Blog Post #3- Investigating a Question About Buffer Solutions

Recently, a question was posed to me by a classmate:
a) Begin by defining what a buffer is and summarize 3-4 buffer characteristics
b) Solve this problem: A buffer is created by combining 150.0mL of 0.25M HCHO2 with 75.0mL of 0.20M NaOH. Determine the pH of the buffer.

While this question is undoubtedly one I need more practice on, the difficulty level does not seem up to par with the assignment. This same type of problem was given to us in the homework, and on our recent test. Personally, I feel that it is significant due to the fact I struggled with a similar type of question on the exam. However, I doubt that the instructor feels the same way.

Here is my answer to this question: 

a) A buffer is a solution which resists a change in pH by neutralizing an added acid or base. Some buffer characteristics are: 

1. A buffer contains significant amounts of both a weak acid and its conjugate base. For example, human blood is composed of the weak acid carbonic acid (H2CO3) and its conjugate base (HCO3-). 
2. When additional base is added to a buffer solution, the weak acid reacts with the base, neutralizing it.
3. When additional acid is added to a buffer solution, the conjugate base reacts with the acid, neutralizing it. 
4. Weak acids and conjugate bases by themselves do not contain sufficient bases or acids respectively to work as neutralizing/buffering agents. 

b) In order to solve this problem, I must first determine the amount of the weak acid that is converted to a conjugate base by the addition of the NaOH. This is found by comparing stoichiometric ratios presented in the equation:

HCHO₂+NaOH→ H₂O+NaCHO₂

_{From this equation, we can then set up a table to track the changes induced by the addition of the NaOH. Please note that the role of water in this reaction does not effect the outcome of our problem and thus has been omitted:}

	HCHO2	NaOH	NaCHO2
Before Addition	0.0375mol	0mol	0mol
Addition	0mol	0.015mol	0mol
After Addition	0.0225mol	≈0mol	0.015mol

Once we have determined the amount of HCHO2 present in equilibrium, we can use the concentrations of the weak acid and the conjucate base, along with the pK_a value which will be plugged into the Henderson-Hasselbalch equation in order to find the pH of the buffer solution.The Henderson-Hasselbalch equation can be used because the "x is small approximation" is relevant in this context due to the fact that the initial concentration of weak acid is much larger than its Ka value.

 The Henderson-Hasselbalch equation is as follows:

pH=pK_a+log[base] /[acid]

The pKa was determined by using the Ka value, 1.8x10^-4 in this equation: 

pKa=-logKa

This value (3.7447), along with the amounts of the weak acid (0.0225mol) and its conjugate base(0.015mol), were plugged in to the Henderson-Hasselbach equation, leaving us with this expression: 

pH=3.7447+log(0.015/0.0225)

Solved algebraically, we are left with: 

pH=3.57

My reference for this problem was:
Tro, Nivaldo J. Chemistry: A Molecular Approach. 2011. Pages 667,714-715, 722-723