Buffers and Titrations (Worksheet)

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Name: ______________________________

Section: _____________________________

Student ID#:__________________________

Work in groups on these problems. You should try to answer the questions without referring to your textbook. If you get stuck, try asking another group for help.

Q1

Consider the following weak acids and their K_a values:

Acetic acid	K_a = 1.8 x 10^–5
Phosphoric acid	K_a1 = 7.5 x 10^–3
Hypochlorous acid	K_a = 3.5 x 10^–8

You are to prepare buffers at pH = 2.8, 4.5, and 7.5. In the 2^nd column write the fomulas for the respective weak acid–conjugate base buffer system that is the best choice for each pH from the listed acids? Explain your reasoning in the 3^rd column.

pH	weak acid–conjugate base combination	Rational
2.8
4.5
7.5

Q2

Reconsider the 100.0 mL solution containing 0.010 mol acetic acid, HC₂H₃O₂, and 0.010 mol sodium acetate, NaC₂H₃O₂, which was first introduced in the reading. Determine the resulting pH if 0.005 mol NaOH is added to the buffer.

Q3

Consider the titration of an acetic acid solution with a sodium hydroxide solution at the following three stages of the titration: (i) before the titration begins, (ii) when the number of moles of sodium hydroxide added is equal to 1/2 the number of moles of acetic acid originally in the beaker, and (iii) at the endpoint. For each of the following questions, select one of the above three stages and explain your reasoning.

When does the reaction solution contain mostly acetate ion?
When does it contain mostly acetic acid?
When does it contain significant amounts of both?
At what point during the titration is the reaction solution's pH at its lowest value? e) At what point is it at its highest value? f ) When is it between the two extreme values?

Q4

Student A claims that she can calculate the pH of a buffer system without knowing the actual concentrations of the acid and conjugate base. Student B disagrees, citing the fact that the buffer equation clearly requires concentrations. Who is correct? Explain.

Q5

The carbonate buffer system is very important in regulating blood pH levels. Carbonic acid is diprotic and therefore has two Ka values:

\[H_2CO_{3} (aq) \rightarrow H^+(aq) + HCO_3^– (aq) \nonumber \]

with \(K_{a1} = 4.2 \times 10^{–7}\)

\[HCO_3^{-}(aq) \rightarrow H^+ (aq) + CO_3^{2–} (aq) \nonumber \]

with \(K_{a2} = 4.8 \times 10^{–11}\)

Since the second dissociation has a \(K_a\) value significantly smaller than that of the first dissociation, it can be assumed to have no effect on the H₂CO₃(aq)/HCO₃^–(aq) equilibrium.

The pH of blood is 7.4. What is the ratio of carbonic acid to bicarbonate ion in blood?

Q6

Biochemical experiments frequently utilize a buffer system based on tris-(hydroxymethyl)aminomethane, (HOCH₂)₃CNH₂, which is also called TRIS or THAM. The pKa of the conjugate acid of TRIS, (HOCH₂)₃CNH₃⁺, is 8.075. What mole ratio of acid-to-base is required to prepare a buffer at the same pH as human blood, pH = 7.4?

Q7

When a strong base is gradually added dropwise to a weak acid, the pH changes at each addition. When the appropriate quantity of base has been added to react with all of the acid, the pH changes sharply, indicating the endpoint of the titration. A plot of pH versus volume of base added gives what is known as a titration curve.

Consider the titration of 25.00 mL of 0.1000 M acetic acid with 0.1000 M NaOH.

(a) Write the equation for the titration reaction.

(b) Determine the volume of NaOH solution required to reach the endpoint.

(b) The chart below has entries for several steps along the titration curve. To calculate the pH at each step, you first must understand the chemistry at that step, and then you can decide the appropriate method to calculate the pH. For each volume listed, (i) list the major species in solution, (ii) determine whether the Ka equation, Kb equation, buffer equation, or solution equilibrium equation is appropriate for the calculation of the solution [H₃O⁺] and pH, and (iii) complete the calculations.

Volume NaOH Added (mL) Major Species Appropriate Equation
[H₃O⁺]
pH
0.00
5.00
12.50
20.00
25.00
30.00

Q8

It is convenient to prepare a buffer by adding sodium hydroxide solution to a weak acid solution, using a pH meter to monitor the pH until the desired buffer pH is reached. This technique accounts for non-ideal conditions that can result in small but significant errors in theoretically-based calculations. (i.e. The Henderson-Hasselbach equation is not perfect and this practical approach overcomes any discrepancy.) a) Explain why this method works. b) How many drops of 1.0 M NaOH should be added to 200.0 mL of 0.050 M acetic acid, Ka = 1.8 x 10^–⁵, to make a buffer with pH = 5.0? Assume that each drop delivers an average volume of 0.05 mL.

Q9

A scientist from a team of chemists and oncologists extracted and purified a natural product from a plant of the datura sp., which was collected in the Amazon Basin. An initial test indicated that the compound was likely a weak monoprotic carboxylic acid. The scientist weighed 252.8 mg of the compound. She used it to prepare 250.00 mL of solution, and then titrated it with 0.01000 M NaOH. Her titration curve is shown below.

How many moles of the compound were in the sample?
What is the molar mass of the compound?
What is the pKa of the compound?