Question

Consider the following acids and bases: Choose substances from the following list that would be the best choice to prepare a \(\mathrm{pH}=9.0\) buffer solution. a. \(\mathrm{HCO}_{2} \mathrm{H}\) b. HOBr c. \(\mathrm{KHCO}_{2}\) d. \(\mathrm{HONH}_{3} \mathrm{NO}_{3}\) \(\mathbf{e} .\left(\mathbf{C}_{2} \mathbf{H}_{5}\right)_{2} \mathrm{NH}\) f. \(\left(C_{2} H_{5}\right)_{2} N H_{2} C l\) g. \(\mathrm{HONH}_{2}\) h. NaOBr

Short answer

The best combination to prepare a buffer solution at pH 9.0 would be HOBr (Hypobromous Acid) and NaOBr (Sodium Hypobromite).

Step-by-step solution

Identify Acids and Bases

We must first identify which compounds are acids and which are bases from the list: a. \(\mathrm{HCO}_{2}\mathrm{H}\) - Acid (Formic Acid) b. HOBr - Acid (Hypobromous Acid) c. \(\mathrm{KHCO}_{2}\) - Base (Potassium Formate) d. \(\mathrm{HONH}_{3}\mathrm{NO}_{3}\) - Acid (Hydroxylammonium Nitrate) e. \((\mathrm{C}_{2}\mathrm{H}_{5})_{2}\mathrm{NH}\) - Base (Diethylamine) f. \((\mathrm{C}_{2}\mathrm{H}_{5})_{2}\mathrm{NH}_{2}\mathrm{Cl}\) - Acid/Base (Diethylammonium chloride, a salt of the acid and conjugate base) g. \(\mathrm{HONH}_{2}\) - Acid (Hydroxylamine) h. NaOBr - Base (Sodium Hypobromite)

Calculate pKa Values

Next, we need to find the pKa values of the acids. The pKa values have been tabulated and can be found in any reference material (such as a textbook or online resource): a. \(\mathrm{HCO}_{2}\mathrm{H}\) - pKa = 3.75 b. HOBr - pKa = 8.6 d. \(\mathrm{HONH}_{3}\mathrm{NO}_{3}\) - pKa = 6.1 g. \(\mathrm{HONH}_{2}\) - pKa = 6.0

Select Suitable Buffer Pair

To create a buffer at pH 9.0, we need to find the acid and its conjugate base combination that works best. The appropriate pKa value should be close to the target pH level: HOBr (pKa = 8.6) is closest to pH 9.0 and is the best option among the weak acids. The conjugate base for HOBr is OBr⁻, which corresponds to NaOBr in the list.

Conclusion

Therefore, the best combination to prepare a buffer solution at pH 9.0 would be HOBr (Hypobromous Acid) and NaOBr (Sodium Hypobromite).

Key concepts

Acid-Base Chemistry

Acid-base chemistry is a fundamental concept in chemistry that revolves around the behaviors of acids and bases. Acids are substances that can donate a proton (H⁺) in a reaction, while bases are substances that can accept a proton. This reaction is essential in determining the characteristics of a substance, such as its pH level.

In the given exercise, we were presented with several compounds, and our first task was to identify whether they were acids or bases. For instance, compounds like \(HCO_2H\) (formic acid) and HOBr (hypobromous acid) are acids. Conversely, bases such as \(KHCO_2\) (potassium formate) and NaOBr (sodium hypobromite) act as proton acceptors.

Understanding acid-base chemistry helps us make informed choices when creating buffer solutions, which maintain a consistent pH level when small amounts of acid or base are added. This principle is essential in various applications, such as pharmaceuticals and biological systems.

pH Calculation

The process of pH calculation is crucial in determining the acidity or basicity of a solution. pH is a scale used to specify the acidity or basicity of an aqueous solution, calculated using the formula \[ \text{pH} = -\log[H^+] \] where \( H^+ \) represents the concentration of hydrogen ions.

In buffer solutions, the objective is to maintain a specific pH. For the buffer solution given in the exercise, the desired pH is 9.0. This target guides the selection of appropriate acids and bases that will work together to stabilize that pH level. A good buffer should have a pKa value close to the desired pH. This allows the buffer to resist changes in pH when small amounts of acid or base are introduced.

Accurate pH calculation not only underpins buffer preparation but also informs many aspects of chemical reactions and solutions, ensuring experimental and application precision.

pKa Values

The pKa value of an acid is a critical parameter in acid-base chemistry, as it measures the strength of an acid in a solution. It is the negative logarithm of the acid dissociation constant (Ka), and it provides insight into how easily an acid donates a proton. Mathematically, it is expressed as \[ pKa = -\log(Ka) \]

In the context of buffer preparation, knowing the pKa is invaluable because it allows you to choose acids and conjugate bases that will effectively stabilize the pH you aim to achieve. For example, in the exercise, HOBr's pKa is 8.6, making it close to the target pH of 9.0, which is suitable for creating a buffer at that pH level.

Thus, understanding pKa values aids in not only choosing the right components for buffer solutions but also predicting the behavior of acids in various chemical environments.

Conjugate Acid-Base Pairs

Conjugate acid-base pairs are two species that differ by a single proton. When an acid donates a proton, it forms its conjugate base, and when a base accepts a proton, it becomes its conjugate acid. This concept is pivotal in maintaining equilibrium within buffer solutions.

In our specific buffer scenario from the exercise, HOBr and its conjugate base \(\text{OBr}^-\) form such a pair. This pairing, along with its partner compound NaOBr (sodium hypobromite), works together to stabilize the solution close to the target pH of 9.0.

Understanding conjugate acid-base pairs allows for the effective design of buffer solutions by selecting pairs that can maintain the desired pH range. This concept applies broadly across chemistry, including biological systems where homeostasis is critical.