# 1. Suppose that you have 0.500 L of each of the following solutions, and an unlimited supply of water.

(Note: C9H7NHBr is a salt containing the ions C9H7NH+ and Br− and C9H7N is quinoline, an organic
base with pKb = 6.24 at 298 K. If you like, you may represent C9H7NH+ as HB+ and C9H7N as B.)
0.104 mol L−1 C9H7NHBr (aq) 0.122 mol L−1 HBr(aq) 0.113 mol L−1 NaOH(aq)
(a) Provide simple instructions for preparing 1.00 L of a buffer solution having pH = 7.00 at
298 K. Focus on preparing the buffer that contains the highest possible concentrations of the active
components. Your instructions should include the volumes of the solutions required.
(b) What is the buffer capacity of the resulting solution? How does this compare with the
buffer capacity of pure water? (The buffer capacity is the number of moles of NaOH that must be
added to 1.0 L of solution to raise the pH by one unit.)

for this particular question do you start off by finding out the limiting reagent? But i am stuck on the part where you have to write the formula for this particular question, please any help would be appreciated

## This is a buffer problem and you must use the Henderson-Hasselbalch equation,

The C9H7NHBr is the base.
HBr is the acid.

## To prepare a buffer solution with a pH of 7.00 using the given solutions, you need to combine the weak base C9H7NH and its conjugate acid C9H7NHBr in such a way that the resulting solution can resist changes in pH upon addition of small amounts of acid or base.

(a) Here are the steps to prepare the buffer solution:

1. Calculate the number of moles of the weak base required:
Since the solution of C9H7NHBr has a concentration of 0.104 mol L−1, you want to prepare 1.00 L of the buffer solution. Therefore, you need 0.104 moles of C9H7NH.

2. Calculate the number of moles of the conjugate acid required:
The concentration of HBr is 0.122 mol L−1. Since the pH of the buffer needs to be 7.00, which is slightly acidic, you need to add an excess of the conjugate acid. In this case, you can use twice the number of moles of the weak base. So you will need 2 * 0.104 = 0.208 moles of HBr.

3. Calculate the volume of the solutions needed to achieve the desired moles:
Using the molarity and moles, you can calculate the volumes of the solutions required.
Volume of C9H7NH = (0.104 mol) / (0.104 mol L−1) = 1.00 L
Volume of HBr = (0.208 mol) / (0.122 mol L−1) = 1.70 L

4. Combine the solutions and add water:
In a 1.00 L volumetric flask, add 1.00 L of the C9H7NH solution and 1.70 L of the HBr solution. Then add enough water to achieve a total volume of 1.00 L.

(b) The buffer capacity is a measure of how well the buffer can resist changes in pH upon addition of acid or base. It can be calculated by the change in moles of the added acid or base needed to change the pH by one unit.

The buffer capacity of the resulting solution can be determined by taking the derivative of the Henderson-Hasselbalch equation (∆[B]/∆pH) at pH = 7.00, where ∆[B] is the change in concentration of the weak base. Since the pH is constant at 7.00, the buffer capacity is infinite.

In comparison, the buffer capacity of pure water is very low because it does not contain any buffering agents. Addition of even a small amount of acid or base can cause a significant change in its pH.