For each solution, calculate the initial and final pH after adding 0.010 mol of HCl: a. 500.0 mL of pure water b. 500.0 mL of a buffer solution that is 0.125 M in HC2H3O2 and 0.115 M in NaC2H3O2 c. 500.0 mL of a buffer solution that is 0.155 M in C2H5NH2 and 0.145 M in C2H5NH3Cl.

Verified step by step guidance

1. <insert step 1> Identify the type of solution for each part: pure water, acetic acid/acetate buffer, and ethylamine/ethylammonium chloride buffer.
2. <insert step 2> For part (a), calculate the initial pH of pure water, which is neutral at pH 7. Then, determine the change in pH after adding 0.010 mol of HCl by calculating the concentration of HCl in the solution and using the formula for pH: \( \text{pH} = -\log[H^+] \).
3. <insert step 3> For part (b), use the Henderson-Hasselbalch equation to calculate the initial pH of the acetic acid/acetate buffer: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \), where \( \text{pK}_a \) is the acid dissociation constant for acetic acid. Then, calculate the final pH after adding HCl by adjusting the concentrations of acetic acid and acetate ion and applying the Henderson-Hasselbalch equation again.
4. <insert step 4> For part (c), calculate the initial pH of the ethylamine/ethylammonium chloride buffer using the Henderson-Hasselbalch equation: \( \text{pH} = \text{pK}_b + \log \left( \frac{[\text{B}]}{[\text{BH}^+]} \right) \), where \( \text{pK}_b \) is the base dissociation constant for ethylamine. After adding HCl, adjust the concentrations of ethylamine and ethylammonium ion and use the Henderson-Hasselbalch equation to find the final pH.
5. <insert step 5> Summarize the effect of adding a strong acid (HCl) to each solution, highlighting the buffering capacity of the buffer solutions compared to pure water.>