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Chapter 17: Problem 21

\(\begin{array}{lll}\text { Given } & 1.00 & \mathrm{L}\end{array}\) of a solution that is \(0.100 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COOH}\) and \(0.100 \mathrm{M} \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COO}\) (a) Over what pH range will this solution be an effective buffer? (b) What is the buffer capacity of the solution? That is, how many millimoles of strong acid or strong base can be added to the solution before any significant change in pH occurs?

Short Answer

Expert verified
The effective pH range of the buffer solution would be between 3.82 and 5.82. The buffer capacity of the solution is 1 mol/L.

Step by step solution

01

Determine the pH range

First, it is important to know that a buffer is most effective when pH is within 1 unit of the pKa of the buffering system. The acid mentioned, CH3CH2COOH ethanol, is a weak acid. Its pKa value is 4.82. So, the effective buffer range of the solution would be the pKa ±1, giving a pH range of 3.82 to 5.82.
02

Calculate the buffer capacity

Buffer capacity refers to the amount of acid or base a buffer can neutralize before the pH begins to change to an appreciable degree. Using the buffer capacity formula, we know \(Buffer Capacity = 0.5 * Volume * (10^(pH-pKa) + 10^(pKa-pH))\). The pKa of ethanol is 4.82, the volume is 1.00 L, and we want to determine the buffer capacity at pH 4.82 (the pKa value). So, \(Buffer Capacity = 0.5 * 1 * (1 + 1) = 1 \). So, the buffer can neutralize up to 1 mol of strong acid or 1 mol of strong base. Therefore, the buffer capacity of the solution is 1 mol/L.

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Most popular questions from this chapter

In the titration of \(20.00 \mathrm{mL}\) of \(0.175 \mathrm{M} \mathrm{NaOH},\) calculate the number of milliliters of \(0.200 \mathrm{M} \mathrm{HCl}\) that must be added to reach a pH of (a) \(12.55,\) (b) \(10.80,\) (c) 4.25

Two solutions are mixed: \(100.0 \mathrm{mL}\) of \(\mathrm{HCl}(\mathrm{aq})\) with \(\mathrm{pH} 2.50\) and \(100.0 \mathrm{mL}\) of \(\mathrm{NaOH}(\text { aq) with } \mathrm{pH} 11.00\) What is the pH of the resulting solution?

What concentration of ammonia, \(\left[\mathrm{NH}_{3}\right],\) should be present in a solution with \(\left[\mathrm{NH}_{4}^{+}\right]=0.732 \mathrm{M}\) to produce a buffer solution with \(\mathrm{pH}=9.12 ?\) For \(\mathrm{NH}_{3}\) \(K_{\mathrm{h}}=1.8 \times 10^{-5}\)

Both sodium hydrogen carbonate (sodium bicarbonate) and sodium hydroxide can be used to neutralize acid spills. What is the pH of \(1.00 \mathrm{M} \mathrm{NaHCO}_{3}(\mathrm{aq})\) and of \(1.00 \mathrm{M} \mathrm{NaOH}(\mathrm{aq}) ?\) On a per-liter basis, do these two solutions have an equal capacity to neutralize acids? Explain. On a per-gram basis, do the two solids, \(\mathrm{NaHCO}_{3}(\mathrm{s})\) and \(\mathrm{NaOH}(\mathrm{s}),\) have an equal capacity to neutralize acids? Explain. Why do you suppose that \(\mathrm{NaHCO}_{3}\) is often preferred to \(\mathrm{NaOH}\) in neutralizing acid spills?

What stoichiometric concentration of the indicated substance is required to obtain an aqueous solution with the pH value shown: (a) \(\mathrm{Ba}(\mathrm{OH})_{2}\) for \(\mathrm{pH}=11.88 ;(\mathrm{b})\) \(\mathrm{CH}_{3} \mathrm{COOH}\) in \(0.294 \mathrm{M} \mathrm{NaCH}_{3} \mathrm{COO}\) for \(\mathrm{pH}=4.52 ?\)
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