Unit 2. Atoms, Molecules, and Ions Upon successful completion of this unit, the students should be able to: 2.1 State and be able to apply the Law of Conservation of Mass, Law of Definite Proportions, and Law of Multiple Proportions. 1. In a combustion reaction, 46.0 g of ethanol reacts with 96.0 g of oxygen to produce water and carbon dioxide. If 54.0 g of water is produced, how much carbon dioxide is produced? 2. A sample of chloroform is found to contain 12.0 g of carbon, 106.4 g of chlorine, and 1.10 g of hydrogen. If a second sample of chloroform is found to contain 30.0 g of carbon, how many grams of chlorine and grams of hydrogen does is contain? Which law are you assuming to be true in order to solve this problem? 3. A sample of water, H 2 O, contains 2.02 g of hydrogen and 16.0 g of oxygen. A sample of hydrogen peroxide, H 2 O 2, contains 2.02 g of hydrogen and 32.0 g of oxygen. Show how this data illustrates the law of multiple proportions. 2.2 Explain the historical development of the atomic concept leading to the modern view of the atom. This would include being able to describe Dalton s atomic theory, Thomson s cathode-ray tube experiment, the plum pudding model, Milliken s oil drop experiment, and Rutherford s gold foil experiment. 1. Is the following finding, expressed to the nearest atomic mass unit, in agreement with Dalton s atomic theory? Explain. An atom of cobalt has a mass of 60 amu and one of nickel also has a mass of 60 amu. 2. Describe in detail the setup, experimental results, and conclusions from the Rutherford gold foil experiment. 3. List one of the main conclusions about the nature of the atom which came from the Rutherford experiment. 4. What are cathode rays made of? 5. What did Milliken determine from his oil drop experiment? 2.3 List the fundamental sub-atomic particles, their basic properties, and describe the basic nature of the atom. 1. Which subatomic particle (proton, neutron, or electron) has the smallest mass? 2. Which subatomic particle has about the same mass as a proton: neutron or electron? 3. Which subatomic particle has a negative charge? 4. True or False. The nucleus accounts for most of the volume of an atom?
5. For the following statements about the subatomic particles, indicate which subatomic particle the statement best describes. a. The neutral particle that is found in the nucleus. b. The particle that determines the atomic number for an element. c. The particle that is lost or gained when an atom forms an ion. 2.4 Define isotope, write (and interpret) atomic symbols of the form A Z X, and solve related problems. 1. Which of the following pairs represent isotopes [circle choice(s)]? a. 78 A 78 A b. 79 A 80 A c. 78 A 77 A d. 80 A 33 34 34 34 34 33 34 35 79 A 2. Define isotope. 3. Write the atomic symbol for a chromium atom having 30 neutrons. 4. What is the mass number of an iodine atom that has 72 neutrons? 5. What is the mass number for an isotope of silicon that has 13 neutrons? 6. Gallium-68 is used in nuclear medicine to detect pancreatic cancer. How many protons, neutrons and electrons does a neutral atom of this isotope have? 7. How many electrons and neutrons are there in a neutral atom of germanium-70? 8. Using the atomic symbols provided, determine the number of protons, neutrons and electrons in the following. Atomic Symbol Number of Protons Number of Neutrons Number of Electrons 37 Cl 59 Ni 3+ 79 Se 2-2.5 Define molecule, ion, atom, cation, anion, monatomic ion, and polyatomic ion and solve related problems. 1. Define cation. 2. Differentiate between the terms monatomic and polyatomic ion. Give an example of each. 3. Consider the following substances: O 2, CO 2, Na 2 O. a. Which is a molecule but not a compound? b. Which is a compound but not a molecule?
2.6 Describe the type of information conveyed by a molecular formula and a structural formula. 1. Describe the type of information conveyed by a molecular formula and a structural formula. 2. The structural formula for octane is shown below. Write the molecular formula of octane. H H H H H H H H H C C C C C C C C H H H H H H H H H 2.7 Correctly use terms associated with the Periodic Table, including Period, Group (Family), atomic number, atomic mass (atomic weight), representative elements (main group elements), transition metals, lanthanides, actinides, alkali metals, alkaline earth metals, halogens, and noble gases. 1. Which term best describes uranium: lanthanide, halogen, noble gas, or actinide? 2. What is the common name for group 2A of the Periodic Table? 3. Give the NAME of the noble gas that is in period 4 of the periodic table. 4. What is the common name of group 17 of the Periodic Table? 5. Using the Periodic Table provided, give a chemical symbol for an element that would belong to the following: a. alkaline earth metal b. metal in row 2 c. halogen d. nonmetal in Group 3A e. transition metal f. metalloid 2.8 Classify elements in the Periodic Table as metals, nonmetals, or metalloids and list and define the general properties of these classifications. 1. Identify hydrogen as a metal or nonmetal. 2. True or False. Helium is more reactive than fluorine. 3. Is barium a metal, nonmetal, or metalloid? 4. An element is found to be ductile in its solid state. Is it likely a metal or nonmetal? 5. Define malleable.
2.9 Identify the elements which exist as solids, those which exist as liquids, and those which exist as gases at room temperature and pressure. 1. Which metal is a liquid at room temperature and pressure? 2. Which is the only nonmetal that exists as a liquid at room temperature and pressure? 3. List all the elements which exist as gases under ordinary conditions. 2.10 Describe the meaning of periodicity as it applies to the Periodic Table, solve related problems, and state that the Periodic Table is arranged so that the elements are 1) in order of increasing atomic number, and 2) so that elements in the same column have similar properties. 1. List one of the two main reasons why the elements are arranged the way they are on the Periodic Table. 2. Estimate the density of selenium (Se) from the following densities (g/cm 3 ): S, 2.07; Te, 6.24; As, 5.72; Br, 3.12. Show how you arrived at your answer. 3. The chemical behavior of selenium would be most similar to that of: arsenic bromine sulfur krypton 2.11 Write the name if given the symbol and symbol if given the name of the following elements: elements 1-56, Pt, Au, Hg, Pb, Rn, Ra, U, and Pu. 1. Give the name and symbol of the element that has an atomic number of 23. 2. Give the names of the following elements. a. Rb b. B c. Sn d. Au e. Ag f. Cs g. Sc 3. Give the names of the following elements. a. Ti b. Ni c. Co d. Pu e. U f. As g. Ne h. V i. He j. Si k. Cl l. N m. Ru o. Zr 4. Write the symbol and name of the element in Period 5 and Group 10 of the Periodic Table.
2.12 Write the symbol (including the charge) and name for the following monatomic ions: all those of Group 1, Group 2, and Group 17; Al 3+, N 3-, P 3-, As 3-, O 2-, S 2-, Se 2-, Te 2-, Fe 2+, Fe 3+, Cu +, Cu 2+, Ag +, Zn 2+, H +, H -. 1. Give the symbol of the stable ion of the following elements: a. aluminum b. barium c. selenium 2.13 Apply the systematic nomenclature (i.e. Roman numeral system) to all possibilities of cations with multiple charges. 1. Use the systematic nomenclature to name the following ions: a. Zr 4+ b. In 3+ c. Fe 2+ 2.14 Apply the older nomenclature system to Fe 2+ (ferrous), Fe 3+ (ferric), Cu + (cuprous), Cu 2+ (cupric) 1. List the common and systematic names of the two common ions of iron. 2.15 Write the symbol (including charge) and name for the following polyatomic ions: NH 4 +, NO 2 -, NO 3 -, SO 3 2-, SO 4 2-, HSO 4 - (hydrogen sulfate and bisulfate name), OH -, CN -, PO 4 3-, HPO 4 2-, H 2 PO 4 -, CO 3 2-, HCO 3 - (hydrogen carbonate and bicarbonate name), ClO -, ClO 2 -, ClO 3 -, ClO 4 -, C 2 H 3 O 2 -, MnO 4 -, Cr 2 O 7 2-, CrO 4 2-. 1. Applying the general principles learned so far in this course, predict the correct name for the IO 4 - ion? 2. Give the formula (including charge) of the following ions: a. nitride b. nitrate c. nitrite d. carbonate e. bicarbonate f. sulfate g. bisulfate 2.16 Write the names and formulas for salts derived from the monatomic and polyatomic ions mentioned above. 1. Name the following compounds. a. BaI 2 b. Na 3 PO 4 c. NH 4 Br d. LiNO 3 e. NaI f. K 2 S g. FeCO 3 h. KMnO 4 i. CuHCO 3
2. Give the formulas for the following compounds. a. potassium bicarbonate b. silver sulfate c. aluminum cyanide d. zinc phosphide e. ammonium sulfite f. cuprous sulfide g. barium dichromate 2.17 Write the names and formulas for binary molecular compounds; students will expected to memorize and apply the following prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-. 1. Name the following compounds. a. S 2 Cl 4 b. NO 2 c. CCl 4 2. Give the formula for the following compounds. a. tetraphosphorus decaoxide b. sulfur hexafluoride c. selenium hexafluoride d. carbon monoxide 2.18 Write the names and formulas for common inorganic acids based upon the anions mentioned above. 1. Give the names of the following. a. H 2 SO 3 (aq) b. HBr c. H 2 CrO 4 d. HCN e. H 3 PO 4 2. Write the formulas of the following acids: a. hydrosulfuric acid b. nitrous acid c. nitric acid d. phosphoric acid 2.19 List the seven elements which exist as diatomic molecules. 1. List the seven elements which exist as diatomic molecules under ordinary conditions. Additional Unit 2 Sample Questions: 1. Give the chemical names for the following compounds based on the chemical formulas listed below. a. Al(ClO) 3 b. FeBr 3 c. Pb(NO 3 ) 2 d. PF 5 e. HClO 4 f. SO 3
2. Write the chemical formulas for the following compounds based on the chemical names provided. a. silver sulfate b. sodium oxide c. magnesium hydroxide d. nitrogen tribromide e. carbon dioxide f. tetraphosphorus decasulfide