Exam 2 Practice

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Mole, Percent Composition, Percent Abundance

  1. Which one of the samples contains the most particles?

    1. 1 mol CO2(g)
    2. 1 mol UF6(g)
    3. 1 mol CH3COCH3(l)
    4. 1 mol He(g)
    5. all contain the same number of particles

    Solution


  2. Which one of the samples has the largest mass?

    1. 1 mol CO2(g)
    2. 1 mol UF6(g)
    3. 1 mol CH3COCH3(l)
    4. 1 mol He(g)
    5. all contain the same number of particles

    Solution


  3. Which one of the samples contains the most atoms?

    1. 1 mol CO2(g)
    2. 1 mol UF6(g)
    3. 1 mol CH3COCH3(l)
    4. 1 mol He(g)
    5. all contain the same number of particles

    Solution


  4. What is the molar mass (in g mol–1 to four significant figures) of Al2(SO4)3 · 18 H2O?

    Solution


  5. What is the molar mass (in g mmol–1 to three significant figures) of (NH4)2HPO4?

    Solution


  6. How many atoms are present in a 0.268 mol sample of CH3OH?

    Solution


  7. How many aluminum atoms are there in a 3.50 g sample of Al2O3?

    Solution


  8. A sample of a compound containing only carbon and oxygen decomposes and produces 24.5 g of carbon and 32.59 g of oxygen. What is the sample?

    1. CO
    2. CO2
    3. CO3
    4. C3O2
    5. C2O

    Solution


  9. Guanidine, HNC(NH2)2, is a fertilizer. What is the percent by mass (to one decimal place) of nitrogen in the fertilizer?

    Solution


  10. Determine the percent by mass (to one decimal place) of Mg in chlorophyll (C55H72MgN4O5), the green pigment in plant cells.

    Solution


  11. The mineral spodumene has the formula LiAlSi2O6 . What is the mass (in g to two decimal places) of lithium in a 438 g sample?

    Solution


  12. How many moles (in normalized scientific notation) of Cs are contained in 595 kg of Cs?

    Solution


  13. Analysis of a sample of a covalent compound showed that it contained 14.4 % hydrogen and 85.6 % carbon by mass. What is the empirical formula for the compound?

    Solution


  14. What is the mass (in g) of 2.6 × 1022 chlorine atoms?

    Solution


  15. How many iron atoms are contained in 354 g of iron?

    Solution


  16. What is the mass (in ng) of 2.33 × 1020 oxygen atoms?

    Solution


  17. What is the mass (in g) of 2.0 × 1024 mercury atoms?

    Solution


  18. What is the mass (in g) of 2.0 × 1024 mercury atoms?

    Solution


  19. The mineral spodumene has the formula LiAlSi2O6. How many lithium atoms are present in a 105 g sample?

    Solution


  20. Find the mass (in g) of 500 atoms of iron.

    Solution


  21. How much Fe (in mol and number of atoms) are in 125.0 g of Fe?

    Solution


  22. Freon-12 (CCl2F2) is used as a refrigerant in air conditioners and as a propellant in aerosol cans. What is the number of freon-12 molecules and what is the mass (in mg) of Cl in a 5.56 mg sample of freon-12?

    Solution


  23. Prevacid (C16H14F3N3O2S) is used to treat gastroesophageal reflux disease (GERD). Determine each of the following:

    1. the molar mass (in g mol–1) of Prevacid
    2. the mass (in g) of fluorine in a 0.75 mol sample of Prevacid
    3. the number of C atoms in a 0.75 mol sample of Prevacid
    4. the mass (in g) of 4.25 × 1021 molecules of Prevacid

    Solution

    Concept: moles; number of particles


  24. Find the percent composition by mass (to one decimal place) of each element in YBa2Cu3O7.

    Solution


  25. Hemoglobin is a protein that transports oxygen in mammals. Hemoglobin is 0.347 % Fe (by mass). Each hemoglobin molecule contains 4 Fe atoms. What is the molar mass (in g mol–1 in standard notation) of hemoglobin?

    Solution


  26. A compound that only contains carbon, hydrogen, and oxygen is 48.64 % C and 8.16 % H (by mass). What is the empirical formula of this substance?

    Solution


  27. Consider four individual samples of phosphine (PH3), water, hydrogen sulfide, and hydrogen fluoride, each with a mass of 121 g. Rank the compounds from the least to the greatest number of hydrogen atoms contained in each sample.

    Solution


Balancing Equations, Solubility

  1. Write a balanced equation for the following reaction by placing appropriate stoichiometric coefficients.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{CH_3OH} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{O_2} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{CO_2} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2O} \end{align*}\]

    Solution


  2. Write a balanced equation for the following reaction by placing appropriate stoichiometric coefficients.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{CH_3NHNH_2} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{O_2} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{CO_2} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2O} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2} \end{align*}\]

    Solution


  3. Write a balanced equation for the following reaction by placing appropriate stoichiometric coefficients.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{Se} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{BrF_5} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{SeF_6} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{BrF_3} \end{align*}\]

    Solution


  4. Determine the proper value for m and n needed to balance each equation.

    \[\begin{align*} \mathrm{B}_m\mathrm{H}_n + \mathrm{3~O_2} \longrightarrow \mathrm{B_2O_3} + \mathrm{3~H_2O} \end{align*}\]

    Solution


  5. Determine the proper value for m and n needed to balance each equation.

    \[\begin{align*} \mathrm{H}_m\mathrm{IO}_n ~\longrightarrow~ \mathrm{H^+} + \mathrm{IO_4^-} + \mathrm{2~H_2O} \end{align*}\]

    Solution


  6. Write a balanced equation for the following reaction by placing appropriate stoichiometric coefficients.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{CH_4} + \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2O} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{CO} + \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2} \end{align*}\]

    Solution


  7. Write a balanced equation for the following reaction by placing appropriate stoichiometric coefficients.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{Ag_2O(s)} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{Ag(s)} + \rule[-1.0pt]{2em}{0.5pt} \mathrm{O_2} \end{align*}\]

    Solution


  8. Chemical equations must be balanced because the resulting coefficients allow us to predict (select all that apply).

    1. the amount of product that can form from a given amount of reactant.
    2. whether the reaction requires a catalyst or not
    3. how much of one reactant is required to react with a given amount of another
    4. how much reactants are required to form a given amount of products
    5. whether the given reaction is possible or not

    Solution


  9. Chemical equations must be balanced because the resulting coefficients allow us to predict (select all that apply).

    aqueous silver sulfate + aqueous barium iodide → solid barium sulfate + solid silver iodide

    1. Ag2SO4 + BaI2 → BaSO4 + AgI
    2. Ag2SO4(l) + BaI2(l) → BaSO4(s) + 2 AgI(s)
    3. Ag2SO4(aq) + BaI2(aq) → BaSO4(s) + 2 AgI(s)
    4. AgSO4(aq) + BaI(aq) → BaSO4(s) + 2 AgI(s)
    5. AgSO4(l) + 2 BaI(l) → Ba2SO4(s) + 2 AgI2(s)

    Solution


  10. Classify each of the following (all that apply) as a

     I. strong electrolyte
     II. weak electrolyte
     III. nonelectrolyte
     IV. strong acid
     V. strong base
     VI. weak acid
     VII. weak base
     VIII. ionic compound
     IX. organic compound

    1. HBr
    2. ammonium carbonate
    3. NaClO4
    4. ethanol
    5. acetic acid
    6. NH3

    Solution


  11. Which of the following compounds are insoluble in water? Select all that apply.

    1. CoCO3
    2. Cu3(PO4)2
    3. AgNO3
    4. Na2S
    5. AgI

    Solution


  12. Which of the following combinations will form a precipitate? Select all that apply.

    1. SrCl2(aq) + Na2S(aq)
    2. KCl(aq) + CaS(aq)
    3. Hg(NO3)2(aq) + Na3PO4(aq)
    4. Ba (NO3)2(aq) + KOH (aq)
    5. NaOH(aq) + FeCl3(aq)

    Solution


  13. Lead(II) nitrate reacts with sodium chloride. Choose the net ionic equation for the reaction.

    1. Pb2+(aq) + Cl(aq) → PbCl(s)
    2. Pb2+(aq) + NO3(aq) + Na+(aq) + Cl(aq) → PbCl(s) + NaNO3(aq)
    3. PbNO3(aq) + NaCl(aq) → PbCl(s) + NaNO3(aq)
    4. Pb2+(aq) + 2 Cl(aq) → PbCl2(s)

    Solution


  14. Aqueous solutions of sodium sulfate and barium chloride react. What is the sum of the coefficients from the balanced net ionic equation?

    Solution


  15. Answer the questions for the following reaction.

       2 HI(aq) + Ca(OH)2(aq) → 2 H2O(l) + CaI2(aq)

    1. Is the acid strong or weak?
    2. Is the base strong or weak?
    3. What is the net ionic equation for the reaction?

    Solution


  16. A reaction between hydrobromic acid and potassium hydroxide occurs. What is the sum of the coefficients from the balanced net ionic equation?

    1. Is the acid strong or weak?
    2. Is the base strong or weak?
    3. What is the net ionic equation for the reaction?

    Solution


  17. Which is the spectator ion in the reaction between potassium carbonate and calcium iodide? Select all that apply.

    1. K+(aq)
    2. CO32–(aq)
    3. Ca2+(aq)
    4. I(aq)

    Solution


  18. What is the sum of the coefficients of the net ionic equation for aqueous sodium hydroxide neutralized by aqueous acetic acid?

    Solution


  19. When the following solutions are mixed together, what precipitate (if any) will form?

    1. FeSO4(aq) + KCl(aq)
    2. Al(NO3)3(aq) + Ba(OH)2(aq)
    3. CaCl2(aq) + Na2SO4(aq)
    4. K2S(aq) + Ni(NO3)2(aq)
    5. Hg2(NO3)2(aq) + CuSO4(aq)
    6. Ni(NO3)2(aq) + CaCl2(aq)
    7. K2CO3(aq) + MgI2(aq)
    8. Na2CrO4(aq) + AlBr3(aq)

    Solution


  20. Which of the following substances are soluble in water? Select all that apply.

    1. aluminum nitrate
    2. magnesium chloride
    3. rubidium sulfate
    4. nickel(II) hydroxide
    5. lead(II) sulfide
    6. barium hydroxide
    7. iron(III) phosphate

    Solution


  21. Write the net ionic equations for the following reactions:

    1. ammonium sulfate and barium nitrate
    2. lead(II) nitrate and sodium chloride
    3. sodium phosphate and potassium nitrate
    4. sodium bromide and rubidium chloride
    5. copper(II) chloride and sodium hydroxide

    Solution


  22. Write the balanced molecular equation, complete ionic equation, and net ionic equation for the following acid-base reactions.

    1. HClO4(aq) + Mg(OH)2(s)
    2. HCN(aq) + NaOH(aq)
    3. HCl(aq) + NaOH(aq)

    Solution


  23. Write a balanced chemical equation between an acid and a base that would have the following salt appear as a product.

    1. potassium perchlorate
    2. cesium nitrate
    3. calcium iodide

    Solution


Oxidation-reduction

  1. Assign oxidation states for all atoms in each of the following compounds.

    1. KMnO4
    2. NiO2
    3. Na4Fe(OH)6
    4. (NH4)2HPO4
    5. P4O6

    Solution


  2. Assign oxidation states for all atoms in each of the following compounds.

    1. Fe3O4
    2. XeOF4
    3. SF4
    4. CO
    5. C6H12O6

    Solution


  3. Specify which reactions are redox reactions and identify the oxidizing agent, reducing agent, the substance being oxidized, and the substance being reduced.

    1. Cu(s) + 2 Ag+(aq) → 2 Ag(s) + Cu2+(aq)
    2. HCl(g) + NH3(g) → NH4Cl(s)
    3. SiCl4(l) + 2 H2O(l) → 4 HCl(aq) + SiO2(s)
    4. SiCl4(l) + 2 Mg(s) → 2 MgCl2(s) + Si(s)
    5. Al(OH)4(aq) → AlO2(aq) + 2 H2O(l)

    Solution


  4. Determine the half-reactions, the amount (in mol) of electrons transferred, and the overall balanced reaction (with phase labels) for the following redox reaction.

    \[\begin{align*} \mathrm{Fe(s)} + \mathrm{Br_2(l)} \longrightarrow \mathrm{FeBr_3(s)} \end{align*}\]

    Solution


  5. Determine the half-reactions, the amount (in mol) of electrons transferred, and the overall balanced reaction (with phase labels) for the following redox reaction.

    \[\begin{align*} \mathrm{Mn(s)} + \mathrm{F_2(g)} \longrightarrow \mathrm{MnF_4(s)} \end{align*}\]

    Solution


  6. Determine the half-reactions, the amount (in mol) of electrons transferred, and the overall balanced reaction (with phase labels) for the following redox reaction.

    \[\begin{align*} \mathrm{Na(s)} + \mathrm{I_2(g)} \longrightarrow \mathrm{NaI(s)} \end{align*}\]

    Solution


  7. Determine the half-reactions, the amount (in mol) of electrons transferred, and the overall balanced reaction (with phase labels) for the following redox reaction.

    \[\begin{align*} \mathrm{Mg(s)} + \mathrm{Cl_2(g)} \longrightarrow \mathrm{MgCl_2(s)} \end{align*}\]

    Solution


  8. Determine the half-reactions, the amount (in mol) of electrons transferred, and the overall balanced reaction (with phase labels) for the following redox reaction.

    \[\begin{align*} \mathrm{Fe(s)} + \mathrm{O_2(g)} \longrightarrow \mathrm{Fe_2O_3(s)} \end{align*}\]

    Solution


  9. Determine the half-reactions, the amount (in mol) of electrons transferred, and the overall balanced reaction (with phase labels) for the following redox reaction.

    \[\begin{align*} \mathrm{Cr^{3+}(aq)} + \mathrm{Mn^{2+}(aq)} \longrightarrow \mathrm{Mn^{6+}(aq)} + \mathrm{Cr(s)} \end{align*}\]

    Solution


Stoichiometry

  1. Balance the following equation.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{Ag_2O(s)} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{Ag(s)} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{O_2(g)} \end{align*}\]

    Additionally, consider a 4.260 × 103 mg sample of impure silver oxide that, when completely decomposes, yields 283 mg of O2(g). Assuming that the silver oxide is the only source of oxygen, what is the mass percent of silver oxide in the sample?

    Solution


  2. Balance the following equation.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{NH_3} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{Cl_2} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{NCl_3} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{HCl} \end{align*}\]

    Additionally, what mass (in g) of HCl is produced if 1.27 g of NH3 reacts with 4.53 g of Cl2?

    Solution


  3. Balance the following equation.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{C_3H_6} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{NH_3} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{O_2} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{C_3NH_3} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2O} \end{align*}\]

    Additionally, determine the following:

    1. What is the limiting reactant if 4.25 g of C3H6 reacts with 3.14 g of NH3 and 6.12 g of O2?
    2. What mass (in g) of C3NH3 can theoretically be produced given the information in (a)?
    3. What mass (in g) of C3NH3, NH3, and O2 would theoretically be leftover given the information in (a)?
    4. If 1.94 g of C3NH3 was produced in an experiment, what is the percent yield?

    Solution


  4. Balance the following equation.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{C_6H_5NO_2} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{C_6H_{14}O_4} ~\overset{\mathrm{Zn}}{ \underset{\mathrm{KOH}}{\longrightarrow}}~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{(C_6H_5N)_2} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{C_6H_{12}O_4} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{H_2O} \end{align*}\]

    Additionally, determine the following:

    1. What volume (in mL) of C6H5NO2 (ρ = 1.20 g mL–1) must be allowed to react with an excess of C6H14O4 to produce 6.21 g of (C6H5N)2 if the percent yield is 83.7 %?
    2. If 0.17 L of C6H5NO2 (ρ = 1.20 g mL–1) and 0.52 L C6H14O4 (ρ = 1.12 g mL–1) react to yield 64.4 g of (C6H5N)2, what is the limiting reactant and what is the percent yield of the reaction?

    Solution


  5. Balance the following equation.

    \[\begin{align*} \rule[-1.0pt]{2em}{0.5pt} \mathrm{Fe} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{Br_2} ~\longrightarrow~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{NaBr} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{CO_2} ~+~ \rule[-1.0pt]{2em}{0.5pt} \mathrm{Fe_3O_4} \end{align*}\]

    Additionally, an alloy contains 84.2 % Fe and 15.8 % Ni (by mass). A 6.41 g sample of the alloy reacts with 8.30 g Na2CO3 with excess Br2.

    1. What mass (in g) of Fe3O4 can be produced?
    2. What mass (in g) of Fe3O4 is produced if the percent yield of the reaction is 94.3 %?

    Solution


  6. A mixture contained no fluorine compound except methyl fluoroacetate, FCH2COOCH3 (M(FCH2COOCH3) = 92.08 g mol–1). When chemically treated, all the fluorine was converted to CaF2 (M(CaF2) = 78.08 g mol–1). The mass of CaF2 obtained was 20.1 g. Find the mass (in g) of methyl fluoroacetate in the original mixture.

    Solution


  7. A 1.62 g sample of a metal chloride, MCl2, is dissolved in water and treated with excess aqueous silver nitrate. The silver chloride that formed weighed 3.46 g. Calculate the molar mass (in g mol–1) of the metal, M, and identify the metal.

    Solution


  8. Which solution has the greatest molar concentration of SO42–?

    1. 0.060 M H2SO4
    2. 0.27 M MgSO4
    3. 0.17 M Na2SO4
    4. 0.098 M Al2(SO4)3
    5. 0.22 M CuSO4

    Solution


  9. A solution is prepared by dissolving 4.25 g NaCl, 0.175 g KCl, and 0.183 g CaCl2 in water. The volume of the solution is 500.0 mL. What is the molar concentration (in mol L–1) of Cl in the solution?

    Solution


  10. In the following reaction, 55.2 mL of potassium sulfate solution was added to excess lead acetate. What is the concentration of K+ in the potassium sulfate solution if 1.23 g of PbSO4 was produced?

    \[\begin{align*} \mathrm{K_2SO_4(aq)} + \mathrm{Pb(C_2H_3O_2)_2(aq)} \longrightarrow \mathrm{2~KC_2H_3O_2(aq)} + \mathrm{PbSO_4(s)} \end{align*}\]

    Solution


  11. You mix 285.0 mL of 1.20 M aqueous lead(II) nitrate with 300.0 mL of 1.55 M aqueous potassium iodide. Determine the following.

    1. the molecular equation for this reaction
    2. the limiting reactant
    3. the final molar concentration (in mol L–1) of Pb2+
    4. the mass (in g) of lead(II) iodide formed
    5. the final molar concentration (in mol L–1) of K+
    6. the final molar concentration (in mol L–1) of NO3

    Solution


  12. If all of the chloride in a 4.106 g sample of an unknown metal chloride is precipitated as AgCl with 70.90 mL of 0.2010 M AgNO3, what is the percentage of chloride in the sample?

    Solution


  13. If all of the chloride in a 4.106 g sample of an unknown metal chloride is precipitated as AgCl with 70.90 mL of 0.2010 M AgNO3, what is the percentage of chloride in the sample?

    Solution


  14. A mixture of BaCl2 and NaCl is analyzed by precipitating all of the barium as BaSO4. After the addition of excess Na2SO4 to a 3.988 g sample of the mixture, the mass of precipitate collected is 2.113 g. What is the mass percentage of barium chloride in the mixture?

    Solution


  15. A 3.00 g sample of an alloy containing only Pb and Sn was dissolved in nitric acid. Sulfuric acid was added to this solution, which precipitated 1.90 g of PbSO4. Assuming that all of the lead was precipitated, what is the percentage of Sn in the sample? (M(PbSO4) = 303.26 g mol–1)

    Solution


  16. You have 76.0 mL of a 2.50 M aqueous solution of Na2CrO4 and 125 mL of a 2.16 M aqueous solution of AgNO3. Calculate the molar concentration (in mol L–1) of CrO42– after the two solutions are mixed together.

    Solution


  17. You have a 75.0 mL 2.50 M aqueous Na2CrO4 solution and 125 mL 2.29 M aqueous AgNO3 solution. What is the molar concentration (in mol L–1) of Ag+ after the two solutions are mixed together?

    Solution


  18. You have 75.0 mL of a 2.50 M aqueous solution of Na2CrO4 and 125 mL of a 2.24 M aqueous solution of AgNO3. Calculate the molar concentration (in mol L–1) of NO3 after the two solutions are mixed together.

    Solution


  19. Combine a 55 mL 1.00 M aqueous silver nitrate solution with a 25 mL 0.55 M silver chloride solution. What mass (in g) of silver chloride is produced?

    Solution


  20. A 0.685 g sample of an unknown diprotic acid requires a 42.57 mL 0.111 M aqueous NaOH solution to be completely neutralized. What is the molar mass (in g mol–1) of the acid?

    Solution


  21. What mass (in g) of NaOH is required to completely react with a 25.0 mL 2.2 M aqueous H2SO4 solution?

    Solution


  22. What volume (in mL) of a 5.00 M hydrofluoric acid solution will completely react with 4.05 g of calcium hydroxide?

    Solution


  23. Sulfamic acid (HSO3NH2) is a strong monoprotic acid that can be used to standardize a strong base. A 0.179 g sample of HSO3NH2 is required to completely neutralize a 19.4 mL aqueous KOH solution. What is the molar concentration (in mol L–1) of the KOH solution?

    Solution


  24. A student weighs out 0.556 g of KHP (M(KHP) = 204.22 g mol–1) and puts it into 36.78 mL of a stock NaOH solution. If that was enough to neutralize the NaOH, what is the concentration of the stock NaOH solution? KHP is a monoprotic acid.

    Solution


  25. A 2.80 g sample of phosphoric acid is added to a 150.0 mL 1.00 M sodium hydroxide solution to give a 151.489 mL mixture and the acid is completely neutralized. Determine the following:

    1. [Na+] (in mol L–1)
    2. [PO43–] (in mol L–1)
    3. [OH] (in mol L–1)

    Solution


  26. A stock solution with a total volume of 1000.0 mL contains 37.1 g Mg(NO3)2. If you take a 20.0 mL aliquot and then dilute it with water to a total volume of 500.0 mL, what is the molar concentration (in g mol–1) of Mg2+ and NO32– in the final solution?

    Solution


  27. Determine the molar concentrations (in g mol–1) the ions present in a solution created from mixing equal volumes of 1.0 M aqueous lead(II) nitrate and 1.0 M aqueous sodium chloride solutions? Assume that the volumes are precise to one decimal place in normalized scientific notation.

    Solution


  28. Determine the molar concentrations (in g mol–1) the ions present in a solution created from mixing equal volumes of 1.0 M aqueous ammonium carbonate and 1.0 M aqueous potassium perchlorate. Assume that the volumes are precise to one decimal place in normalized scientific notation.

    Solution


  29. Determine the molar concentration (in g mol–1) of the salt produced by a reaction between a 200. mL 0.100 M aqueous HCl solution with a 100. mL 0.50 M aqueous KOH solution.

    Solution


  30. What volume (in mL) of a 0.100 M aqueous HNO3 solution is required to neutralize a 50.0 mL 0.150 M aqueous Ba(OH)2 solution?

    Solution


  31. Hydrogen cyanide is produced industrally from a reaction between gaseous ammonia, oxygen, and methane.

    \[\begin{align*} \mathrm{2~NH_3(g)} + \mathrm{3~O_2(g)} + \mathrm{2~CH_4(g)} \longrightarrow \mathrm{2~HCN(g)} + \mathrm{6~H_2O(l)} \\ \end{align*}\]

    If 5.00 × 103 kg of each reactant react, what mass (in kg) of each product would be produced (assuming a 100 % yield)?

    Solution


  32. Acrylonitrile (C3H3N) is the starting material for many synthetic carpets and fabrics and is produced by the following reaction.

    \[\begin{align*} \mathrm{2~C_3H_6(g)} + \mathrm{2~NH_3(g)} + \mathrm{3~O_2(g)} \longrightarrow \mathrm{2~C_3H_3N(g)} + \mathrm{6~H_2O(g)} \\ \end{align*}\]

    If 15.0 g C3H6, 5.00 g NH3, and 8.00 g O2 react, what mass (in g) of acrylonitrile can be produced (assuming a 100 % yield)?

    Solution


  33. Calcium chloride is a strong electrolyte and is used to ``salt’’ streets in the winter to melt ice and snow. Write a net ionic reaction to show how this substance breaks apart when it dissolves in water.

    Solution


  34. A solution of ethanol in water is prepared by dissolving 75.0 mL of ethanol (ρ = 0.79 g cm–3) in enough water to make a 250.0 mL solution. What is the molar concentration (in mol L–1) of the ethanol in this solution?

    Solution


  35. Which of the following aqueous solutions contains the largest number of ions?

    1. 100.0 mL of 0.100 M NaOH
    2. 50.0 mL of 0.200 M BaCl2
    3. 75.0 mL of 0.150 M Na3PO4

    Solution


  36. If 12.0 g of AgNO3 is available, what volume (in L) of 0.25 M AgNO3 can be prepared?

    Solution


  37. A solution is prepared by dissolving 10.8 g ammonium sulfate in enough water to make 100.0 mL of stock solution. A 10.00 mL aliquot is taken and 50.00 mL of water is added. What is the molar concentration (in mol L–1) of ammonium ions and sulfate ions in the final solution?

    Solution


  38. What mass (in g) of Na2CrO4 is required to precipitate all of the silver ions from a 75.0 mL 0.100 M aqueous solution of AgNO3?

    Solution


  39. What mass (in g) of iron(III) hydroxide precipitate can be produced by reacting a 72.0 mL 0.105 M aqueous iron(III) nitrate solution with a 125 mL 0.150 M aqueous sodium hydroxide solution?

    Solution


  40. A 100.0 mL 0.200 M aqueous potassium hydroxide solution is mixed with a 100.0 mL 0.200 M aqueous magnesium nitrate solution.

    1. Write a balanced chemical equation for the reaction that occurs.
    2. Determine the precipitate that forms (if any).
    3. Determine the mass (in g) of precipitate that forms (if any).
    4. Determine the molar concentration (in g mol–1) of each ion in solution after the reaction goes to 100 % completion.

    Solution


  41. A 1.42 g sample of a pure, metal (M) containing compound (M2SO4) was dissolved in water and treated with an excess of aqueous calcium chloride. All the sulfate ions precipitated as calcium sulfate which was collected, dried, and found to be 1.36 g. What is the identity and standard atomic weight of the metal?

    Solution


  42. What volume (in mL) of each of the following bases will will completely react with 25.0 mL of 0.200 M HCl?

    1. 0.100 M NaOH
    2. 0.0500 M Sr(OH)2
    3. 0.250 M KOH

    Solution


  43. A 25.0 mL sample of HCl(aq) requires 24.16 mL of 0.106 M NaOH for complete neutralization. What is the molar concentration (in mol L–1) of the original HCl(aq) solution?

    Solution


  44. 5.00 g of barium chloride was added to 225 mL of a 1.40 M solution of sodium sulfide.

    1. Write the balanced molecular equation and include phase labels.
    2. Write the full ionic equation and include phase labels.
    3. Write the net ionic equation and include phase labels. If there is no net ionic equation, write “no net ionic equation.”
    4. Indicate the pricipitate (if any).
    5. Determine the limiting reactant.

    Solution


  45. A precipitate forms when titanium(IV) chloride is added to water. Two water molecules react and form four HCl molecules. What is the identity of the precipitate?

    1. What is the identity of the precipitate?
    2. What is the molar concentration (in mol L–1) of H+ ions if 2.00 g of TiCl4 was added to enough water to give a 100.0 mL solution?

    Solution


  46. A 123 mL sample of 0.210 M aqueous magnesium chloride forms a precipitate when mixed with 324 mL 0.120 M aqueous sodium hydroxide.

    1. How much (in g) precipitate is formed?
    2. What is the molar concentration (in mol L–1) of the magnesium(2+) ion?
    3. What is the molar concentration (in mol L–1) of the sodium(1+) ion?

    Solution


  47. If 275 mL of a 0.125 M aqueous NaCl solution and 375 mL of a 0.575 M aqueous Na2SO4 solution are mixed, determine the molar concentrations (in mol L–1) of the following:

    1. chloride ions
    2. sulfate ions
    3. sodium ions

    Solution


Thermochemistry Basics

  1. Sugar is melted in a pot and its temperature is measured as it heats. In this scenario, what is the system?

    1. the pot and sugar
    2. the stove
    3. the entire kitchen
    4. the rest of the universe

    Solution


  2. Which of the following is not a type of potential energy (select all that apply)?

    1. Energy held in chemical bonds
    2. Energy resulting from intramolecular attractions
    3. Energy from the random motion of molecules
    4. Energy of a ball dropping from a height

    Solution


  3. Which of the following is true of heat (select all that apply)?

    1. Heat is a form of thermal energy.
    2. Heat is the transfer of thermal energy.
    3. Heat is the action of forces through a distance.
    4. A negative heat in the system means the surroundings loses energy.
    5. A negative heat in the system means the system loses energy.

    Solution


  4. Which of the following are false (select all that apply)?

    1. Energy can be converted from one type to another.
    2. Energy is the capacity to do work.
    3. Kinetic energy is energy resulting from condition, position, or composition.
    4. Potential energy is energy transferred between a system and its surroundings as a result of a temperature difference.

    Solution


  5. A block of ice absorbs heat and melts. The value q for the system is:

    1. Positive
    2. Negative
    3. Zero
    4. There is not enough information to determine.

    Solution





Questions are written by UGA Chemistry unless otherwise indicated and translated with minor tweaks into an online format by Eric Van Dornshuld.