Answer Sheet for Sample Problems for Chemistry Level 1 Final Exam 2016 Study Guide Electrons in Atoms Chapter 5 1. What is the frequency of green light, which has a wavelength of 4.90 x 10-7 m? 8 c 3.00x10 m / s c= = 4.90 x 10-7 m, c = 3.00 x 10 8 m/s =? Hz 7 4.90x10 m 6.12 x 10 14 Hz 2. An x-ray has a wavelength of 1.15 x 10-10 m. What is its frequency? c c= = 1.15 x 10-10 m, c = 3.00 x 10 8 m/s =? Hz 2.61 x 10 18 Hz 8 3.00x10 / m s 10 1.15x10 m 3. Which type of electromagnetic radiation has the most energy: Microwaves, X-Rays, or Visible light? Why? Ans: X-Rays, b/c they have the shortest wavelength and highest frequency of the types listed and frequency and energy are directly related. 4. What is the energy of a 9.50 x 10 13 Hz radiation? E=h = 9.50 x 10 18 Hz, h = 6.626 x 10-34 J. s =? J E = h = (6.626 x 10-34 J. s )( 9.50 x 10 18 s -1 ) = 6.29 x 10-15 J 5. The element with the electron configuration of [Kr]5s 2 4d 10 5p 1 is indium (In) 6. Barium (Ba) 7. Write out the electron configurations for the following elements: (For the exam make sure you know how to do both the complete and the noble gas short cut for electron configurations) a. Tc Ans: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 5 or [Kr]5s 2 4d 5 b. Aluminum Ans: 1s 2 2s 2 2p 6 3s 2 3p 1 or [Ne] 3s 2 3p 1 8. Why are Cr and Cu exceptions to writing the electron configurations of elements? They are exceptions because instead of following the pattern of filling up the d orbitals Cr actually has a half filled s and d orbitals and Cu has a half filled s and completely filled d orbitals. The reason why this occurs is because the electrons in the d-sublevel are more stable with this electron configuration than the expected configuration. 9. a. Ca b. As c. Sn.... a. 2 Ca. b. 5.As. c. 4. Sn... 9. Using the following quantum numbers, what is the maximum number of electrons that the atom could have? n=3, l = 1 ml = -1, 0, 1 The answers is 6 because the numbers stand for 3p and the three orbitals
Periodic Table and Periodic Trends Chapter 6 1. Compare & contrast how Mendeleev s periodic table was organized to how Moseley organized it. Ans: Mendeleev used atomic mass & properties of elements to order the elements. This resulted in some elements being out of order and not exactly fitting in with the others in a group. Moseley used the atomic number of each element as well as the properties of the elements and this fixed Mendeleev s table. 2. What is the periodic law? Ans: When elements are arranged by increasing atomic number, there is a repeating pattern to their chemical and physical properties. 3. Identify each of the following as a metal, nonmetal, or metalloid and as representative elements or transition elements: a. Neon b. Arsenic c. Zinc d. Magnesium Ans: a. Nonmetal, b. metalloid, c. metal, d. metal, representative representative transition representative 4. Ans: a. Alkali metals = 4. group 1 b. Noble gases = 1. group 18 c. Alkaline earth metals= 3. group 2 d. Halogens = 2. group 17 5. Rank the following elements by increasing atomic radius: carbon, aluminum, oxygen, potassium O, C, Al, K 6. Why does fluorine have a higher ionization energy than iodine? The e- in F are closer to the nucleus so they are more difficult to remove. 7. Why do elements in the same family generally have similar properties? Same number of valence e- Chapters 8 & 9 Chemical Names and Formulas Practice Problems: 1. Write the symbol and name for the ion formed when: a. An aluminum atom becomes an ion Al 3+ Aluminum ion b. A fluorine atom becomes an ion F - Fluoride ion c. A nitrogen atom gains three electrons N 3- Nitride ion 2. Write the formula (including charge) for each polyatomic ion: a. Hydroxide OH - b. Nitrate NO 3 - c. Sulfate SO 4 3. Write the formulas for the following compounds they are a mixture of ionic and molecular compounds: a. Potassium sulfate K 2 SO 4 d. Magnesium hydroxide Mg(OH) 2 b. Carbon tetrabromide CBr 4 e. Diphosphorus trioxide P 2 O 3 c. Silver chloride AgCl d. Iron(II) oxide FeO
4. Name the following compounds they are a mixture of ionic and molecular compounds: a. CS 2 Carbon disulfide d. Cu 2 O Copper(I) oxide b. NaI Sodium iodide e. K 2 S Potassium sulfide c. NH 4 Br Ammonium bromide f. N 2 O 5 Dinitrogen pentoxide Chapter 9 Covalent Bonding Sample Problems: 1. Draw electron dot structure and state the shape of each of the following: a. PCl 3 trigonal pyramidal c. BF 3 trigonal planar b. CO 2 linear d. CO 3 trigonal planar but remember it has resonance 2. State the type of bonds (nonpolar covalent, polar covalent, or ionic) found in the following compounds, and circle the polar molecules. a. CO polar covalent c. SF 6 polar covalent b. NH 3 polar covalent d. CaO ionic Chapter 10 Chemical Reactions Sample problems: Balance the following & state what type of reaction it is: 1. 2 AgNO 3 + H 2 S Ag 2 S + 2 HNO 3 Double Replacement 2. 2 HCOOH + O 2 2 CO 2 + 2 H 2 O Complete Combustion 3. ALREADY BALANCED Single Replacement 4. 2 SO 2 + O 2 2 SO 3 Synthesis
5. Predict the products in the following chemical reactions; make sure to balance the reactions. a. Zn (s) + 2 HCl (aq) H 2 (g) + ZnCl 2 (aq) b. H 2 SO 4 (aq) + 2 KOH (aq) 2 H 2 O (l) + K 2 SO 4 (aq) 6. Write a balanced net ionic equation for the following reaction and state what the spectator ions are: Pb(NO 3 ) 2 (aq) + H 2 SO 4 (aq) PbSO 4 (s) + HNO 3 (aq) Complete ionic equation: Pb 2+ + 2NO 3 - + 2H + + SO 4 PbSO 4 (s) + 2H + + 2NO 3 - Net ionic equation: Pb 2+ + SO 4 PbSO 4 (s) Spectator ions: H + and NO 3 - Identify what is oxidized (O), what is reduced (R), and the oxidation numbers on all the elements. 0 0 +2-2 1. 2Zn + O 2 2ZnO zinc was oxidized, oxygen was reduced Not Covered in 2016 0 0 +2-1 2. Mg + I 2 MgI 2 magnesium was oxidized, iodine was reduced +1-2 0 0 +1-1 3. H 2 S + Cl 2 S + 2HCl sulfur was oxidized, chlorine was reduced Chapter 13 States of Matter Sample Problems: 1. At 25 C, the vapor pressure of compound A is 40 mmhg and the vapor pressure of compound B is 50 mmhg. Which compound will evaporate faster? Which compound has the higher boiling point? Explain. Compound B will evaporate faster because there is more vapor above the solution to begin with so it takes less energy to make the compound turn into a gas. Compound A has the higher boiling point because it has the lower vapor pressure so, less of the compound is in the gaseous state so it will take more energy to cause A to boil therefore you need a higher temperature. 2. Name two changes of state (e.g. freezing, melting, etc.) that require energy. Melting, boiling (vaporizing), sublimation
3. State the three kinds of intermolecular attractive forces. Rank these forces from weakest to strongest (for molecules of about the same size). Dispersion, dipole-dipole, hydrogen bond Not Covered in 2016 4. Which compound would you expect to have the higher boiling point, CH 4 or NH 3? Explain. NH 3 because these molecules can hydrogen bond so it takes more energy (higher temperature) to break those IMF and cause ammonia to boil. Chapter 14 Gases and Section 13.1 (14.4 covered with Stoichiometry) Sample Problems: 1. Which gas diffuses faster, nitrogen or carbon monoxide? Why? Neither because they have the same molar mass. 2. A gas at 155 kpa and 25 o C occupies a container with an initial volume of 1.00 L. By changing the volume, the pressure of the gas increases to 605 kpa as the temperature is raised to 125 o C. What is the new volume? P1V1 T1 = P2V2 T2 V2 = (155kPa)(1L)(398K) (298K)(605kPa) = 0. 342L 3. A sample of air has a volume of 6.00 L at 100 kpa. What volume will it occupy at 25.0 kpa if the temperature remains constant at 30 o C? P1V1 T1 = P2V2 T2 V2 = (100kPa)(6L)(303K) (303K)(25.0kPa) = 24 L 4. A steel cylinder that has a volume of 20.0 L is filled with nitrogen gas to a final pressure of 25.0 atm at 28 o C. How many moles of nitrogen gas does the cylinder contain? PV = nrt n = (25 atm)(20.0l) (0.0821 Latm molk )(301K) = 20. 2 mol 5. A gas mixture containing oxygen, nitrogen, and carbon dioxide has a total pressure of 32.9 kpa. If the pressure of oxygen is 6.6 kpa and the pressure of nitrogen is 23.0 kpa, what is the pressure of CO 2 in the mixture? P T = P O2 + P N2 + P CO2 P = 32.9 kpa (6.6 kpa + 23.0 kpa) = 3.3 kpa Chapter 15 Solutions Sample Problems: 1. 36.0 g of glucose are dissolved to make 2.00 liters of solution. What is the molarity of the solution? (The molar mass of glucose is 180 g/mol). M = mol L 36.0g( 1 mol 180 g ) = 0. 20 mol 0.20 mol M = = 0. 1M 2.00 L 2. How many grams of solute are in 335 ml of 0.425 M KNO 3? M = mol L mol mol = M L = (0. 425 ) (0. 335 L) = 0. 1424 mol (101.1g ) = 14. 4 g KNO L 1 mol 3
3. If 250 ml of 2.0 M CaCl 2 is diluted to a final volume of 700 ml, what is the concentration of the diluted solution? M1V1 = M2V2 M2 = M1V1 V2 (2. 0M)(250 ml) = = 0. 714 M 700mL 4. Calculate the percent by mass of 3.55 g of NaCl dissolved in 88 g water. 3. 55g x100 = 3. 88% (3. 55g + 88g) 5. What mass of water must be added to 255.0g KCl to make a 15.00 percent by mass aqueous solution? 255.0g 216.75g x100 = 15% 255g = 38.25g + 0.15x (255.0g + x) Chapter 19 Acids and Bases Sample Problems: 0.15 = 0.15x 0.15 x = 1445g 1. An aqueous solution tastes bitter and turns litmus blue. Is the solution acidic or basic? BASIC 2. Classify the following as an Arrhenius acid or an Arrhenius base: a. H 2 S ACID c. Mg(OH) 2 BASE b. RbOH BASE d. H 3 PO 4 ACID 3. Given the concentration of either hydrogen ion or hydroxide ion, calculate the ph, poh and concentration of the other ion at 298 K: a. [H+] = 1.0 x 10-4 M b. [OH-] = 1.3 x 10-2 M ph = -log(1x10-4 ) = 4 poh = -log (1.3x10-2 ) = 1.89 poh = 14 4 = 10 ph = 14-1.89 = 12.11 [OH-] = 1 x 10-10 M [H+] = 10-12.11 = 7.76x10-13 M 4. Calculate [H+] and [OH-] in each of the following solutions at 298 K. a. ph = 3.00 b. ph = 5.24 [H+] = 1 x 10-3 M [OH-] = 1 x 10-11 M [H+] = 10-5.24 = 5.75 x 10-6 M [OH-] = 1x10 14 5.57 x 10 6 = 1. 74 x 10 9 M 5. How many milliliters of 0.225 M HCl would be required to titrate 6.00 g KOH? 6.00g KOH ( 1 mol KOH 56.1 g HCl + KOH H 2 O + KCl 1 mol HCl ) ( ) = 0.107mol HCl ( 1 L ) = 0.475 L = 475 ml 1 mol KOH 0.225 mol HCl