Sodium chloride, \( NaCl \), is known for its role as a neutral salt in solution. Unlike other salts, NaCl does not contribute to changing the pH because neither sodium \( Na^{+} \) nor chloride \( Cl^{-} \) ions picks up or donates protons in solution.
When added to a solution with other ionizable compounds, NaCl can increase the ionic strength of the solution:

• This increased ionic strength can influence the behavior of other ions present by altering activity coefficients, which can influence reactions' equilibrium positions.
• In the context of sodium acetate solution, NaCl's presence mainly changes ionic strength without affecting pH significantly, since sodium acetate's acetate ion \( CH_{3}COO^{-} \) does not share a common ion effect with the neutral Cl ion.

Thus, the addition of sodium chloride typically leaves the pH unchanged, although it might influence the solution in subtle ways related to ionic interactions.