Grade 11 Chemistry Exam Review Answers

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1 Grade 11 Chemistry Exam Review Answers 1. On the periodic table below, label the following groups of elements: metals, non-metals, metalloids, transition metals, halogens, noble gases, alkali metals, and alkaline earth metals. 2. Fill in the following chart with the number of electrons, protons, neutrons, atomic number, mass number, and the element notations Atom Atomic Number Mass Number # of Protons # of neutrons K P Al As W P # of electrons 3. Draw Bohr diagrams and Lewis dot diagrams for the following elements and ions: Li, O, and F 4. Draw Bohr diagrams of the following IONS: Mg2+, B3+, and N3-

2 5. Consider the following electron configuration for neutral atoms: i) 1s 2 2s 2 2p 6 3s 2 3p 5 Cl ii) 1s 2 2s 2 2p 6 iii) 1s 2 2s 2 2p 6 3s 2 3p 6 Ar iv) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 v) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 K Ne Ca a) Which of these would you expect to be an alkaline earth metal? V b) List the five atoms in order of increasing atomic radii? Ne, Cl, Ar, K, Ca, c) Which of these would you expect to have the lowest ionization energy? V d) Which of the following would you expect to be in the same family? II and III 6. Consider the following atoms: 16 8 X, 17 9Y, and 17 8Z Which of these atoms (if any) are isotopes of the same element? Explain your reasoning. X and Z are isotopes, because they have the same number of protons, but differ number of neutrons. 7. Find the average atomic mass of P when P-30 is in 77% abundance and the P- 32 is in 23% abundance. AAM = (30 x 0.77) + (32 x 0.23) = amu 8. Use the following given below to calculate the average atomic mass of element X ISOTOPE ABUNDANCE MASS (amu) 302 X X X Average atomic mass = a.m.u 9. What is ionization energy, electron affinity, and electronegativity? Ionization Energy: The energy required to remove an electron from a gaseous atom Electron Affinity: The energy that accompanies the addition of an electron to a gaseous atom Electronegativity: The tendency of an atom to attract electrons to itself when it is chemically combined with another element. 10. All of the trends in the periods table can be explained by two major factors: Effective nuclear charge, an distance of electrons from nucleus. For the following elements decide which one has a larger atomic radius, electronegativity, ionization energy, and reactivity. Largest Atomic Radius Largest Electronegativity Largest Ionization Energy Greatest Reactivity O Po Po O O O B N B N N N Ni Pt Pt Ni Ni Pt Rationale/Justification for choice. Po electrons further from nucleus, easier to remove electrons N is holding electrons closets due to greater effective nuclear charge Ni is holding electrons more tightly

3 A u Cu Au Cu Cu Au As N As N N N I Cl I Cl Cl Cl Au has e - at greater distances, thus more willing to lose e - N is holding onto e - more tightly, more greedy for electrons I is holding onto e - more tightly, more greedy for electrons 11. Describe what happens to the atomic radius and effective nuclear charge across a row and down a group. As you move to the right on the periodic table, you get more protons in the nucleus which will have a greater pulling force for the electrons, bringing them closer to the nucleus, therefore making a smaller atomic radius to the right. 12. Describe how the reactivity is linked to ionization energy for the metals and electron affinity for the non- metals. The lower the ionization energy, the more an atom wants to get rid of its electron, therefore the faster it will react making a greater reaction. Therefore the Lower the ionization energy, the greater the reactivity. 13. What is the octet rule and what does it mean for any element that has its octet full? All atoms are striving for a full outer orbital of electrons. Therefore, atoms will lose or gain electrons to get a full outer orbital of electrons 14. What is the bonding capacity of the following elements: P, S, Cl, N, O, and F 3,2,1,3,2 15. Define the following terms: alloy, metalloid, solution, hydrate Alloy: solid solution of metals Metalloid: Semi-metals - "staircase" line from the top of Group 13 to the bottom of Group 16 Solution: a homogeneous mixture of a solvent and one or more solutes Hydrate: a compound that has a specific number of water molecules bonded to each formula unit 16. Describe the difference between ionic bonds and covalent bonds (non-polar and polar covalent). Be sure to include which type of elements are joined together for each type of bond, how strong the bond is, if the electrons are either transferred or shared. Ionic Bonds Between a Metal (or NH4+) and a Non-Metal Bond Transfer of Electrons Covalent Bonds Between 2 Non-Metals Bond Shares Electrons Polar (electrons are shared unequally) 17. Determine the bond type in the following compounds using electronegativity values, and indicate if they are ionic or covalent bonds and draw the appropriate Lewis structures. a) LiF : Ionic Bond b) F2 : Covalent c) MgS : Ionic Bond d) CO2 : Covalent 18. Describe polar covalent, non-polar covalent, pure covalent, and ionic bond with respect to electro-negativities. Pure covalent =0, non-polar covalent 0< <0.4, polar covalent 0.4< <1.7, ionic bond > Distinguish fully between polar and non-polar substances, and their ability to dissolve in different solvents. Polar dissolves in polar, Non-Polar dissolves in Non-Polar. Polar molecules have an overall difference in the charge while a non-polar molecule will not have a difference in charge.

4 20. Name the following compounds: NaF sodium fluoride LiBr lithium bromide CoO cobalt (ii) oxide Sn(O2)2 tin (iv) peroxide MgH2 magnesium hydride SbBr5 antinomy pentabromide SO2 sulfur dioxide NaBrO2 sodium bromite Li2O2 lithium peroxide HgO mercury (ii) oxide AgOH silver hydroxide CS2 carbon disulfide HBr(aq) hydrobromic acid HBrO2(aq) bromous acid AlN aluminum nitride Ba(ClO4)2 barium perchlorate Cu(NO3)2 copper (ii) nitrate Na2CO2 sodium carbonite HBrO3(aq) bromic acid Mg(MnO2)2 magnesium permanganate LiH2PO4 lithium dihydrogen phosphate LiCN lithium cyanide 21. Write the chemical formula for the following: sodium bromide NaBr copper (ii) nitrate Cu(NO3)2 ferric hydroxide Fe(OH)3 potassium sulfite K2SO3 barium carbonate BaCO3 manganese (iv) oxide MnO2 silver perchlorate AgClO4 hypobromous acid HBrO(aq) aluminum hypobromite Al(BrO)3 lithium phosphite Li3PO3 nikel (ii) sulfate NiSO4 dihydrogen monoxide H2O carbon tetrachloride CCl4 zinc hydrogen phosphate ZnHPO4 antinomy (v) perbromate Sb(BrO4)5 magnesium dichromate MgCr2O7 lithium manganate Li2MnO4 bimuth (v) sulfide Bi2S5 22. Distinguish between intermolecular and intramolecular forces. Intermolecular forces are forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions); i.e dipole-dipole, hydrogen bonding, London dispersion. They are weak compared to the intramolecular forces, the forces which keep a molecule together (i.e covalent and ionic bonds). 23. In an experiment designed to compare the reactivity of 4 metals (W, X,Y, and Z), the metallic elements were reacted with aqueous solutions containing their ions and an anion A. The following observations were made when the single displacement reactions were attempted. a) Z(s) + WA(aq) ZA(aq) + W(s) b) W(s) + XA(aq) NR c) X(s) + ZA(aq) XA(aq) + Z(s) d) W(s) + YA(aq) WA(aq) + Y(s) Y< W<X<Z 24. Balance the following equations. a) 2 Na + 2 H2O 2 NaOH + H2 b) CaCO3 + 2 HCl CaCl2 + H2O + CO2 c) 5 FeCl2 + KMnO4 + 8 HCl 5 FeCl3 + KCl + MnCl2 + 4 H2 O

5 d) 2 KClO3 2 KCl + 3 O2 e) C3H8 + 5 O2 3 CO2 + 4 H2O f) 2 Al + 3 H2SO4 Al2(SO4)3 + 3 H2 g) 2 Fe2O3 + 3 C 4 Fe + 3 CO2 h) CaC2 + 2 H2O Ca(OH)2 + C2H2 25. Define the law of conservation of mass and the law of definite proportions In any physical or chemical reaction, matter is not created nor destroyed, it is conserved. In any chemical compound, the elements are always combined in the same proportion by mass. 26. What type of reaction are the following equations. a) 2 Na + 2 H2O 2 NaOH + H2 Single Displacement b) Mg + 2 HCl MgCl2 + H2 Single Displacement c) 2 KClO3 2 KCl + 3 O2 Decomposition d) C3H8 + 5 O2 3 CO2 + 4 H2O Combustion e) 2 Al + 3 H2SO4 Al2(SO4)3 + 3 H2 Single Displacement f) 2 Fe2O3 + 3 C 4 Fe + 3 CO2 Single Displacement g) 2 Fe(OH)3 + 3 H2SO4 Fe2(SO4)3 + 6 H2O DD, Neutralization h) 3 H2 + N2 2 NH3 Synthesis i) CH3OH + O2 CO2 + 2 H2O Combustion j) BaSO4 + 2 NaCl BaCl2 + Na2SO4 Double Displacement k) 2 H2 + O2 2 H2O Synthesis 27. Complete the following reactions. Remember to balance them and indicate states. a) Fe(s) + CuSO4 (aq) FeSO4 (aq) + Cu (s) b) Cu (s) + AgNO3 (aq) CuNO3 (aq) + Ag (s) c) Mg (s) + 2 HCl (aq) MgCl2 (aq) + H2 (g) d) Na2CO3 (aq) + CuSO4 (aq) CuCO3 (s) + 2 Na2SO4 (aq) e) (NH4)2SO4 (aq) + CaCl2 (aq) CaSO4 (s) + 2 NH4Cl (aq) f) Ag + CuSO4 No Reaction g) Mg (s) + H2SO4 (aq) MgSO4 (aq) + H2 (g) h) HCl (aq) + NaCO3 (aq) NaCl (aq) + CO2 (g) + H2O (l) i) 3 CaS + Ni2(SO4)3 3 CaSO4 + Ni2S3

6 28. For the following pairs of aqueous solutions write the balanced chemical equation, total ionic equation and net ionic equation. a) sodium sulfide and iron (II) sulfate Na2S(aq) + FeSO4(aq) Na2SO4(aq) + FeS(s) Na+(aq) + S2-(aq) + Fe2+(aq) + SO42-(aq) Na+(aq) + SO42-(aq) + FeS(s) Fe2+(aq) + S2-(aq) FeS(s) b) ammonium iodide and silver nitrate NH4I(aq) + AgNO3(aq) NH4NO3(aq) + AgI(s) NH4+(aq) + I-(aq) + Ag+(aq) + NO3-(aq) NH4+(aq) + NO3-(aq) + AgI(s) Ag+(aq) + I-(aq) AgI(s) c) potassium carbonate and iron (II) nitrate K2CO3(aq) + Fe(NO3)2(aq) 2KNO3(aq) + FeCO3(s) 2K+(aq) + CO32-(aq) + Fe2+(aq) + 2NO3(aq) 2K+(aq) + 2NO3(aq) + FeCO3(s) Fe2+(aq) + CO32-(aq) FeCO3(s) 29. Distinguish fully between the following terms: a) Limiting reagent and excess reagent b) Actual yield and theoretical yield c) Empirical formula and molecular formula 30. What is the molar mass of the following? a) NH3 = (1.01) b) AgNO3 = (16) d) C6H12O6 = 6(1.01) + 12(1.01) + 6(16) f) Mg(BrO4)2 = ( (16)) = g/mol = g/mol = g/mol = g/mol 31. How many moles of AlBr3 are there if there is: recall. n=m/m a) 4.6 grams n=4.6/266.7 n= mol b) 3.9 grams n=3.9/266.7 n= mol c) 2.3 X 106g n=2.3 X 106/266.7 n=8625 mol 32. How many grams of B2S3 is there if there is: a) 3.8 moles m=nm n=3.8 X m=447 g b) 2.9 moles m=nm n=2.9 X m=341 g c) 7.3 X 104 m=nm n=7.3x104 X m=8.59 X 106 g 33. How many atoms of He are there if there are: a) 4.7 moles =4.7 x 6.022x1023 =2.83 X 1024 atoms b) 2.9 moles =2.9 x 6.022x10 23 =1.75 X 1024 atoms c) 8.6 moles =8.6 x 6.022x10 23 =5.18 X 1024 atoms 34. How many atoms of O is there in: a) 7.8 moles of AgNO3 =7.8 X 6.022x1023 X 3 =1.4 X 1025 atoms b) 9.2 moles of CO2 =9.2 X 6.022x10 23 x 2 =1.1 X 1025 atoms c) 3.3 moles of O2 =3.3 X 6.022x10 23 X 2 =4.0 X 1024 atoms 35. Calculate the percentage by mass of each element in the following compounds. a) CH4 %C=1(12)/(1(12)+4(1)) X 100% %H=4(1)/(1(12)+4(1)) X 100% b) C3H8 %C=3(12)/(3(12)+8(1) X 100% %H=8(1)/(3(12)+8(1) X 100% =74.82% =25.17% =81.68% =18.32%

7 c) NaHCO3 %Na=23/( (16)) X 100% %H=1/( (16)) X 100% %C=12/( (16)) X 100% %O=3(16)/( (16)) X 100% =27.37% =1.20% =14.30% =57.14% 36. Ammonium nitrate, NH4NO3, and ammonium sulphate, (NH4)2SO4, are both used as fertilizers. Show by calculation which compound has the greater percentage by mass of nitrogen. NH4NO3 %N=2(14)/(14+4(1)+14+3(16) X 100% =35.00% (NH4)2SO4 %N=2(14)/(2(14)+8(1)+32+4(16) X 100% =21.20% 37. Determine the empirical formula of each of the following compounds. The percentage composition by mass is given. a) 85.7% carbon, 14.3% hydrogen Assume 100g C n=m/m =85.7/12.01 = H n=m/m =14.3/1.01 = Divide both by CH2 b) 62.6% lead, 8.4% nitrogen, 29.0% oxygen Assume 100g Pb n=m/m =62.6/207.2 =0.302 N n=m/m =8.4/14.01 =0.6 O n=m/m =29/16 =1.81 Divide by PbN2O6 38. Chemical analysis of one of the constituents of gasoline showed that it consists of 92.29% C and 7.71% H by mass. If the molecular mass of the compound is 78 g/mol, determine its molecular formula. Assume 100g C n=m/m =92.29/12.01 =7.7 H n=m/m =7.71/1.01 =7.63 Divide by 7.63 CH 78 = X ( ) X=6 C 6H6 39. A compound of silicon and fluorine was analyzed and found to consist of 33.0% silicon and 67.0% fluorine by mass. The molecular mass of the compound was determined by experiment to be 170 g/mol. What is the molecular formula of the compound? Assume 100g Si n=m/m =33/28.09 =1.17 F n=m/m =67/19 =3.53 Divide by 1.17 SiF3 170 = X ( (19)) X=2 Si2F6 40. A crystal is added to a solution, Describe what happens to the crystal if the solution is unsaturated, saturated, and supersaturated. Unsaturated: crystal will dissolve, saturated: crystal will remain undissolved, supersatured: will disturbed equilibrium and excess salt will crystalize

8 41. Consider the following solubility curves. a) What is the solubility of KNO3, at 50 oc? 80g/100 gh2o b) What must the temperature be to create a saturated solution of KNO3, using 50 g of the salt and 50 g of water? >60 oc c) What mass of KNO3 will precipitate if a saturated solution of KNO3 at 60 oc is cooled to 10 oc? ~ 80g 42. Finish the following equations taking into account the solubility rules. Circle all spectator ions (the ones that remain as ions) a) NaI + Pb(NO3)2 PbI2 + 2Na+ + 2NO3 - b) CaCl2 + NaCO3 CaCO3 + 2Na+ + 2Cl - c) BaBr2 + Na2SO4 BaSO4 + 2Na+ + 2Br - d) CuSO4 + K2CO3 CuCO3 + 2K+ + SO4 2- e) NH4OH + FeI2 Fe(OH)2 + 2NH4+ + 2I g of sodium chloride is dissolved in 100 g of water. What is the concentration in: a) % by mass 25 g/125 g x 100 = 20 % m/m b) mol/l (density of solution is 1.15 g/ml) 25 g x (I mol/58.44g) = mol 125 x (1 ml/1.15g) = ml Thus, molar concentration = mol / L = 3.9 mol/l 44. What is the concentration of a solution of 3.7 moles of CH4? a) 1L c=n/v =3.7/1 =3.7 mol/l b) 7L c=n/v =3.7/7 =0.529 mol/l c) 450 ml c=n/v =3.7/0.45 =8.22 mol/l 45. What is the volume of 5.8 moles of a solution of NaCl? a) 4.5 moles/l V=n/c =5.8/4.5 =1.29 L b) 2.9 moles/l V=n/c =5.8/2.9 =2 L 46. What is the concentration of a solution of 8L of 4.7 moles/l, if it is diluted to 200L?. C 1 V 1 =C 2 V 2 (4.7)(8) = (200)c2 =0.188 mol/l 47. What is the concentration of a solution of 200 ml of 3.2 mol/l, if it is diluted to 1L. c1v1=c2v2 (0.2)(3.2) = (1)c2 =0.64 mol/l 48. What is the concentration of sulfate ions in the solution that has 4.0 g of ammonium sulfate dissolved in 350 ml of water? (NH4)2SO4 2 NH4+ + SO4 2- moles of ammonium sulfate: n=m/m n= 4.0g/ 132g/mol n = mol moles of ammonium sulfate = moles of sulate c = n/ V c = mol/ L c = mol/l = 0.087M

9 49. What mass of lead (II) nitrate is required to make 4.5 L of a 0.75 mol/l solution? 4.5 L x 0.75M x g/mol = 1.1 kg 50. The water in a swimming pool has tested positive for lead at a concentration of 4.9 ppm. What mass of lead would there be in a 60,000 L pool? (hint: Density of water????) 4.9 ppm = (x/60000 g) x 100% x = 2940 g 51. Calculate the concentrated (in mol/l) of a 7.50% by mass acetic acid solution if the density is 1.02 g/ml. Since, (7.50 g/100g) x 100 = 7.4 % then, 1 gram of solution has 7.50 g acid, and 92.5 g solvent) g x (1 mol/60.06 g) = mol, assuming 100 g of solution, 100 g x (1 ml/1.05 g) = ml = L C = moles of solute/volume of solution = / L = 1.3 M 52. Concentrated phosphoric acid has a concentration of 18.0 M. What volume of concentrated phosphoric acid is needed to make 8.00 L of 1.50 mol/l solution? 0.667M 53. What is the difference between an Arrhenius acid/base and a Bronsted-Lowry acid/base. Arrhenius Acid: Any substance that increases the hydronium ion concentration in neutral water Arrhenius Base: Any substance that increases the hydroxide ion concentration in neutral water Bronsted-Lowry Acid: Any substance that acts as a proton donor Bronsted-Lowry Base: Any substance that acts as a proton acceptor 54. Write an equation that demonstrates how pure water can act as both as a Bronsted- Lowry Acid and Bronsted-Lowry base at the same time. 55. List three properties of acids, and three properties of bases. 56. Distinguish fully between strong acids and weak acids. Strong acids: fully ionize, high concentration of H3O +, low ph ~1-3 Weak acids: partially ionize, lower concentration of H3O +, slightly higher ph ~ If you perform an acid base titration and place 50 ml of 0.25 mol/l acetic acid into an Erlenmyer and read the burette initially at 3.5 ml and at the endpoint at 25.2 ml, what is the concentration of the base? cava=cbvb (50)(0.25) = ( )cB =0.58 mol/l 58. If the hydronium ion concentration of a solution is 2.5 X 10-6, what is the hydroxide ion concentration, the ph, and the poh? [H+][OH-]=1 x [OH-]=4 x 10-9 mol/l ph = -log[h3o+] poh = -log[oh-] ph = 5.6 poh = If the hydroxide ion concentration of a solution is 4.7 X 10-4, what is the hydronium ion concentration, the ph, and the poh.? [H+][OH-]=1 x [H+]=2.13 x mol/l ph = -log[h3o+] poh = -log[oh-] ph = poh = If the ph of a solution is 3.4, what is the poh, the hydroxide concentration, and the hydronium ion concentration? ph + poh = 14 poh = 10.6 [H+]=10-pH [H+] = 3.98 x 10-4 mol/l [OH-]=10-pOH [OH-] = 2.5 x mol/l

10 61. If the poh of a solution is 8.4, what is the ph, the hydroxide concentration, and the hydronium ion concentration? ph + poh = 14 ph = 5.6 [H+]=10-pH [H+] = 2.5 x 10-6 mol/l [OH-]=10-pOH [OH-] = 3.98 x 10-9 mol/l 62. Describe Boyle's Law, Charle's Law, and the Pressure Temperature Law. Boyle's Law: Pressure is inversely proportional to Volume P1V1=P2V2 Charle's Law: Volume is directly proportional to Temperature V1/T1=V2/T2 Pressure-Temperature Law: Pressure is directly proportional to Temperature P1/T1=P2/T2 63. What is the difference between STP and SATP? STP: Standard Temperature and Pressure P=101.3 kpa, T= 0oC SATP: Standard Ambient Temperature and Pressure P=100 kpa, T=25oC 64. What is the volume of 1 mol of a gas at STP (Molar Volume)? 22.4 L 65. Calculate the number of: a) Moles in 45.2 L of carbon dioxide gas at STP: 2.00 mol b) Molecules in 45.2 L of cabon dioxide gas at STP: 1.22 x c) Moles in 45.2 L of oxygen gas at STP: 2.00 mol d) Moles of 45.2 L of carbon dioxide gas at 25 oc and 97.0 kpa: 1.77 mol e) Volume of 45.2 g of carbon dioxide gas at STP: 1.03 mol x 22.4 L/mol = 23.1 L f) Volume of 45.2 g of oxygen gas at STP: 1.41 mol x 22.4 l/mol = 31.6 L 66. What is the density of methane gas (CH4) at 20 oc and 115 kpa? D= m/v, assuming 1 mol and using PV= nrt n= (1 mol)(8.314)(293k)/115 kpa = 21.2 L D = g/21.2 L = g/l 67. What is the pressure of He if it begins at 78oC with a pressure of 56 kpa and then gets changed to a temperature of 109oC? P1/T1=P2/T2 P1/ = 56/ P1=60.5 kpa 68. What is the Temperature of O2 gas if it starts at 45oC and a volume of 7L and ends at a volume of 30L? V1/T1=V2/T2 7/ = 30/T2 T2= K 69. What would be the new temperature of a gas that begins at 56oC with a pressure of 90 kpa and changes to a pressure of 207 kpa? P1/T1=P2/T2 90/ = 207/T2 T2=757 K 70. What would be the volume of He if it begins at 4oC at of volume of 8L and the temperature then increases by 7oC. V1/T1=V2/T2 8/ = V2/ V2=8.09 L 71. What would be the pressure of F2 gas if the initial pressure was 67 kpa and the volume increased from 4L to 9L. P1V1=P2V2 (67)(4) = P2(9) P2=29.8 kpa 72. What is the initial volume of Cl2 if it ends with a volume of 6L? The initial pressure is 145 kpa with a temperature of 70oC and ends with a pressure of 200 kpa with a temperature of 45oC. P1V1/T1=P2V2/T2 (145)V1/ = (200)(6)/ V1=8.93 L V=16.6 L

11 73. A sample of gas has a volume of 150mL at 260K and 92.3 kpa. What will the new volume be at 376K and 123 kpa? P1V1/T1=P2V2/T2 (92.3)(0.150)/260 = (123)(V2)/376 V2=163mL 74. In a large syringe, 48 ml of ammonia gas at STP is compressed to 24 ml and 110 kpa. What must the new temperature of the gas be? P1V1/T1=P2V2/T2 (101.3)(0.048)/(273) = (110)(0.024)/T2 T2 = K 75. What is the temperature of Ar if there is 3.6 g, 560 kpa, and 7.4L? Remember R=8.31 kpa L/mol K n=m/m n=3.6/39.95 n= mol PV=nRT (560)(7.4)=(0.0901)(8.314)T T=5532 K 76. A balloon contains 2.0 L of He gas at STP. How many moles of He are present? PV = nrt n = PV/RT n = (101.3)(2.0)/(8.314)(273) n= V = n R T/ P V = L 77. A student collected 245 ml of unknown gas X over water at an atmospheric pressure of kpa and a temperature of 20 oc. If the mass of the gas is g, determine the molar mass of gas X. Vapour pressure of water at 20 oc = Patm = kpa kpa = kpa is pressure of dry gas Using PV= nrt, n = (106.26Kpa)(0.245L)/8.314 (293K) = mol STIOCHIOMETRY REVIEW PROBLEMS 78.Use the following equation for all three problems: 2 H2 (g) + O2 (g) 2 H2O (g) a) How many moles of water are produced when 5.00 moles of oxygen are consumed? b) If 3.00 moles of water are produced, how many grams of oxygen must be consumed? c) How many moles of hydrogen gas mush be used, given that data in problem b)? d) What mass of water is produced when 105 g of hydrogen gas reacts with excess oxygen? a) 10.0 mol H 2O, b) 48.0 g O 2 c) 3.00 mol H 2 and d) 937 g of H 2O 79. What mass of sodium carbonate must be used to produce L of carbon dioxide at 24 oc and 103 kpa according to the following neutralization reaction. HCl(aq) + Na2CO3 (s) --. NaCl(aq) + CO2 (g) + H2O (l) 45.8 g 80. Given: C2H5OH + 3O2 2CO2 + 3H2O a) What mass of oxygen is required to react with 1200 g of ethanol? 2500g o2 b) If 655 g of water is produced, what mass of ethanol is burned? 558 g ethanol 81. Sodium Hydroxide, when mixed with hydrochloric acid, will produce sodium chloride and water. a) Write the balanced equation for the reaction. b) What mass of sodium hydroxide must be used to produce 225 g of Sodium Chloride? NaOH + HCl NaCl + H2O =154 g of NaOH

12 82. The sulphur dioxide produced in this reaction, and in other reactions similar to it, is responsible for much of the acid rain that falls on North America. What mass of iron(iii) oxide can be obtained by the roasting of 774 g of the sulphide ore? What mass of sulphur dioxide is produced? 4 FeS + 7 O2 2 Fe2O3 + 4 SO 2 =703 g of Fe2O3 and =564 g of SO2 83. One of the reactions used in the smelting of copper ores to produce copper and sulfur dioxide involves reacting copper (I) oxide with copper (I) sulfide. When 250 kg of copper (I) oxide is heated with 125 kg of copper (I) sulfide, 285 kg of copper is recovered. a) write the balanced chemical equation for this reaction. 2 Cu2O + Cu2S 6 Cu + SO2 b) Determine the limiting reagent. Cu2S is limiting c) Calculate the theoretical yield using stoichiometry 299 kg d) Determine the percentage yield of copper. 95.3% 84. Solid carbon dioxide (dry ice) may be used for refrigeration. Some of this carbon dioxide is obtained as a by- product when hydrogen is produced from methane in the following reaction. CH4 + 2H2O CO2 + 4H2 a) What mass of carbon dioxide should be obtained from the complete reaction of 1250 g of methane? =3430 g b) If the actual yield obtained is 3000 g, what is the percentage yield? =87.5% 85. Ammonium nitrate is an important compound used both as a fertilizer and as an explosive. It is produced by reacting ammonia with concentrated nitric acid. NH3 + HNO3 NH4NO3 a) What mass of ammonium nitrate can theoretically be produced from the reaction of g of ammonia with excess nitric acid? =1762 g b) If the percentage yield is 88.5%, what mass of ammonium nitrate is actually obtained? =1560 g 86. During the formation of 35 g of potassium chloride, what volume of oxygen gas was produced when measured at 10oC and 100 kpa? 2 KClO3 2KCl + 3O2 17 L 87. If 129 g of oxygen gas and 300 g of propane are mixed and allowed to react as shown below, determine the volume of water vapour formed at 116 kpa and 120 oc? C3H8 (g) + O2 (g) CO2 (g) + H2O (l) 90.7 L 88. How many moles of the following are required to manufacture 5.0 mol of ammonia? N2 + 3H2 2NH3 a) Nitrogen = 2.5 mol required and b) Hydrogen = 7.5 mol required 89. C3H8 + 5O2 3CO2 + 4H2O How many moles of oxygen are required to react with: a) 3.0 mol of propane: 15 mol b) 20.0 mol of propane: 100 mol c) 0.2 mol of propane 1 mol

13 90. What volume of M sodium hydroxide are needed to neutralize 100 ml of 7.50 M hydrobromic acid solution? = 1.44 L 91. If I have 4.l5 g of Iron Sulfide, how many grams of Iron (III) Oxide will be made? 4FeS + 5O2 2 Fe2O3 + 4SO2 =3.77 g 92. If I have 3.8 g of oxygen, how much water will be made? C2H5OH + 3O2 2CO2 + 3 H2O =2.1 g 93. If I have 4.5 g of sodium and 6.4 grams of chlorine, which is the limiting reactant and which is in excess? 2Na + Cl2 2NaCl Cl2 is the limiting reactant and Na is in excess 94. If I have 3.2 g of aluminum and 5.4 g of bromine, which is the limiting reactant and which is in excess? Br2 is the limiting reactant and Al is in excess 95. What mass of carbon dioxide should be obtained from the complete reaction of 1250 g of methane and 2000 g of water? CH4 + 2H2O CO2 + 4H2 =2440 g of CO2 96. If I have 4.3 g of N2H4 and 6.8 g of O2, what will the percent yield be if in a Reaction, I get 4.9 g of HNO3 produced? 2N2H4 + 7O2 4HNO3 + 2H2O =7.7 g of HNO2 and %yield = 64% 97. Consider the following reaction: C3H8 + 5O2 3CO2 + 4H2O a) If 1.5 L of propane gas are burned in a barbecue what volume of carbon dioxide is produced and what volume of oxygen is consumed? (Assume all gases are at STP) 7.5 L of oxygen and 4.5 L of carbon dioxide b) If 35 g of propane gas is burned in a barbecue what volume of water vapor is produced assuming SATP? 78.8 L of oxygen c) If 35 g of propane gas is burned in a barbecue what volume of oxygen gas is consumed at SATP? 97.9 L of oxygen 98. Excess Pb(II) reacts with 25 ml of 1.5 M hydrofluoric acid to produce hydrogen gas at 22 O C and 88.5 kpa. How many liters of dry hydrogen gas are collected? Pb + 2 HCl PbCl2 + H2 = 0.52 L of H ml of 0.20 mol/l sodium carbonate solution and ml of 0.10 mol/l calcium nitrate solution are mixed together. Calculate the mass of precipitate that will form, and the concentrated of each of the ions remaining in solution. Na2CO ml ml 0.20 M 0.10 M Ca(NO3)2 CaCO3 (s) + 2 NaNO3 = 2.0 g CaCO3 will form and [Na + ] and [NO 3 - ] remaining in solution = 0.10M g of copper metal is reacted with 2.50 L of 3.00 mol/l nitric acid solution. Calculate how much of the copper metal remains after the reaction is complete. 3 Cu + 8 HNO3 4 H2O + 3 Cu(NO3)2 + 2 NO 500 g 2.50 L 3.00 M = 320 g of Cu Remains