Acid Base and Titration pH & Buffer Solutions Worksheet

Direction: For exercises that require problem solving, you need to show the calculation set up leading to the answer in order to receive credit. Make sure your answers contain the correct units. You will need a Table of Ionization Constants to look up the Ka values of some acids.

A. Acid Base and Titration

    1. Does the pH of the solution increase, decrease, or stay the same when youAdd solid sodium oxalate, Na2C2O4, to 50.0 mL of 0.015 M oxalic acid, H2C2O4 ? Explain for credit. B. Add solid ammonium chloride to 75 mL of 0.20 M HF ? Explain for credit. C. Add 20.0 g of NaCl to 1.0 L of 0.10 M sodium acetate, CH3COONa ? Explain for credit. D. Add aqueous ammonium nitrate to 30.0 mL of 1.0 M NH3 ? Explain for credit. 2
    2. What is the pH of the solution that results from adding 30.0 mL of 0.015 M KOH to 50.0 mL of 0.015 M benzoic acid ? (Benzoic acid, Ka = 6.3 x 10–5 ) Answer: pH = 4.38
    3. What mass of sodium acetate, CH3COONa, must be added to 1.00 L of 0.10 M acetic acid (Ka = 1.8 x 10–5 ) to give a solution with a pH = 4.50 ? Assume the volume of the solution remains constant upon addition of sodium acetate. (Sodium acetate, molar mass = 82.0 g/mol) Answer: 4.7 g 3
    4. A buffer is composed of 1.360 g of KH2PO4 and 5.677 g of Na2HPO4. (Molar masses: KH2PO4 = 136.08 g/mol; Na2HPO4 = 141.96 g/mol) (H3PO4: Ka1 = 7.5 x 10–3 ; Ka2 = 6.2 x 10–8 ; Ka3 = 3.6 x 10–13) A. What is the pH of the buffer solution ? Answer: pH = 7.81 B. What additional mass (in g) of KH2PO4 must be added to decrease the pH of the buffer solution made in part A by 0.5 unit ? Answer: 2.9 g 4
    5. Which of the following combinations would be the best choice to buffer the pH of a solution at approximately 9 ? For each combination, explain why you would or would not choose it.
      •  HCl and NaCl
      • NH3 and NH4Cl
      • CH3COOH and CH3COON
    6. Which of the following combinations would be the best choice to buffer the pH of a solution at approximately 7 ? For each combination, explain why you would or would not choose it. A. H3PO4 and NaH2PO4 B. NaH2PO4 and Na2HPO4 C. Na2HPO4 and Na3PO4 5
    7. Assume you dissolve 0.235 g of the weak acid benzoic acid, C6H5COOH, in enough water to make 1.00 x 102 mL of solution and then titrate the solution with 0.108 M NaOH. [Benzoic acid: Ka = 6.3 x 10–5 ; MM = 122.1 g/mol)] C6H5COOH(aq) + OH– (aq)  C6H5COO– (aq) + H2O() A. What is the pH of the benzoic acid solution prior to titration ? Answer: pH = 2.96 B. What is the pH and the concentrations of Na+ and C6H5COO– ions at the equivalence point ? Answer: pH = 8.21; [Na+ ] = [C6H5COO– ] = 0.0163 M 6
    8. 25.0 mL of 0.100 M benzoic acid (Ka = 6.3 x 10–5 ) is being titrated with 0.150 M NaOH. What color indicator would be most suitable for this titration ? Why ? Indicator pKa Indicator pKa Thymol Blue 2.0 Bromothymol Blue 6.8 Methyl Orange 3.2 Phenol Red 7.8 Bromocresol green 4.6 Phenolphthalein 8.8 Methyl Red 5.4 Alizarin yellow 11.0
    9. Describe how to prepare a 1.00 L buffer solution from stock solutions of 2.00 M NH3 and 1.50 M NH4Cl to have a pH of 9.50. Calculate the volumes, in mL, of NH3 and NH4Cl that you would mix together to create 1 liter of buffer. The Ka of NH4 + is 5.6 x 10–10 . [Hint: One approach is to use the Henderson-Hasselbach equation to calculate the ratio of NH3/NH4 + ion; another equation needed is the sum of the volumes of NH3 and NH4Cl is equal to 1.00 L. Combining these two equations should lead to the answers.] Answers: VNH3 = 0.57 L; VNH4Cl = 0.43 L 7
    10. For each of the following cases, write a balanced equation for the reaction and determine whether the pH of the mixture is less than 7, equal to 7, or greater than 7. Explain for credit. A. Equal volumes of 0.10 M acetic acid, CH3COOH, and 0.10 M KOH are mixed. B. 25 mL of 0.015 M NH3 is mixed with 25 mL of 0.015 M HCl. C. 150 mL of 0.20 M HNO3 is mixed with 75 mL of 0.40 M NaOH. D. 25 mL of 0.45 M H2C2O4 is mixed with 25 mL of 0.90 M NaOH. 8 
    11. A buffer solution with a pH of 12.00 consists of Na3PO4 (MM = 163.9 g/mol) and Na2HPO4 (MM = 142.0 g/mol). The volume of the solution is 200.0 mL. [H3PO4: Ka1 = 7.5 x 10−3 ; Ka2 = 6.2 x 10−8 ; Ka3 = 3.6 x 10−13] A. Which component of the buffer is present in a larger amount (i.e. HPO4 2− or PO4 3− ) ? B. If the concentration of Na3PO4 is 0.400 M, what mass of Na2HPO4 is present ? Answer: 32 g Na2HPO4 C. Which component of the buffer (i.e. PO4 −3 or HPO4 −2 ) must be added to raise the pH of the buffer from 12.00 to 12.25 ? What additional mass, in g, of that component is required ? [Note: Mass must be either in Na3PO4 or Na2HPO4] Answer: about 10 g (state the component) 9 B. Solubility and Ksp
    12. Name two insoluble salts of each of the following ions. [See tables of solubility guidelines on last page.] a) SO4 2– : b) Ni2+ : c) Br– : d) Zn2+
    13. Which compound in each pair is more soluble in water than is predicted by a calculation from Ksp ? Explain your choice. a) AgI or Ag2CO3 b) PbS or PbCl2 c) AgCl or AgCN
    14. You have 95 mL of a solution that has a lead(II) concentration of 0.0012 M. Will PbCl2 precipitate when 1.20 g of solid NaCl is added ? [NaCl, MM = 58.44 g/mol; Ksp of PbCl2 = 1.17 x 10−5 ] 10
    15. You place 1.234 g of solid Ca(OH)2 in 1.00 L of pure water at 25C. The pH of the solution is found to be 12.68. Estimate the value of Ksp for Ca(OH)2. Answer: Ksp = 5.5 x 10–5 16. Determine the solubility, in mg/mL, of barium fluoride, BaF2 (MM = 175 g/mol). The Ksp of BaF2 is 1.8 x 10–7 . A. in pure water Answer: 0.62 mg/mL B. in water containing 5.0 mg/mL KF (molar mass = 58.1 g/mol) Answer: 4.3 x 10–3 mg/mL 11
    16. Calculate the equilibrium constant for the following reaction: Zn(OH)2(s) + 2 CN− (aq)  Zn(CN)2(s) + 2 OH− (aq) The involved equilibria are: Zn(OH)2(s) Ksp = 3.0 x 10−17 Zn(CN)2(s) Ksp = 8.0 x 10−12 Does the equilibrium lie predominantly to the left or to the right ? Can zinc hydroxide be transformed into zinc cyanide by adding a soluble salt of the cyanide ion ?
    17. If 55 mg of lead(II) sulfate (PbSO4, MM = 303.3 g/mol) is placed in 250 mL of pure water, does all of it dissolve ? If not, how much dissolves ? (PbSO4, Ksp = 2.5 x 10–8 ) Answer: 12 mg dissolves 12
    18. The cations of Ba2+ and Sr2+ can be precipitated as very insoluble sulfates: BaSO4 Ksp = 1.1 x 10−10 SrSO4 Ksp = 3.4 x 10−7 A. If you add Na2SO4 to a solution containing these metal cations, each with a concentration of 0.10 M, which is precipitated first, BaSO4 or SrSO4 ? Why ? B. What will be the concentration of the first ion that precipitates (Ba2+ or Sr2+) when the second, more soluble salt begins to precipitate ? Answer: 3.2 x 10−5 M
    19. The Ca2+ ion in hard water can be precipitated as CaCO3 (Ksp = 3.4 x 10–9 ) by adding soda ash, Na2CO3. If the Ca2+ ion concentration in hard water is 0.010 M and if the Na2CO3 is added until the carbonate ion concentration is 0.050 M, what percentage of the Ca2+ ions have been removed from the water ? (You may neglect the hydrolysis of water by carbonate ion.) Answer: Essentially all (> 99.99%) of the Ca2+ has been removed.

 GUIDELINES TO SOLUBILITY OF IONIC COMPOUNDS IN WATER

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