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Chapter 15: Problem 3

Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain. How does the amount of each solution added change the effectiveness of the buffer?

Short Answer

Expert verified
Mixing acetic acid and sodium hydroxide forms a buffered solution consisting of acetic acid (weak acid) and sodium acetate (conjugate base), which help maintain a near constant pH. The balanced equation for the reaction is: \(CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O\). The buffer capacity depends on the concentrations of the weak acid and its conjugate base. Adjusting the amounts of acetic acid and sodium hydroxide solutions affects the concentration ratio, which in turn influences the buffer's effectiveness in maintaining a constant pH.

Step by step solution

01

1. Write the reaction equation

Write the balanced equation for the reaction between acetic acid and sodium hydroxide, forming sodium acetate and water. \[CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O\]
02

2. Understand the buffer system

In the reaction, acetic acid (weak acid) and sodium acetate (conjugate base) form a buffer system. This is because they can both donate or accept protons (H+) depending on whether an acid or base is added to the solution, helping to maintain the pH levels constant. For example, if an acid is added to the buffer system, the conjugate base \(CH_3COO^-\) will accept protons: \[CH_3COO^- + H^+ \rightarrow CH_3COOH\] If a base is added to the buffer system, the acetic acid will donate protons: \[CH_3COOH + OH^- \rightarrow CH_3COO^- + H_2O\]
03

3. Determine the buffer capacity

The buffer capacity is defined as the amount of acid or base that can be added to a buffered solution before its pH changes significantly. The buffer capacity depends on the concentrations of the weak acid and its conjugate base. A higher concentration of both the weak acid and its conjugate base will result in a higher buffer capacity, meaning it can resist pH changes more effectively.
04

4. Explore the effect of changing the amount of each solution

Changing the amount of acetic acid or sodium hydroxide in the solution will affect the buffer’s concentration ratio and therefore its effectiveness. If more acetic acid is added, the solution will become more acidic, reducing its ability to resist pH changes caused by adding more acid. If more sodium hydroxide is added, the solution will become more basic, making it less effective in resisting pH changes caused by adding a base. A balance between the concentration of acetic acid and sodium acetate is necessary for an effective buffer system that is able to resist pH changes.
05

5. Conclusion

Mixing solutions of acetic acid and sodium hydroxide can produce a buffered solution, with acetic acid and sodium acetate acting as the weak acid and conjugate base, respectively. Their concentrations determine the buffer capacity, and thus the effectiveness of the buffer in maintaining a constant pH. Adjusting the amounts of acetic acid and sodium hydroxide solutions will affect the concentration ratio of the weak acid and its conjugate base, which in turn influences the buffer's effectiveness.

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Most popular questions from this chapter

Th pH of blood is steady at a value of approximately 7.4 as a result of the following equilibrium reactions: $$ \mathrm{CO}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \leftrightharpoons \mathrm{H}_{2} \mathrm{CO}_{3}(a q) \leftrightharpoons \mathrm{HCO}_{3}-(a q)+\mathrm{H}^{+}(a q) $$ The actual buffer system in blood is made up of $\mathrm{H}_{2} \mathrm{CO}_{3}\( and \)\mathrm{HCO}_{3}$ - One way the body keeps the pH of blood at 7.4 is by regulating breathing. Under what blood ph conditions will the body increase breathing and under what blood pH conditions will the body decrease breathing? Explain.

A student titrates an unknown weak acid, HA, to a pale pink phenolphthalein end point with 25.0 \(\mathrm{mL}\) of 0.100\(M \mathrm{NaOH}\) . The student then adds 13.0 \(\mathrm{mL}\) of 0.100 \(\mathrm{M} \mathrm{HCl}\) . The pH of the resulting solution is \(4.70 .\) How is the value of \(\mathrm{p} K_{2}\) for the unknown acid related to 4.70\(?\)

Sketch the titration curves for a diprotic acid titrated by a strong base and a triprotic acid titrated by a strong base. List the major species present at various points in each curve. In each curve, label the halfway points to equivalence. How do you calculate the pH at these halfway points?

Tris (hydroxymethyl)aminomethane, commonly called TRIS or Trizma, is often used as a buffer in biochemical studies. Its buffering range is $\mathrm{pH} 7\( to \)9,\( and \)K_{\mathrm{b}}\( is \)1.19 \times 10^{-6}$ for the aqueous reaction $$ \left(\mathrm{HOCH}_{2}\right)_{3} \mathrm{CNH}_{2}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons\left(\mathrm{HOCH}_{2}\right)_{3} \mathrm{CNH}_{3}^{+}+\mathrm{OH}^{-} $$ a. What is the optimal pH for TRIS buffers? b. Calculate the ratio \([T R I S] /\left[T R I S H^{+}\right]\) at \(p H=7.00\) and at p H=9.00 c. A buffer is prepared by diluting 50.0 \(\mathrm{g}\) TRIS base and 65.0 \(\mathrm{g}\) TRIS hydrochloride (written as TRISHCl) to a total volume of 2.0 \(\mathrm{L}\) What is the pH of this buffer? What is the pH after 0.50 \(\mathrm{mL}\) of 12 \(\mathrm{MHCl}\) is added to a 200.0 -mL portion of the buffer?

A buffer solution is prepared by mixing 75.0 \(\mathrm{mL}\) of 0.275 \(\mathrm{M}\) fluorobenzoic acid $\left(\mathrm{C}_{7} \mathrm{H}_{5} \mathrm{O}_{2} \mathrm{F}\right)\( with 55.0 \)\mathrm{mL}$ of 0.472 \(\mathrm{M}\) sodium fluorobenzoate. The \(\mathrm{pK}_{\mathrm{a}}\) of this weak acid is \(2.90 .\) What is the pH of the buffer solution?
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