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A.P. Chemistry-Semester 1 Final NET IONICS – TRY THESE • • Part A -20 scantron questions-no calculator – Can use periodic table and pink sheet • Part B- Short Essay – no calculator – Can use periodic table and pink sheet • Part C - can use calculator Write the net ionic equation for the reactants and the products for the laboratory situations described below. In all cases, a reaction occurs. Assume that solutions are aqueous unless otherwise indicated. Omit formulas for any ions or molecules that are unchanged by the reaction (net ionic). You must balance the equations. 1) Example: A strip of magnesium is added to a solution of silver nitrate. Mg + 2Ag+ ---> Mg 2+ +2 Ag (a) Solid calcium carbonate is strongly heated. CaCO3 ---> CaO + CO2 – Can use periodic table and pink sheet (b) A piece of nickel metal is immersed in a solution of copper(II) sulfate. Ni + Cu2+ ---> Ni2+ + Cu hydrated ions acceptable with correct charge 1 point for Ni(OH)2 as product (c) Equal volumes of equimolar solutions of sodium hydrogen phosphate and hydrochloric acid are mixed. HPO42¯ + H+ ---> H2PO4¯ incorrect charge on H2PO4¯ when only one product occurs, 1 point only 1 product point for transfer if H+ from an ionic reactant to product when a phosphate species is incorrectly but consistently written. (d) Chlorine gas is bubbled into a solution of sodium bromide. Cl2 + 2 Br- ---> 2 Cl¯ + Br2 no credit for monatomic Cl as reactant or Br as product (g) Drops of liquid dinitrogen trioxide are added to distilled water. N2O3 + H2O ---> 2 HNO2 1 product point for H+ + NO2¯ (h) Solutions of potassium permanganate and sodium oxalate are mixed in an acidic solution. 16 H+ +2 MnO4¯ + 5C2O4 2¯ ---> 2Mn2+ + 10CO2 + 8 H2O (e) Ammonia gas is bubbled into a solution of ethanoic (acetic) acid. NH3 + HC2H3O2 ---> C2H3O2¯ + NH4+ 1 product point for NH4C2H3O2 1 point for NH3 + H+ ---> NH4 + (f) Solid ammonium carbonate is added to a saturated solution of barium hydroxide. (NH4)2CO3+ Ba 2++ 2 OH¯ ---> 2 NH3 +BaCO3 +2H2O 1 product point for either NH3 or BaCO3 2 product points for all three species correct *Note about barium hydroxide = marginally soluble-mostly insoluble unless small amounts being formed or in hot water or in an acidic environment in this case given as a solution. 1. During the gravimetric dehydration of a hydrate what factors could cause you to have too large of a percentage of water driven off? Too little? This comes from the experiment we did about dehydrating a hydrate. The percentage of water expected to be driven off comes from the formula itself. Some factors that would cause you to have too much are: spattering some of the original sample or excessive heating that causes it to decompose or not waiting long enough for crucible to cool. Too little would be caused by incomplete heating. 2. If you assigned to only use one solvent to separate Ca2+ from Na+ and K+, which acid would you use? Hydrochloric, Sulfuric or Nitric? You need to choose the one that only precipitates Ca2+. Sulfuric acid is the right choice. Know those rules! 1 3. Given a list of gases, what criteria can you use to determine which one deviates the most from ideal gas behavior? At very large pressures, several factors become more important…Larger atoms deviate more, so do those atoms or ions or molecules with very strong intermolecular forces. 4. 6. 100. mL of 0.100 M Al(NO3)3 is added to 50.0 mL of 0.35 M Ba(NO3)2. What is the molarity of the NO31- ion in the resulting solution? 0.100L 0.100 mol Al(NO3)3 3 NO3- ions = 0.0300 mol NO3- What volume of water must be added to dilute 140. mL of a 1.00 M NaOH solution to 0.25 M solution? M1V1 = M2V2 (140. mL)(1.00M) = (.25M)V2 V2 = 560. mL Remember in solutions volume is additive so (560.-140.) = amount of water added = 420 mL 5.Equal number of moles of different gases are placed into a container with a small hole, what criteria determines which effuses faster? How could you determine their partial pressures when the system reaches equilibrium? 1.L 1 mol Al(NO3)3 0.0500 L 0.35 mol Ba(NO3)2 2 NO3- ions = 0.0350 mol NO31L 1 mol Ba(NO3)2 (0.0300 mol + 0.0350 mol) NO3- = 0.433 M NO3(.100 L + .0500 L) According to Graham’s Law, the smaller the molar mass the faster it effuses. Each gases partial pressure is directly related to its mass/mole ratios. 7. Complete the trends, using arrows… Period Group Atomic Radius First Ionization Energy Electronegativity 9. What is the empirical formula for a compound that is 68.4 % Cr and the rest is Oxygen? 68.4 g Cr 1 mol = 1.32 mol Cr 31.6 g O 1 mol =1.98 mol O 52.0 g 16.0 g Activity of Metals 1.32 mol Cr/1.32 = 1 8. What happens to the ionization energy of calcium when you remove the first electron, then the second electron and then the third electron? Because calcium would like to lose 2 electrons to be stable we would expect the huge jump to be between the second and the third ionization energies. 4 5 2 3 4 6 12. If you know the molality of a solution, what would you need to know to determine its mole fraction? Molarity? Percent by mass? Molality = mol solute / kg solvent You would need to Mol fraction = mol solute / total mol know the formula for Molarity = mol solute / L solution % mass = g solute / total grams both the solute and solvent, and the density of the solvent. = 1.5 Empirical Formula = Cr2O3 10. Rank the following molecules in order from least to greatest dipole moment; NO2, H2, CH4, CHCl3. Least = H2, CH4, NO2, CHCl3= greatest 13. 3 1 1.98 mol O/1.32 Give reasons to explain the reason that CaS has a higher melting point than LiF. Both are ionic, so the greater the charge, the greater the IM forces. Since Ca2+ and S2- their IM forces are greater resulting in a higher melting point. 14. What types of hybridization of the carbon atom are found in the compound CH3CHCCH2? Number the carbons 1-4 from left to right and draw the molecule… 1. C in CH3 = sp3 2. C in CH= sp2 3. C = sp 4. C in CH2= sp2 2 15. What mass of Calcium Chloride can be made from the reaction of 17.9 g of Sodium Chloride with 43.0 g of Calcium Nitrate? First write a balanced equation… 1 + ___Ca(NO 3)2 2___NaCl 1 2 ___CaCl 2 + ___NaNO3 Next Find the limiting reagent… 17.9 g NaCl 1 mol = .306 mol NaCl 16. 0.50 L of a 0.75 M aluminum sulfate solution is added to 0.10 L of a 1.20 M potassium sulfate solution. How many moles of barium chloride are necessary to completely precipitate all the sulfate ions? 2 Al3+ + 3 SO42- Al2(SO4)3 (0.50 L x 0.75 mol/L) = 0.375 mol Al2(SO4)3 = 1.13 mol SO42.306 mol NaCl 1 mol Ca(NO3)2 58.5 g 43.0 g Ca(NO3)2 1 mol = .262 mol Ca(NO3)2 2 mol NaCl = .153 mol Ca(NO3)2 K2SO4 2 K+ + SO42- (0.10 L x 1.20 mol/L) = 0.120 mol K2SO4 = 0.120 mol SO42- Total moles of sulfate = (1.13 + 0.120 mol) = 1.25 mol NaCl = L.R. 164 g .306 mol NaCl 1 mol CaCl2 111 g CaCl2 = 17.0 g CaCl2 Ba2+ + SO42BaSO4 which means a 1:1 ratio and therefore you will need 1.25 mol of BaCl2 2 mol NaCl 1 mol CaCl2 17.Give the set of quantum numbers for the last electron in Ag and As. Ag = 4 d 9 n=4 As = 4 p 3 n=4 l=2 l=1 ml = +1 ms =± ½ m1 = +1 ms = + or – ½ 18. What is the molecular mass of a substance where 1.33 g of the gas at 120oC and 780. mm Hg has a volume of 950. mL? 19. A 5.00 L flask contains 4.00 mol of ammonia gas and 6.00 mol of hydrogen chloride gas. There is also enough helium to have a partial pressure of 1.50 atm. The temperature in the flask is 125oC. a. What is the total pressure in the flask? nHe = (1.50 atm)(5.00L)= 0.230 mol= 1.50 atm (0.0821) (398 K) n = PV/RT = (1.03 atm) (0.950 L) = 0.0303 mol Pt = (.230 mol + 4.00 mol + 6.00 mol) 1.50 atm = 66.7 atm .230 mol (0.0821)( 393 K) 1.33 g / 0.0303 mol = 43.9 g/mol b. Calculate the density in g/L of the gas mixture in the flask 0.230 mol 4.00 g = 0.920 g He c. What is the mole fraction of each remaining substance if 4.00 moles of gaseous ammonium chloride is made? Write the balanced equation (no Helium reacts Duh!) 1 mol 4.00 mol 17.0 g . = 68.0 g NH3 1 mol 6.00 mol 36.5 g = 219 g HCl 1 mol (0.920 + 68.0 + 219 g ) g/ 5.00 L = 57.6 g/L NH3(g) +HCl (g) NH4Cl(g) Which is the limiting reactant? 4 mol + 6 mol 4 mol?? NH3 = L.R. Now calculate the moles present AFTER the reaction for Each Species!!! Total moles after = 6.23 mol Mol NH3 = 0 mol Mol HCl = ( 6.00 – 4.00) mol = 2.00 mol Mol NH4Cl = 4.00 mol Mol He = 0.230 mol XHCl = 2.00/6.23 =.321 XNH4Cl = 4.00/6.23 = .642 XHe = 0.230/6.23 = .0369 3 20. A compound contains only C, H and N. It is 58.51% C and 7.37% H by mass. Helium effuses through a porous frit 3.20 times as fast as the compound does. Determine the empirical and molecular formulas of this compound. 4.88 mol C/ 2.44 = 2 7.30 mol H / 2.44 = 3 Empirical formula = C2H3N efm = 41.03 g/mol 3.20 = √X 1 = √4.00 21. Draw Lewis dot diagrams for the following species: carbonate ion, carbon dioxide and carbon monoxide. CO2 CO32- CO 2.44 mol N/2.44 =1 X = 40.96 g/mol Molecular formula = C2H3N b. Describe the molecular shape expected for each species. Include any resonance structures. CO32- will be trigonal planar with 3 resonance structures CO2 is linear CO is linear c. Compare the bond strength and bond length of the carbon-oxygen bonds in the three species. Carbon monoxide would have the strongest bond with the shortest length due to the fact that it is a triple bond. Carbon dioxide would be next, with a double bond. The carbonate ion has resonance which means each bond is like a 1 1/3 strength and a length shorter than a single bond. 22. a. Draw the complete Lewis dot structures for the molecules CF4 and SF4. 23. Oxygen is found in the atmosphere as diatomic gas, O2, and as ozone, O3. a. Draw the Lewis structures for both molecules. b. b. In terms of molecular geometry, account for the fact that the CF4 molecule is nonpolar, whereas the SF4 molecule is polar. CF4 is symmetrical with equal pull in all directions, and this is nonpolar. SF4 has a lone pair on the central atom and is not symmetrical, therefore it is a polar molecule. Use the principles of bonding and molecular structure to account for the fact that ozone has a higher boiling point than diatomic oxygen. Ozone is nonpolar, but does have a lone pair on the central atom, this creates a small IM force like a dipole, whereas oxygen is symmetrical and has less IM forces. Ozone has a greater molar mass. c. Use the principles of bonding and molecular structure to account for the fact that ozone is more soluble than diatomic oxygen in water. d. Explain why the two bonds in O3 are the same length and are longer than the bond length of the bond in diatomic oxygen. Same as letter “b” Ozone has resonance which means each bond is like a 1 ½ strength and a length shorter than a single bond but longer than a double bond. 24. When NH3 gas is introduced at one end of a long tube, while HCl gas is introduced simultaneously at the other end, a ring of white ammonium chloride is observed to form in the tube after a few minutes. This ring is closer to the HCl end of the tube than the NH3 end. The molecules of gas are in constant motion so the HCl and NH3 diffuse along the tube. Where they meet, NH4Cl(s) is formed. Since HCl has a higher molar mass, its velocity (average) is lower, therefore, it doesn’t diffuse as fast as the NH3. 25. Explain each of the following using your knowledge of intermolecular forces and molecular structure. a. Br2 has a higher boiling point than Cl2. Both have a dipole moment of zero, but bromine has a larger molar mass. b. F2 has a greater bond length than O2. Oxygen has a double bond and fluorine has a single bond. Single bonds are longer than double bonds. c. LiF has a higher melting point than NaCl. The Coulomb’s force in LIF is larger due to the smaller atom size, therefore a higher melting point. d. C4H10 has a higher boiling point than CH4. Both are nonpolar, CH4 is symmetrical and has a smaller molar mass. Butane has greater IM forces due to its larger mass. 4