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Chapter 17: Problem 19

Which of the following solutions is a buffer? (a) \(0.20 \mathrm{M}\) formic acid (HCOOH), (b) \(0.20 \mathrm{M}\) formic acid \((\mathrm{HCOOH})\) and \(0.20 \mathrm{M}\) sodium formate (HCOONa), (c) \(0.20 \mathrm{M}\) nitric acid \(\left(\mathrm{HNO}_{3}\right)\) and \(0.20 \mathrm{M}\) sodium nitrate \(\left(\mathrm{NaNO}_{3}\right),\) (d) both b and \(\mathrm{c},(\mathbf{e})\) all of \(\mathrm{a}, \mathrm{b},\) and \(\mathrm{c}\)

Short Answer

Expert verified
The buffer is option (b): 0.20 M formic acid and 0.20 M sodium formate.

Step by step solution

01

Understanding Buffer Solutions

A buffer is a solution that resists changes in pH upon the addition of an acid or a base. Buffers typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid.
02

Analyzing Each Option for Buffer Composition

Look at each solution to see if it comprises a weak acid or base and its corresponding conjugate partner. - (a) contains only formic acid, which is a weak acid, with no conjugate base present. Thus, it is not a buffer. - (b) contains both formic acid (a weak acid) and sodium formate, which provides the conjugate base (formate ion). Thus, this combination can act as a buffer. - (c) contains nitric acid, a strong acid, and sodium nitrate. Since strong acids don't form buffers, this solution isn't a buffer.
03

Conclusion Based on Analysis

From our analysis, only option (b) forms a solution with both a weak acid (formic acid) and its conjugate base (formate ion), fulfilling the requirements for a buffer solution.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Weak Acid
A weak acid is an acid that partially dissociates into its ions in aqueous solution. This means not all the acid molecules release their hydrogen ions (H⁺) in water.

Because of this partial dissociation, weak acids have higher pH values when compared to strong acids of the same concentration. For example, formic acid ( HCOOH ) is a weak acid because it does not fully ionize in water.

Weak acids are essential in the formation of buffer solutions, as they can donate H⁺ ions when necessary without altering the pH significantly. This makes them effective in maintaining a stable pH environment.
Conjugate Base
In a chemical equilibrium, the term "conjugate base" refers to the species that remains after an acid has donated a proton. It's the counterpart of an acid.

For instance, in the dissociation of formic acid, the formate ion ( HCOO⁻ ) becomes the conjugate base. It plays a vital role in buffering solutions by accepting H⁺ ions when excess acid is added.

Conjugate bases are crucial to the effectiveness of buffer systems. They work hand-in-hand with their corresponding weak acids to neutralize added acids or bases, thereby resisting drastic pH changes.
Buffer Capacity
Buffer capacity is a measure of a buffer solution's ability to resist pH changes upon the addition of an acid or base.

A buffer solution operates efficiently within a certain range, with its capacity dependent on the concentrations of the weak acid and its conjugate base.

The higher these concentrations, the greater the buffer capacity. This means they can handle more significant amounts of added acids or bases without substantial pH shifts. It's crucial for many biological and chemical applications, such as maintaining the proper working environment in biological systems.
Formic Acid
Formic acid, with the chemical formula HCOOH , is the simplest carboxylic acid. It's a weak acid commonly found in nature, notably in the venom of ants and other insects.

Formic acid plays a significant role in science and industry, used in producing textiles, leather, and cleaning agents.

Because it acts as a weak acid, formic acid is valuable in buffer solutions, pairing with its salt counterpart, such as sodium formate, to create an effective pH-regulating environment.
Sodium Formate
Sodium formate ( HCOONa ) is the sodium salt of formic acid. In solution, it dissociates into sodium ions and formate ions ( HCOO⁻ ).

It's the conjugate base of formic acid and is integral in forming a buffer when combined with formic acid.

Sodium formate is not only instrumental in buffers but also in various applications, such as a de-icing agent and as an additive in drilling fluids in the oil and gas industry. Its role as a conjugate base is pivotal for maintaining stable pH levels in solutions.

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Most popular questions from this chapter

What is the \(\mathrm{pH}\) at \(25^{\circ} \mathrm{C}\) of water saturated with \(\mathrm{CO}_{2}\) at a partial pressure of \(111.5 \mathrm{kPa}\) ? The Henry's law constant for \(\mathrm{CO}_{2}\) at \(25^{\circ} \mathrm{C}\) is \(3.1 \times 10^{-4} \mathrm{~mol} / \mathrm{L}-\mathrm{kPa} .\)

A solution contains three anions with the following concentrations: \(0.20 \mathrm{MCrO}_{4}^{2-}, 0.10 \mathrm{MCO}_{3}^{2-},\) and \(0.010 \mathrm{MCl}^{-}\). If a dilute \(\mathrm{AgNO}_{3}\) solution is slowly added to the solution, what is the first compound to precipitate: \(\mathrm{Ag}_{2} \mathrm{CrO}_{4}\left(K_{4 p}=1.2 \times 10^{-12}\right), \mathrm{Ag}_{2} \mathrm{CO}_{3}\left(K_{4 p}=8.1 \times 10^{-12}\right)\) or \(\mathrm{AgCl}\left(K_{\mathrm{sp}}=1.8 \times 10^{-10}\right) ?\)

Salts containing the phosphate ion are added to municipal water supplies to prevent the corrosion of lead pipes. (a) Based on the \(\mathrm{pK}_{\mathrm{ad}}\) values for phosphoric acid \(\left(\mathrm{p} K_{\mathrm{at}}=7.5 \times 10^{-3}\right.\), \(\left.\mathrm{p} K_{a 2}=6.2 \times 10^{-8}, \mathrm{p} K_{a 3}=4.2 \times 10^{-13}\right)\) what is the \(\mathrm{K}_{\mathrm{b}}\) value for the \(\mathrm{PO}_{4}^{3-}\) ion? (b) What is the pH of a \(1 \times 10^{-3}\) \(M\) solution of \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) (you can ignore the formation of \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) and \(\left.\mathrm{H}_{3} \mathrm{PO}_{4}\right) ?\)

In nonaqueous solvents, it is possible to react HF to create \(\mathrm{H}_{2} \mathrm{~F}^{+} .\) Which of these statements follows from this observation? (a) HF can act like a strong acid in nonaqueous solvents, (b) HF can act like a base in nonaqueous solvents, (c) HF is thermodynamically unstable, \((\mathbf{d})\) There is an acid in the nonaqueous medium that is a stronger acid than HE.

Suppose that a \(10-\mathrm{mL}\) sample of a solution is to be tested for \(1^{-}\) ion by addition of 1 drop \(\left(0.2 \mathrm{~mL}\right.\) ) of \(0.10 \mathrm{M} \mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\). What is the minimum number of grams of \(I^{-}\) that must be present for \(\mathrm{Pbl}_{2}(s)\) to form?
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