Question

1. Chemical Calculations 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g) MM = 17.0 g/mol MM = 32.0 g/mol MM = 30.0 g/mol MM = 18.0 g/mol a) How many moles of H atoms are present in 45.1 g of NH3? b) How many grams of NO correspond to 8.25 x 10^24 molecules of NO? c) How many moles of O2 correspond to 25.0 g O2? 2. Balance the following chemical equation. Classify the reaction as combination, decomposition, displacement, exchange, or combustion. FeCl3(aq) + Na2CO3(aq) → Fe2(CO3)3(aq) + NaCl(aq) 3. Assign oxidation numbers to every element in the redox reaction below. What gets oxidized and what gets reduced in the following chemical equation? Cl2(g) + 2 NaBr(aq) → 2 NaCl(aq) + Br2(g) Oxidized: Cl2(g) Reduced: NaBr(aq) 4. Using the kinetic molecular theory, explain why the solubility of a gas increases as pressure increases. 5. Write a balanced chemical equation (using equilibrium arrows when necessary) for the ionization of each of the following substances when they dissolve in water. Mg(NO3)2(s), magnesium nitrate – a strong electrolyte: 2 Mg(NO3)2(s) → 2 Mg^2+(aq) + 4 NO3^-(aq) HF(l), hydrofluoric acid – a weak electrolyte: HF(l) ⇌ H+(aq) + F^-(aq) 6. Using the solubility rules in the lecture notes or textbook, predict whether the following are soluble or insoluble. a) Ca3(PO4)2 - Insoluble b) BaSO4 - Insoluble c) PbCl2 - Insoluble d) Na2S - Soluble 7. Predict the products and write the balanced chemical equation for the precipitation reaction of ammonium sulfate with lead(II) nitrate. Use solubility rules to determine the phase labels. (NH4)2SO4(aq) + Pb(NO3)2(aq) → 2 NH4NO3(aq) + PbSO4(s) 8. Solution Concentrations a) How many grams of glucose must be added to 275 g of water in order to prepare a 25.0% (m/m) glucose solution. b) Calculate the volume (in mL) of a 2.5 M AgNO3 solution needed to obtain 0.065 moles of AgNO3. c) You have 355 mL of a 1.75 M glucose solution. What is the diluted concentration after adding 875 mL of water to the glucose solution? 9. Equilibrium & Le Châtelier’s Principle. 2 SO2(g) + 2 H2O(g) → 2 H2S(g) + 3 O2(g) ΔH = –125 kJ a) At equilibrium, what are the two conditions that must be met? b) Write the equilibrium constant expression for the equilibrium reaction shown above. c) Calculate Keq at 125°C for the equilibrium reaction above if a reaction mixture at equilibrium contains [SO2(g)] = 0.025 M, [H2O(g)] = 0.025 M, [H2S(g)] = 2.5 M, and [O2(g)] = 2.0 M Using the equilibrium reaction above, predict the direction that equilibrium shifts under the following conditions. d) Increasing the concentration of H2O(g): e) Decreasing the concentration of SO2(g): f) Increasing pressure: g) Decreasing temperature: 10. Acid-Base Ionization Reactions (Don’t forget to show charges on any acid or base ions) a) Write a chemical equation showing HPO42– acting as a Brønsted-Lowry acid in H2O. HPO42–(aq) + H2O(l) → H2PO4–(aq) + OH–(aq) b) Write a chemical equation showing HPO42– acting as a Brønsted-Lowry base in H2O. HPO42–(aq) + H2O(l) ⇌ PO43–(aq) + H3O+(aq) c) Write a chemical equation showing NH4+ acting as a Brønsted-Lowry acid in H2O. NH4+(aq) + H2O(l) ⇌ H3O+(aq) + NH3(aq) d) Write a chemical equation showing CH3COOH acting as a Brønsted-Lowry acid in H2O. CH3COOH(aq) + H2O(l) ⇌ H3O+(aq) + CH3COO–(aq) 11. What is the conjugate acid of HCO3–? H2CO3 What is the conjugate base of CH3NH3+? CH3NH2 12. Conjugate Acids & Bases and Relative Strengths of Acids a) Identify the Brønsted acid & its conjugate base and the Brønsted base & its conjugate acid below. H2PO4–(aq) + HS–(aq) → H3PO4(aq) + S2–(aq) Brønsted acid: H2PO4– Conjugate base: HPO42– Brønsted base: HS– Conjugate acid: H2S b) Using the table on p.7 of the Chapter 10 Lecture Notes, what is the stronger acid? CIRCLE: H3PO4 OR HS– c) If the Ka value for HNO2 is 4.5 x 10–4 and the Ka value for HCN is 4.9 x 10–10, which one is the stronger acid? CIRCLE: HNO2 OR HCN 13. Acid-Base Chemistry Calculations a) If the solution pH is 5.58, calculate [OH–]. Is this solution acidic or basic? b) What is the pH of a solution if [OH–] = 1.9 x 10–3? Is this solution acidic or basic? c) Calculate the pH of this phosphate buffer if it is composed of 0.125 M H2PO4– and 0.675 M HPO42–. For H2PO4–, Ka = 6.31 x 10–8

          1. Chemical Calculations
4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
MM = 17.0 g/mol
MM = 32.0 g/mol
MM = 30.0 g/mol
MM = 18.0 g/mol

a) How many moles of H atoms are present in 45.1 g of NH3?
b) How many grams of NO correspond to 8.25 x 10^24 molecules of NO?
c) How many moles of O2 correspond to 25.0 g O2?

2. Balance the following chemical equation. Classify the reaction as combination, decomposition, displacement, exchange, or combustion.
FeCl3(aq) + Na2CO3(aq) → Fe2(CO3)3(aq) + NaCl(aq)

3. Assign oxidation numbers to every element in the redox reaction below. What gets oxidized and what gets reduced in the following chemical equation?
Cl2(g) + 2 NaBr(aq) → 2 NaCl(aq) + Br2(g)
Oxidized: Cl2(g)
Reduced: NaBr(aq)

4. Using the kinetic molecular theory, explain why the solubility of a gas increases as pressure increases.

5. Write a balanced chemical equation (using equilibrium arrows when necessary) for the ionization of each of the following substances when they dissolve in water.
Mg(NO3)2(s), magnesium nitrate – a strong electrolyte:
2 Mg(NO3)2(s) → 2 Mg^2+(aq) + 4 NO3^-(aq)

HF(l), hydrofluoric acid – a weak electrolyte:
HF(l) ⇌ H+(aq) + F^-(aq)

6. Using the solubility rules in the lecture notes or textbook, predict whether the following are soluble or insoluble.
a) Ca3(PO4)2 - Insoluble
b) BaSO4 - Insoluble
c) PbCl2 - Insoluble
d) Na2S - Soluble

7. Predict the products and write the balanced chemical equation for the precipitation reaction of ammonium sulfate with lead(II) nitrate. Use solubility rules to determine the phase labels.
(NH4)2SO4(aq) + Pb(NO3)2(aq) → 2 NH4NO3(aq) + PbSO4(s)

8. Solution Concentrations
a) How many grams of glucose must be added to 275 g of water in order to prepare a 25.0% (m/m) glucose solution.
b) Calculate the volume (in mL) of a 2.5 M AgNO3 solution needed to obtain 0.065 moles of AgNO3.
c) You have 355 mL of a 1.75 M glucose solution. What is the diluted concentration after adding 875 mL of water to the glucose solution?

9. Equilibrium & Le Châtelier’s Principle.
2 SO2(g) + 2 H2O(g) → 2 H2S(g) + 3 O2(g) ΔH = –125 kJ
a) At equilibrium, what are the two conditions that must be met?
b) Write the equilibrium constant expression for the equilibrium reaction shown above.
c) Calculate Keq at 125°C for the equilibrium reaction above if a reaction mixture at equilibrium contains [SO2(g)] = 0.025 M, [H2O(g)] = 0.025 M, [H2S(g)] = 2.5 M, and [O2(g)] = 2.0 M
Using the equilibrium reaction above, predict the direction that equilibrium shifts under the following conditions.
d) Increasing the concentration of H2O(g):
e) Decreasing the concentration of SO2(g):
f) Increasing pressure:
g) Decreasing temperature:

10. Acid-Base Ionization Reactions (Don’t forget to show charges on any acid or base ions)
a) Write a chemical equation showing HPO42– acting as a Brønsted-Lowry acid in H2O.
HPO42–(aq) + H2O(l) → H2PO4–(aq) + OH–(aq)

b) Write a chemical equation showing HPO42– acting as a Brønsted-Lowry base in H2O.
HPO42–(aq) + H2O(l) ⇌ PO43–(aq) + H3O+(aq)

c) Write a chemical equation showing NH4+ acting as a Brønsted-Lowry acid in H2O.
NH4+(aq) + H2O(l) ⇌ H3O+(aq) + NH3(aq)

d) Write a chemical equation showing CH3COOH acting as a Brønsted-Lowry acid in H2O.
CH3COOH(aq) + H2O(l) ⇌ H3O+(aq) + CH3COO–(aq)

11. What is the conjugate acid of HCO3–? H2CO3
What is the conjugate base of CH3NH3+? CH3NH2

12. Conjugate Acids & Bases and Relative Strengths of Acids
a) Identify the Brønsted acid & its conjugate base and the Brønsted base & its conjugate acid below.
H2PO4–(aq) + HS–(aq) → H3PO4(aq) + S2–(aq)
Brønsted acid: H2PO4–
Conjugate base: HPO42–
Brønsted base: HS–
Conjugate acid: H2S

b) Using the table on p.7 of the Chapter 10 Lecture Notes, what is the stronger acid?
CIRCLE: H3PO4 OR HS–
c) If the Ka value for HNO2 is 4.5 x 10–4 and the Ka value for HCN is 4.9 x 10–10, which one is the stronger acid?
CIRCLE: HNO2 OR HCN

13. Acid-Base Chemistry Calculations
a) If the solution pH is 5.58, calculate [OH–]. Is this solution acidic or basic?
b) What is the pH of a solution if [OH–] = 1.9 x 10–3? Is this solution acidic or basic?
c) Calculate the pH of this phosphate buffer if it is composed of 0.125 M H2PO4– and 0.675 M HPO42–. For H2PO4–, Ka = 6.31 x 10–8

Added by Kimberly M.

1. Which of the following aqueous solutions are good buffer systems? A. 0.14 M acetic acid + 0.13 M nitric acid. 0.13 M calcium hydroxide + 0.22 M calcium chloride. 0.39 M potassium nitrate + 0.23 M calcium nitrate. 0.24 M hydrochloric acid + 0.16 M potassium chloride. 0.27 M ammonia + 0.39 M ammonium nitrate. B.0.34 M calcium bromide + 0.29 M sodium bromide 0.24 M nitric acid + 0.16 M potassium nitrate 0.15 M potassium hydroxide + 0.24 M potassium bromide 0.10 M hydrocyanic acid + 0.20 M sodium cyanide 0.27 M ammonia + 0.40 M sodium hydroxide C. 0.10 M nitrous acid + 0.11 M sodium nitrite 0.27 M ammonium bromide + 0.36 M ammonia 0.25 M hydroiodic acid + 0.24 M sodium iodide 0.19 M sodium hydroxide + 0.29 M sodium chloride 0.40 M hypochlorous acid + 0.22 M potassium hypochlorite 2a. A solution contains 0.155 M sodium acetate and 0.496 M acetic acid. The pH of this solution is b. A solution contains 0.482 M sodium cyanide and 0.406 M hydrocyanic acid. The pH of this solution is c. A solution contains 0.400 M sodium nitrite and 0.225 M nitrous acid. The pH of this solution is 3a. A solution contains 0.496 M NH4I and 6.61×10-2 M ammonia. The pH of this solution is b. A solution contains 0.406 M NH4I and 0.187 M ammonia. The pH of this solution is c. The compound ethylamine is a weak base like ammonia. A solution contains 0.189 M C2H5NH3+ and 0.135 M ethylamine, C2H5NH2. The pH of this solution is 4a. A buffer solution is 0.498 M in CH3COOH and 0.247 M in CH3COONa. If Ka for CH3COOH is 1.8×10-5, what is the pH of this buffer solution? b. A buffer solution is 0.458 M in H2S and 0.392 M in KHS. If Ka1 for H2S is 1.0*10-7 , what is the pH of this buffer solution? c. A buffer solution is 0.312 M in NaHSO3 and 0.305 M in Na2SO3. If Ka for HSO3- is 6.4*10-8, what is the pH of this buffer solution? 5a. An aqueous solution contains 0.24 M ammonia. One liter of this solution could be converted into a buffer by the addition of: (Assume that the volume remains constant as each substance is added.) 0.25 mol HClO4 0.24 mol BaCl2 0.12 mol HClO4 0.12 mol KOH 0.25 mol NH4Cl b. An aqueous solution contains 0.26 M hydrofluoric acid. One Liter of this solution could be converted into a buffer by the addition of: (Assume that the volume remains constant as each substance is added.) 0.27 mol KF 0.131 mol NaOH 0.13 mol HBr 0.26 mol KCl 0.27 mol HBr c. An aqueous solution contains 0.30 M ammonium bromide. One liter of this solution could be converted into a buffer by the addition of: (Assume that the volume remains constant as each substance is added.) 0.30 mol HNO3 0.14 mol NaOH 0.29 mol NH3 0.14 mol HNO3 0.29 mol KBr 6a. A buffer solution made from CH3COOH and CH3COONa has a pH of 4.43. If pKa for CH3COOH is 4.74, what is the [CH3COO-]/[CH3COOH] in the buffer? b. The pKa value for H2CO3 is 6.38. What mole ratio of KHCO3 to H2CO3 is needed to prepare a buffer with a pH of 6.18? c. The pKa value for HCO3- is 10.30. What mole ratio of Na2CO3 to NaHCO3 is needed to prepare a buffer with a pH of 9.97. 7a. How many grams of solid sodium acetate should be added to 1.00 L of a 0.298 M acetic acid solution to prepare a buffer with a pH of 5.362 ? b. How many grams of solid sodium cyanide should be added to 1.50 L of a 0.248 M hydrocyanic acid solution to prepare a buffer with a pH of 8.418? c. How many grams of solid ammonium chloride should be added to 0.500L of a 0.147 M ammonia solution to prepare a buffer with a pH of 8.370 ? 8A. A 1 liter solution contains 0.538 M hydrocyanic acid and 0.404 M sodium cyanide. Addition of 0.202 moles of hydrochloric acid will: (Assume that the volume does not change upon the addition of hydrochloric acid.) Raise the pH slightly Lower the pH slightly Raise the pH by several units Lower the pH by several units Not change the pH Exceed the buffer capacity B. A 1 liter solution contains 0.313 M acetic acid and 0.417 M potassium acetate. Addition of 0.078 moles of barium hydroxide will: (Assume that the volume does not change upon the addition of barium hydroxide.) Raise the pH slightly Lower the pH slightly Raise the pH by several units Lower the pH by several units Not change the pH Exceed the buffer capacity C. A 1 liter solution contains 0.357 M hypochlorous acid and 0.268 M potassium hypochlorite. Addition of 0.295 moles of hydrobromic acid will: (Assume that the volume does not change upon the addition of hydrobromic acid.) Raise the pH slightly Lower the pH slightly Raise the pH by several units Lower the pH by several units Not change the pH Exceed the buffer capacity 9A. A buffer solution is made that is 0.368 M in HF and 0.368 M in NaF. If Ka for HF is 7.20*10-4 , what is the pH of the buffer solution? pH = ______ Write the net ionic equation for the reaction that occurs when 0.105 mol HNO3 is added to 1.00 L of the buffer solution. (Use the lowest possible coeffieicents, Omit state of matter. Use H3O+ instead of H+) B. A buffer solution is made that is 0.420 M in H2CO3 and 0.420 M in KHCO3. If Ka1 for H2CO3 is 4.20*10-7, what is the pH of the buffer solution? pH = ____ Write the net ionic equation for the reaction that occurs when 0.085 mol HCl is added to 1.00 L of the buffer solution. (Use the lowest possible coefficients. Omit states of matter. Use H3O+ instead of H+) C. A buffer solution is made that is 0.378 M in HNO2 and 0.378 M in NaNO2. If Ka for HNO2 is 4.50*10-4, what is the pH of the buffer solution? pH = Write the net ionic equation for the reaction that occurs when 0.085 mol KOH is added to 1.00 L of the buffer solution. (Use the lowest possible coefficients. Omit states of matter.) 10A. The pKa value for CH3COOH is 4.74. Would a buffer prepared from CH3COOH and CH3COONa with a pH of 4.24 be considered to be an effective buffer? _____ (Yes or No) A buffer in which the mole ratio of NaCH3COO to CH3COOH is 0.49 has a pH of 4.43. Would this buffer solution have a greater capacity for added acid (H3O+) or added base (OH-)? B. A buffer solution that is 0.458 M in HCN and 0.458 M in KCN has a pH of 9.40. The addition of 0.01 mol of OH- to 1.0 L of this buffer would cause the pH to _________ (increase slightly, increase by 2 units, decrease slightly, decrease by 2 units or not change) The capacity of this buffer for added OH- could be increased by the addition of 0.139 mol _________ (of the weak acid, or of the salt). C. A buffer solution that is 0.378 M in HNO2 and 0.378 M in NaNO2 has a pH of 3.35. Addition of which of the following would increase the capacity of the buffer for added OH-? (Select all that apply.) both HNO2 and NaNO2 HNO2 pure water NaNO2 none of the above 11A. A buffer solution contains 0.492 M KHSO3 and 0.372 M K2SO3. If 0.0479 moles of potassium hydroxide are added to 225 mL of this buffer, what is the pH of the resulting solution? (Assume that the volume does not change upon adding potassium hydroxide) B. A buffer solution contains 0.356 M NaHCO3 and 0.347 M Na2CO3. If 0.0608 moles of hydrobromic acid are added to 250. mL of this buffer, what is the pH of the resulting solution ? (Assume that the volume does not change upon adding hydrobromic acid) C. A buffer solution contains 0.292 M acetic acid and 0.443 M sodium acetate. If 0.0188 moles of nitric acid are added to 125 mL of this buffer, what is the pH of the resulting solution ? (Assume that the volume does not change upon adding nitric acid) 12A. A buffer solution contains 0.306 M NaH2PO4 and 0.299 M K2HPO4. Determine the pH change when 0.081 mol KOH is added to 1.00 L of the buffer. pH change = _____ B. A buffer solution contains 0.420 M KHSO3 and 0.256 M K2SO3. Determine the pH change when 0.052 mol HCl is added to 1.00 L of the buffer. pH change = ____ C. Determine the pH change when 0.080 mol HBr is added to 1.00 L of a buffer solution that is 0.378 M in HNO2 and 0.356 M in NO2-. pH change = ______

Adi S.

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