Respan Dataset Explorer

The buffer capacity of a buffer solution would not be affected by decreasing the concentration of the base component.


	True
	False

**False.**

The buffer capacity of a buffer solution is directly related to the concentrations of the acid and base components in the buffer. Decreasing the concentration of the base component would reduce the buffer capacity, meaning the solution would be less able to resist changes in pH upon the addition of acids or bases. Therefore, the statement is false.

The buffering capacity of each buffer in this experiment will be determined using pH indicators.


	True
	False

**False.**

While pH indicators can help visualize changes in pH and identify the approximate pH range of a buffer, they do not determine the buffering capacity itself. Buffering capacity is typically assessed by directly measuring the pH changes upon the addition of a strong acid or strong base to the buffer solution and calculating how much acid or base can be added before the pH changes significantly. Buffers are characterized by their ability to maintain pH within a certain range, but this capacity is quantified through experimental measurements rather than solely through pH indicators.

A larger buffer capacity means that:


	more buffer is needed to neutralize a strong base
	more buffer is needed to neutralize a strong acid
	the buffer is less resistant to changes in pH
	the buffer is more resistant to changes in pH

A larger buffer capacity means that:

**the buffer is more resistant to changes in pH.**

This means that a buffer with higher capacity can effectively neutralize larger amounts of strong acids or bases without experiencing significant changes in pH. Therefore, the correct statement is that the buffer is more resistant to changes in pH.

A buffer solution can be made from a weak acid and its conjugate base.


	True
	False

**True.**

A buffer solution can be made from a weak acid and its conjugate base. This combination allows the buffer to effectively resist changes in pH when small amounts of strong acids or strong bases are added. The weak acid can react with added bases, while the conjugate base can react with added acids, thus maintaining the pH of the solution.

A buffer solution can be made from a weak base and its conjugate acid.


	True
	False

**True.**

A buffer solution can indeed be made from a weak base and its conjugate acid. This combination also allows the buffer to resist changes in pH when small amounts of strong acids or bases are added. The weak base can neutralize added acids, while the conjugate acid can neutralize added bases, thus maintaining the pH of the solution.

In the equation:

HA ⇔ H+ + A-

the buffer ratio is [A-] / [H+]


	True
	False

**False.**

In the equation:

\[ HA ⇔ H^+ + A^- \]

the buffer ratio should be expressed as:

\[
\text{Buffer Ratio} = \frac{[A^-]}{[HA]}
\]

where \([A^-]\) is the concentration of the conjugate base and \([HA]\) is the concentration of the weak acid. The ratio of \([A^-]\) to \([H^+]\) does not represent the buffer capacity or the buffer ratio in a typical buffer system.

A buffer solution maintains a relatively steady H+ concentration by shifting the equilibrium between a conjugate acid-base pair.


	True
	False

**True.**

A buffer solution maintains a relatively steady \( \text{H}^+ \) concentration (and therefore a stable pH) by utilizing the equilibrium between a conjugate acid-base pair. When an acid or base is added to the buffer, the equilibrium shifts to either consume the added acid (using the conjugate base) or the added base (using the conjugate acid), thereby minimizing changes in \( \text{H}^+ \) concentration and maintaining stability in pH.

A buffer is prepared using acetic acid, CH3COOH, (a weak acid) and sodium acetate, CH3COONa (which provides acetate ions, the conjugate base), according to the following proportions:

 

Volume of CH3COOH(aq):   113.0 mL

Concentration of CH3COOH(aq):  1.233 M

 

Volume of CH3COONa(aq):  111.0 mL

Concentration of CH3COONa(aq):  1.272 M

 

What is the concentration of the conjugate acid, CH3COOH, component in the final buffer solution?

Give your answer in M, with correct significant digits.

 

Number
 M

To calculate the concentration of the conjugate acid (\( \text{CH}_3\text{COOH} \)) component in the final buffer solution, we need to determine the moles of acetic acid and the total volume of the buffer solution.

### Step 1: Calculate moles of \( \text{CH}_3\text{COOH} \)

Using the volume and concentration of acetic acid:

\[
\text{Moles of } \text{CH}_3\text{COOH} = \text{Concentration} \times \text{Volume}
\]
\[
\text{Moles of } \text{CH}_3\text{COOH} = 1.233 \, \text{M} \times 0.113 \, \text{L} = 0.139 \, \text{moles}
\]

### Step 2: Calculate moles of \( \text{CH}_3\text{COONa} \)

Using the volume and concentration of sodium acetate:

\[
\text{Moles of } \text{CH}_3\text{COONa} = \text{Concentration} \times \text{Volume}
\]
\[
\text{Moles of } \text{CH}_3\text{COONa} = 1.272 \, \text{M} \times 0.111 \, \text{L} = 0.141 \, \text{moles}
\]

### Step 3: Calculate the total volume of the buffer solution

\[
\text{Total Volume} = \text{Volume of } \text{CH}_3\text{COOH} + \text{Volume of } \text{CH}_3\text{COONa}
\]
\[
\text{Total Volume} = 113.0 \, \text{mL} + 111.0 \, \text{mL} = 224.0 \, \text{mL} = 0.224 \, \text{L}
\]

### Step 4: Calculate the concentration of \( \text{CH}_3\text{COOH} \) in the final buffer

\[
\text{Concentration of } \text{CH}_3\text{COOH} = \frac{\text{Moles of } \text{CH}_3\text{COOH}}{\text{Total Volume}}
\]
\[
\text{Concentration of } \text{CH}_3\text{COOH} = \frac{0.139 \, \text{moles}}{0.224 \, \text{L}} \approx 0.6205 \, \text{M}
\]

### Final Answer

The concentration of the conjugate acid \( \text{CH}_3\text{COOH} \) in the final buffer solution is:

**0.620 M** (rounded to three significant figures based on the significant figures of the given volumes and concentrations).

A buffer is prepared using acetic acid, CH3COOH, (a weak acid, pKa = 4.75) and sodium acetate, CH3COONa (which provides acetate ions, the conjugate base), according to the following proportions:

 

Volume of CH3COOH(aq):   137.0 mL

Concentration of CH3COOH(aq):  1.189 M

 

Volume of CH3COONa(aq):  112.0 mL

Concentration of CH3COONa(aq):  1.127 M

 

Calculate the theoretical pH of this buffer.

	
4.73

	
4.77

	
4.64

	
4.86

	
4.75

To calculate the theoretical pH of the buffer solution made from acetic acid and sodium acetate, we can use the Henderson-Hasselbalch equation:

\[
\text{pH} = \text{pKa} + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)
\]

where:
- \(\text{pKa} = 4.75\) (given for acetic acid)
- \([\text{A}^-]\) is the concentration of the conjugate base (sodium acetate).
- \([\text{HA}]\) is the concentration of the weak acid (acetic acid).

### Step 1: Calculate moles of \( \text{CH}_3\text{COOH} \) (weak acid)

Using the volume and concentration of acetic acid:

\[
\text{Moles of } \text{CH}_3\text{COOH} = \text{Concentration} \times \text{Volume} = 1.189 \, \text{M} \times 0.137 \, \text{L}
\]
\[
\text{Moles of } \text{CH}_3\text{COOH} = 0.1636 \, \text{moles}
\]

### Step 2: Calculate moles of \( \text{CH}_3\text{COONa} \) (conjugate base)

Using the volume and concentration of sodium acetate:

\[
\text{Moles of } \text{CH}_3\text{COONa} = \text{Concentration} \times \text{Volume} = 1.127 \, \text{M} \times 0.112 \, \text{L}
\]
\[
\text{Moles of } \text{CH}_3\text{COONa} = 0.1262 \, \text{moles}
\]

### Step 3: Calculate concentrations of acid and base in the final buffer

**Total Volume of the Buffer:**

\[
\text{Total Volume} = 137.0 \, \text{mL} + 112.0 \, \text{mL} = 249.0 \, \text{mL} = 0.249 \, \text{L}
\]

**Concentration of \( \text{CH}_3\text{COOH} \) (weak acid):**

\[
[\text{CH}_3\text{COOH}] = \frac{0.1636 \, \text{moles}}{0.249 \, \text{L}} \approx 0.656 \, \text{M}
\]

**Concentration of \( \text{CH}_3\text{COONa} \) (conjugate base):**

\[
[\text{CH}_3\text{COONa}] = \frac{0.1262 \, \text{moles}}{0.249 \, \text{L}} \approx 0.507 \, \text{M}
\]

### Step 4: Apply the Henderson-Hasselbalch equation

\[
\text{pH} = 4.75 + \log\left(\frac{0.507}{0.656}\right)
\]

Calculating the log term:

\[
\frac{0.507}{0.656} \approx 0.772
\]

Now, calculate the logarithm:

\[
\log(0.772) \approx -0.113
\]

Thus, the pH calculation becomes:

\[
\text{pH} = 4.75 - 0.113 \approx 4.64
\]

Therefore, the theoretical pH of this buffer is:

**4.64**.

Pyridine is a weak base with the formula, C5H5N, the Kb for Pyridine is 1.8 x 10-9

In aqueous solution, pyridine partially dissociates according to the following reaction:

C5H5N + H2O ⇔ C5H5NH+ + OH-

Use the Kb equation to calculate the pH of the aqueous pyridine solution described below:

Volume:  400 mL

Concentration:  0.2851 M

Since this is a weak base, you can assume the amount of base dissociated is << 5% of the total amount of base present.

Report your answer with the correct number of significant digits.

To calculate the pH of a 0.2851 M solution of pyridine (\( \text{C}_5\text{H}_5\text{N} \)), we can set up the equilibrium expression based on its dissociation:

\[
\text{C}_5\text{H}_5\text{N} + \text{H}_2\text{O} \rightleftharpoons \text{C}_5\text{H}_5\text{NH}^+ + \text{OH}^-
\]

The base dissociation constant (\( K_b \)) for this reaction is given by:

\[
K_b = \frac{[\text{C}_5\text{H}_5\text{NH}^+][\text{OH}^-]}{[\text{C}_5\text{H}_5\text{N}]}
\]

### Step 1: Calculate the moles of \( \text{C}_5\text{H}_5\text{N} \)

\[
\text{Moles of } \text{C}_5\text{H}_5\text{N} = \text{Concentration} \times \text{Volume}
\]
\[
\text{Moles of } \text{C}_5\text{H}_5\text{N} = 0.2851 \, \text{M} \times 0.400 \, \text{L} = 0.11404 \, \text{moles}
\]

### Step 2: Set up the equilibrium expression

Let \( x \) be the concentration of \( \text{OH}^- \) produced at equilibrium. 

At equilibrium:
- \([\text{C}_5\text{H}_5\text{N}] = 0.2851 - x \approx 0.2851\) (since \( x \) is small)
- \([\text{C}_5\text{H}_5\text{NH}^+] = x\)
- \([\text{OH}^-] = x\)

The \( K_b \) expression is:

\[
K_b \approx \frac{x^2}{0.2851}
\]

### Step 3: Solve for \( x \)

Substituting \( K_b \):

\[
1.8 \times 10^{-9} = \frac{x^2}{0.2851}
\]

Multiplying both sides by 0.2851 gives:

\[
x^2 = 1.8 \times 10^{-9} \times 0.2851 = 5.13078 \times 10^{-10}
\]

Taking the square root:

\[
x = \sqrt{5.13078 \times 10^{-10}} \approx 2.26 \times 10^{-5} \, \text{M}
\]

### Step 4: Calculate the pOH and then the pH

\[
[\text{OH}^-] = 2.26 \times 10^{-5} \, \text{M}
\]

Calculate pOH:

\[
\text{pOH} = -\log(2.26 \times 10^{-5}) \approx 4.44 
\]

Now calculate pH using:

\[
\text{pH} + \text{pOH} = 14
\]

Thus:

\[
\text{pH} = 14 - \text{pOH} = 14 - 4.44 \approx 9.56
\]

### Final Result

The pH of the 0.2851 M pyridine solution is approximately:

**9.56** (rounded to four significant figures).