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Carbonic acid (H2CO3) produces a relatively acidic buffer when combined with it…

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Wize University Chemistry Textbook > Buffers and Titrations

Preparing Buffers with Desired pH

1 Activity

Carbonic acid (H2CO3) produces a relatively acidic buffer when combined with it's conjugate base HCO3- (pKa of carbonic acid = 3.6). If a solution of carbonic acid and bicarbonate has a pH of 4.2 what can we say about the relative concentrations of the two components?

More Preparing Buffers with Desired pH Questions:
Buffer Solutions
Consider a buffer containing 17 grams of sodium acetate (CH3COONa\rm CH_3COONa) and 21g of acetic acid (CH3COOH\rm CH_3COOH) Ka=1.8×10−5\rm K_a = 1.8 \times 10^{-5} in 1L of water.
a) What is the pH\rm pH of this buffer?
b) A graduate student finds a bottle of NaOH\rm NaOH(aq) but the concentration is unknown. To find the concentration the student adds 250mL of the solution to the buffer from part a) and observes a +0.4 change in the pH\rm pH. What is the concentration of the unknown NaOH\rm NaOH?
To a 100 mL of 0.1M of CH3COOH\rm CH_3COOH (acetic acid) is added 25mL of 0.17M NaOH\rm NaOH. The Ka\rm K_a of acetic acid is 1.8×10−51.8 \times 10^{-5}. Calculate the pH\rm pH of the solution. Please report your answer to two decimal places.
What are the [H3O+] and pH of a buffer containing 0.25 M propanoate and 0.15 M propanoic acid? (Ka of propanoic acid = 1.3 x 10-5)
pH and Henderson Hasselbalch
How many moles of NaCH3COO should be added to 700mL of a 0.15M of CH3COOH to get a pH of 4.8? Assume no volume change. pKa=4.75
moles of NaCH3COO?
V=700mL (CH3COOH)
Practice: Buffers/Titrations
Buffers and Titrations - Long Answer Practice Problem
A student prepares a buffer solution in the lab by titrating bleach (sodium hypochlorite, NaCℓO) with HCℓ until they produce a 1:1 buffer. Ka for HClO = 3.5 x 10-8.
a. What is the pH of their buffer solution?
Buffers: Citrate Buffer
A student prepares 350.00 mL of 0.186 M sodium citrate (Na3C6H5O7; pKb = 7.60) solution.
a) Calculate the volume (in mL) of 0.150 M HCl required to prepare a citrate buffer with pH = 7 using 25.00 mL of the sodium citrate solution.
b) Calculate the pH of a solution containing 25.00 mL of 0.223 M sodium hydrogen citrate (Na2HC6H5O7) and 25.00 mL of the sodium citrate solution.
Buffers: Preparing a Buffer at Specified pH
Use the following list of chemicals to answer the following questions.
0.50 M HF
0.25 M CH3COONa
Buffer Chemistry
Calculate the volume of 1.10 M NaOH required to prepare a buffer with pH = 9.79 from a 35.00 mL 0.100 M HCN solution (pKa (HCN) = 9.40).
Buffering Action
Determine the volume of 1.00 M BeO required to prepare a buffer with pH of 3.00 using a mixture of 150.0 mL containing 0.0800 M HF and 0.0500 M NaF.
Calculate the pH of a buffer containing 0.100 M HF and 0.250 M NaF after 1.25 mL of 1.30 M NaOH has been added to 75.00 mL of this buffer.
Buffer Chemistry
Calculate the pH of a buffer prepared by adding 1.50 mL of 0.900 M NaOH to 25.00 mL of 0.100 M HCN.
Buffer Chemistry
Calculate the pH of a buffer prepared by dissolving 0.350 g of KCN into 30.00 mL of 0.125 M HCN.
Buffer Chemistry
Calculate the pH of a buffer formed by combining 23.00 mL of 0.250 M HCN with 27.00 mL of 0.150 M KCN.
pKa HCN = 9.40
Buffering Action
Determine the volume of 0.750 M BeO required to prepare a buffer with pH of 5.00 using a mixture of 150.0 mL containing 0.0250 M CH3COOH and 0.0100 M CH3COO- (pKa = 4.75).
Buffering Action
Calculate the pH of a buffer containing 0.200 M HF and 0.150 M NaF after 1.20 mL of 1.25 M HCl has been added to 100.0 mL of this buffer.
Buffer Chemistry
Calculate the volume of 0.100 M NaOH required to prepare a buffer with pH = 9.91 from a 25.00 mL 7.50 x 10-2 M HCN solution (pKa (HCN) = 9.40).
Buffer Chemistry
Calculate the pH of a buffer prepared by adding 2.50 mL of 0.900 M NaOH to 25.00 mL of 0.320 M HCN.
Buffer Chemistry
Calculate the pH of a buffer prepared by dissolving 0.225 g of NaClO into 40.00 mL of 0.750 M HClO.
Buffer Chemistry
Calculate the pH of a buffer formed by combining 17.00 mL of 0.130 M HCN with 33.00 mL of 0.0850 M KCN.
Henderson Hasselbalch: Buffer solutions
A buffer is formed by dissolving 5 g5\ g of sodium benzoate, C6H5COONa\rm C_6H_5COONa, and 6 g6\ g of benzoic acid, C6H5COOH\rm C_6H_5COOH (Ka=6.6×10−5\rm K_a =6.6 \times 10^{-5}) in 1 L1\ L of water.
a) What is the pH\rm pH of this buffer?
b) The buffer is split into two beakers, 500 mL500\ mL each. To the first beaker is added 1 mL1\ mL of 1.0 M1.0\ M HCl\rm HCl and to the second is added 1 mL1\ mL of 1.0 M1.0\ M NaOH\rm NaOH. Calculate the pH\rm pH of each beaker.
Henderson Hasselbalch: Buffer solutions
90 mL90\ mL of 1.4 M1.4\ M HCl\rm HCl was added to 120 mL120\ mL or 1.5 M1.5\ M of bioxalate (HC2O42−\rm {HC_2O_4}^{2-}). The Ka\rm K_a of oxalic acid (H2C2O4\rm H_2C_2O_4), is 5.9×10−25.9 \times 10^{-2}. Calculate the pH\rm pH of the resulting solution.
Buffer solutions
To a solution of acetic acid (Ka=4.0×10−4\rm K_a = 4.0 \times 10^{-4}) was added 0.2 mol0.2\ mol of NaOH\rm NaOH. After the addition of NaOH\rm NaOH the volume was 1.4 L1.4\ L and the pH\rm pH was 4.1.
a) Calculate the number of moles of acetic acid that were originally in solution.
b) (True/False) This solution is a buffer, because pH=pKa±1\rm pH = pK_a \pm 1
An aqueous solution measured to be 0.5 M in sodium acetate and 0.3 M in acetic acid(Ka=1.8×10−5)\left(K_a=1.8\times10^{-5}\right) . Is this a buffer? and what is the pH?
pH and Acid/Base Equilibrium
A buffer is formed by dissolving 10g of sodium benzoate, C6H5COONa\rm C_6H_5COONa, and 12g of benzoic acid, C6H5COOH\rm C_6H_5COOH (Ka=6.6×10−5\rm K_a =6.6 \times 10^{-5}) in 1L of water.
a) What is the pH\rm pH of this buffer? (Two decimal Places)
b) The buffer is split into two beakers, 500mL each. To the first beaker is added 1mL of 1.0 M HCl\rm HCl and to the second is added 1mL of 1.0M NaOH\rm NaOH. Calculate the pH\rm pH of each beaker.
A buffer is prepared by adding 0.021 moles of HI to 250mL of 0.71M of (CH3)2NH(aq). What is the pH of this buffer? Kb((CH3)2NH(aq))=5.4x10-4 (Assume there is no change in volume)
Calculate the pH of a buffer solution that contains 0.25 M benzoic acid (C6H5COOH) and 0.15 M sodium benzoate (C6H5COONa). (Ka = 6.5 x 10-5 for benzoic acid)
A buffer is formed by dissolving 5g of sodium benzoate, C6H5COONa, and 6g of benzoic acid, C6H5COOH (Ka =6.5 x 10-5) in 1 L of water.
Buffer Solutions
What are the [H3O+] and pH of a buffer containing 0.25 M propanoate and 0.15 M propanoic acid? (Ka of propanoic acid = 1.3 x 10-5)
Henderson Hasselbalch Equation: pH of Buffer Solutions
Consider 500 mL of a sodium acetate/acetic acid buffer in which the concentration of sodium acetate is 0.50 M and that of acetic acid is 0.10 M. (Ka acetic acid= 1.8×10-5)
(a) What is the initial pH of the buffer?
(b) What is the pH after 100 mL of 0.10 M NaOH is added?
Buffer Solutions
Find the pH of a buffer that consists of 1.0 M sodium phenolate (C6H5ONa) and 1.2 M phenol (C6H5OH). (Ka of phenol = 1x10^-8)
Based on the given ionization constants, which of the following pairs of substances should be dissolved together in water to produce a buffer with pH = 9.85?
AcidKaBaseKbAcetic acid (CH3COOH)1.8×10−5Ammonia (NH3)1.8×10−5Bicarbonate (HCO3−)4.8×10−11Acetate (CH3COO−)5.6×10−10Carbonic acid (H2CO3)4.2×10−7Aniline (C6H5NH2)3.8×10−10Hydrazoic acid (HN3)1.9×10−5Hydrazine (N2H4)8.5×10−7Hydrofluoric acid (HF)6.8×10−4Methylamine (CH3NH2)4.4×10−4Phenol (C6H5OH)1.3×10−10Pyridine (C5H5N)1.5×10−9\def\arraystretch{2} \begin{array}{c|c|c|c} \hline \rm Acid &\rm K_a & \rm Base & \rm K_b\\ \hline \rm Acetic\ acid \ (CH_3COOH) &1.8 × 10^{−5} & \rm Ammonia \ (NH_3) & 1.8 × 10^{−5} \\ \hline \rm Bicarbonate \ (HCO_3^−) & 4.8 × 10^{−11} & \rm Acetate \ (CH_3COO^−) & 5.6 × 10^{−10} \\\hline \rm Carbonic \ acid \ (H_2CO_3) &4.2 × 10^{−7} &\rm Aniline \ (C_6H_5NH_2) & 3.8 × 10^{−10}\\ \hline \rm Hydrazoic \ acid \ (HN_3) &1.9 × 10^{−5} &\rm Hydrazine \ (N_2H_4) & 8.5 × 10^{−7}\\ \hline \rm Hydrofluoric \ acid \ (HF) & 6.8 × 10^{−4}& \rm Methylamine \ (CH_3NH_2) &4.4 × 10^{−4}\\ \hline \rm Phenol \ (C_6H_5OH) & 1.3 × 10^{−10}&\rm Pyridine \ (C_5H_5N) &1.5 × 10^{−9}\\ \hline \end{array}
Let us compare two buffer solutions that you could make from an aqueous solution of the weak acid, HA (aq). In both cases, you start from an identical initial solution. In both buffers, the target buffer ratio is 1A–: 2HA.