Buffer Solution: Mixing Acids & Salts

Oct 16, 2025 by ADMIN 38 views

Hey guys! Let's dive into a cool chemistry problem involving buffer solutions. A student whipped up a buffer solution by mixing 0.1M $CH_3COOH$ (acetic acid) and 0.1 M $CH_3COONa$ (sodium acetate). We're also given that the $pKa$ of $CH_3COOH$ is 4.76. Our mission? To figure out which statements about this buffer are true. Sounds fun, right? Don't worry, we'll break it down step by step so it's super clear.

Understanding Buffer Solutions: The Basics

First off, what exactly is a buffer solution? Think of it like a chemical bodyguard. Its primary job is to resist changes in pH when you add a little bit of acid or base. This is super important in all sorts of situations, from our own blood to industrial processes. A buffer solution typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid. In our case, we've got acetic acid ( $CH_3COOH$ ), which is a weak acid, and its conjugate base, acetate ( $CH_3COO^-$ ), which comes from the sodium acetate ( $CH_3COONa$ ).

The secret to a buffer's power lies in the equilibrium between the weak acid and its conjugate base. When we add a strong acid (like $HCL$ ), the conjugate base in the buffer will react with it, neutralizing the added acid. Conversely, if we add a strong base (like $NaOH$ ), the weak acid in the buffer will neutralize it. This keeps the pH relatively stable, which is the magic behind how a buffer solution works. To really get a grasp of this concept, we need to think about the equilibrium and how each component interacts to maintain the balance. We're talking about dynamic equilibrium here, so it is a continuous process of reaction. That's why buffer solutions can stay at a stable pH.

To really nail this down, it's essential to grasp the acid dissociation constant, or Ka, and its negative logarithmic form, pKa. The Ka is a measure of how readily an acid donates a proton. pKa, on the other hand, is just a handy way to express Ka, making the numbers easier to work with. A lower pKa value indicates a stronger acid, and a higher pKa value indicates a weaker acid. In our scenario, the $pKa$ of $CH_3COOH$ is 4.76. This number tells us something about the strength of acetic acid.

Now, let's talk about the Henderson-Hasselbalch equation, which is our secret weapon for calculating the pH of a buffer solution. It’s a game-changer! The equation is: $pH = pKa + log rac{[A^-]}{[HA]}$ , where $[A^-]$ is the concentration of the conjugate base, and $[HA]$ is the concentration of the weak acid. This equation is super useful for making predictions about the pH. This equation helps us to calculate the pH of a buffer solution, given the $pKa$ of the weak acid and the concentrations of the acid and its conjugate base. In our case, since the concentrations of $CH_3COOH$ and $CH_3COONa$ are the same (0.1 M), the pH of the buffer will be close to the $pKa$ value.

Keep in mind that the effectiveness of a buffer is greatest when the concentrations of the weak acid and its conjugate base are roughly equal. This is because the buffer can neutralize similar amounts of both added acid and base. A well-designed buffer will resist changes in pH, holding its ground even when faced with external acid or base additions. Therefore, we will be able to answer the question using all these fundamental concepts, and understanding how the buffer resists change in pH.

Key Takeaway:

Buffer solutions resist changes in pH. They contain a weak acid and its conjugate base (or a weak base and its conjugate acid).
The $pKa$ of an acid indicates its strength.
The Henderson-Hasselbalch equation helps us calculate the pH of a buffer.

Analyzing the Statements about the Buffer

Alright, time to get down to the specifics. We have a buffer solution made from 0.1M $CH_3COOH$ and 0.1 M $CH_3COONa$ , and the $pKa$ of $CH_3COOH$ is 4.76. With all this in mind, let's analyze each statement about the buffer solution.

Statement Analysis

Let’s go through each statement and see if it holds up. To do this, we'll apply our knowledge of buffer solutions and use the Henderson-Hasselbalch equation to check our understanding. Remember, the key is the equilibrium between the weak acid ( $CH_3COOH$ ) and its conjugate base ( $CH_3COO^-$ ), and how they interact to maintain a stable pH. It's also important to consider the concentrations of the acid and the conjugate base, which in our case are equal. Equal concentrations are a great starting point for a buffer, and that's one of the main factors that make it effective.

We also should think about how adding an acid or base will change the reaction. Acids will be neutralized by the conjugate base, and bases will be neutralized by the weak acid. Therefore, the buffer will resist changes in pH. Keep in mind that we need to analyze each answer. So, take your time and read it carefully, making sure you fully understand what the statement is actually saying. This will help you to pick out the correct answers without making any mistakes, and guarantee your ability to answer any question related to buffer solutions.

When we're done analyzing each statement, we will be able to tell what is right and what is wrong. By understanding the fundamentals of buffer solutions, the Henderson-Hasselbalch equation, and how acids and bases interact, we can confidently identify the correct statements and understand why they are true. So, let’s get started and break down the specifics of each possible answer.

The Correct Statements:

Now, let's analyze each possible statement about the buffer solution.

"The pH of the solution is approximately 4.76." This statement is correct. Because the concentrations of the weak acid ( $CH_3COOH$ ) and its conjugate base ( $CH_3COONa$ ) are equal (0.1 M), the pH of the buffer solution is approximately equal to the $pKa$ of the weak acid. The Henderson-Hasselbalch equation tells us that when the concentrations of the acid and conjugate base are equal, the $log rac{[A^-]}{[HA]}$ term becomes log(1), which equals 0. Thus, $pH = pKa + 0$ . So, the pH is approximately 4.76.
"The buffer has a high buffer capacity." This statement is correct. Buffer capacity is the ability of a buffer to resist changes in pH upon the addition of an acid or base. Since the concentrations of the weak acid and its conjugate base are relatively high and equal, the buffer solution has a good capacity to neutralize added acid or base, therefore, this statement is right.
"The concentration of $CH_3COOH$ is equal to the concentration of $CH_3COO^-$ ." This statement is correct. The problem states that both $CH_3COOH$ and $CH_3COONa$ are 0.1 M. The $CH_3COONa$ dissociates in water to produce $CH_3COO^-$ ions. Therefore, at the beginning, we have the same concentration of both the weak acid and its conjugate base.
"If a small amount of strong acid is added, the pH will decrease slightly." This statement is correct. Buffer solutions resist changes in pH when acids or bases are added. However, adding a strong acid will consume some of the conjugate base ( $CH_3COO^-$ ), shifting the equilibrium and slightly decreasing the pH. The magnitude of the change will be small compared to what would happen in the absence of a buffer, but a decrease is still expected.
"If a small amount of strong base is added, the pH will increase slightly." This statement is correct. Adding a strong base consumes some of the weak acid ( $CH_3COOH$ ), which will shift the equilibrium and slightly increase the pH. Again, the change will be minimal due to the buffer's action.

Incorrect Statements

"The solution contains only $CH_3COOH$ molecules." This statement is incorrect. The solution contains both $CH_3COOH$ and $CH_3COO^-$ ions from the dissociation of $CH_3COONa$ .
"The pH of the solution is 7." This statement is incorrect. The pH of an acetic acid/acetate buffer is not neutral (pH 7). It is determined by the $pKa$ of the acetic acid and the relative concentrations of the acid and conjugate base. In this case, the pH is close to the $pKa$ (4.76).
"The solution is highly acidic." This statement is incorrect. While the pH is acidic, it is not highly acidic, as a pH of 4.76 is not as acidic as, for example, a pH of 1 or 2. Also, a buffer solution moderates the acidity.

Conclusion: Mastering Buffer Solutions

Alright, guys, we've made it through! We've successfully navigated the world of buffer solutions. We've seen how they work, how to calculate their pH, and how to predict their behavior. Remember, buffer solutions are super important, so understanding them is a win. Keep practicing, and you'll become a buffer expert in no time. If you have any questions, don’t hesitate to ask! Happy studying!