Chem 111 Test out & Study guide

STUDY GUIDE AND SAMPLE EXAMINATION FOR CHEMISTRY 11100 This Study Guide describes topics to be mastered prior to attempting the examination to establish credit in Purdue's Chemistry 11100. The material can be studied from many of the textbooks on the market. A list of several of these is given below. IMPORTANT: 1. Read this material thoroughly if you contemplate trying for advanced placement (and extra credit which counts toward graduation). 2. Study all the material listed in the outline. 3. Work many practice problems. 4. When you feel prepared for it, take the sample examination. 5. Come to the actual examination rested and confident. SUBJECT MATTER - A Brief Outline The subject matter of any General Chemistry course is varied. Chemistry 11100 deals with metrication, scientific units and measurements, the classification of matter, chemical and physical changes, chemical formulas and equations, energy in chemical reactions, characteristics of reactions, atomic weights, stoichiometry, atomic structure from both historical and modern aspects, periodic relationships among the elements, chemical bonding, solubility and solutions, chemical nomenclature, and oxidation-reduction. Interspersed in the preceding topics is descriptive chemistry pertinent to the various topics, e.g., hydrogen and oxygen are studied as examples of real gases and water as a common liquid. The course is taken primarily by students in the College of Agriculture and the College of Health and Human Sciences and is designed to meet their needs. It is also taken by smaller numbers of students in the College of Liberal Arts and the College of Technology. A knowledge of high-school chemistry is presumed, and the course goes far beyond the high school level. A weekly three-hour laboratory is an integral part of the course and includes measurement techniques, the use of the analytical balance, titration techniques and experiments with various elements and compounds. The topics presented in the outline should be studied prior to attempting the sample examination included with this study guide. These topics are broken down under several headings according to those found in the texts currently in use at Purdue. In preparing for the examination it is important to work many problems. However, the problems may test your recall of relations like D = mass/volume or M = moles/Liter, and also should be used to measure your understanding of the concepts and principles involved. At the end of this Study Guide you will find a sample examination over this material. Allow one hour for the exam. Naturally, it does not cover every point. No examination extending over a reasonable time period can do that. SPECIAL NOTE: A word of advice concerning the taking of the actual examination for credit. No one does well on an examination when he/she is excessively fatigued. Therefore, you are urged to provide yourself an adequate rest period before taking the actual examination. If your trip to the campus necessitates travel into the late hours of the night or an extremely early departure from your home, you may be well advised to allow for a one night's rest in the Lafayette area before taking the examination. Many students who are unsuccessful with the examination tell us that failing to take the above precautions contributed strongly to this result. Most such students find that their first year was somewhat less rewarding than it might have been because of the time spent retracing materials studied in high school. Please consult your advanced credit schedule for the actual time and place of the examination. It is usually given both morning and afternoon. The texts referred to below may be ordered from any of the local book stores: Local Book Stores: University Book Store, 360 State Street, W. Lafayette, IN Village Book Cellar, 308 W. State Street, W. Lafayette, IN 47906. 47906. Follett's Purdue Book Store, Purdue Service Center, W. Lafayette, IN 47906. Texts: Introduction to Chemistry, 2nd ed (2010) by Bauer/Birk/Marks; McGraw-Hill. General Chemistry, the Essential Concept by Chang; McGraw-Hill. Introductory Chemistry by Tro; Pearson/Prentice-Hall. 2 STUDY GUIDE FOR CHEMISTRY 11100 1. Fundamentals. • • • The metric system. Basic SI units and prefixes, unit conversions. Significant figures - identify the number of significant figures in a number; rules for addition and subtraction, multiplication and division. 2. The mole, chemical formulas, and chemical equations. Know the relationship between a mole of material and its weight, volume, concentration, or number of atoms or molecules. Given the name of any "common" chemical compound, write the correct formula; or given the formula, write the name. Balance simple equations by inspection. Given the formula for any compound, (a) name the elements that make up the compound; (b) indicate the number of each kind of atom represented; (c) determine the mass of each element in one mole of the compound; and (d) calculate the molecular (formula) weight. Given a sample of a compound (or its mass) and its formula, calculate, (a) the number of moles of the compound in the sample; (b) the number of moles of any one element in the sample; (c) the number of molecules (formula units) in the sample; and (d) the number of atoms of any one element in the sample. 3. Chemical stoichiometry. Use Avogadro's Hypothesis and information from any two of the following to calculate the third: (a) (b) (c) the ratio between the volume of gaseous reactants and products the formula of the product the formula of the reacting elements Example: One volume of hydrogen (H 2 ) reacts with one volume of chlorine (Cl 2 ) to produce two volumes of hydrogen chloride. What is the formula of hydrogen chloride? Given a sample of some compound and a way to find the mass of each element, find the percent (by weight) of each element. Given the percent (by weight) of each element in a compound, determine the simplest (empirical) formula for the compound. Given all reactants and products of a reaction (with the correct formulas) and any one of the following pieces of information, calculate any other piece of information on the list. (a) (b) (c) (d) (e) the number of moles of one of the products or reactants the mass of any product or reactant the number of molecules of any product or reactant the number of atoms or mass of any single element involved concentration and volume of a solution of reactant or product 3 Know the definition of exothermic and endothermic and how these terms are related to chemical reactions. 4. Properties of liquids and solutions. Describe what occurs when a solid changes to a liquid or a liquid changes to a gas using a microscopic model. (a) What differences are there in the arrangement of the atoms of molecules; (b) what happens to the average speed of the particles; (c) what changes occur in the distance between particles; (d) what temperature changes occur, why? Name several physical properties of liquids which are characteristic of these substances and can therefore be used to identify an unknown. Describe how you would measure each of the properties listed above. Describe differences in solids, liquids and gases in terms of macroscopic properties such as: (a) density, (b) compressibility, (c) viscosity, (d) vapor pressure. Describe the effect that varying any of the following will have on the vapor pressure over a liquid: (a) volume of the liquid, (b) surface area of the liquid, (c) volume of gas over the liquid, (d) temperatures. Describe the various processes or "steps" that must occur in order for a given substance to dissolve. Decide whether the energy associated with that process favors dissolution. Take a chemical off of the shelf and prepare a specified volume of solution of some given concentration. (e.g., How would you prepare 200 mL of 0.1 M NaCl?) Calculate the moles (or mass) of solute in a given volume of solution of known concentration. (e.g., How many grams of NaCl are in 50 ml of a 1.5 M solution?) Use a solution of some known concentration to prepare a solution of a specified lower concentration by proper dilution. (e.g., Make 100 ml of 0.22 M HCl from a solution of 2.5 M HCl.) 5. Periodic chemistry. By referring only to the Periodic Table you should be able to give the common valence and oxidation number for any element in Group IA, IIA, IB, IIB, IIIB, IVB, VB, VIB, and VIIB. Write the formula (including the proper charge) for any of the following polyatomic ions: perchlorate, chlorate, chlorite, hypochlorite, permanganate, dichromate, chromate, cyanide, oxalate, acetate, hydroxide, sulfate, sulfite, carbonate, nitrate, nitrite, and phosphate. Identify a period, row, group, column, and family on the Periodic Table. Indicate whether or how some given property such as melting point, density, ionization energy, electronegativity, atomic weight or atomic radius is a periodic function of atomic number. Use the Periodic Table to predict properties of elements in a given family. 4 6. Electronic structure of atoms. Know the experiments which have helped us build a modern model of the atom. Given the symbol for an element or ion and a Periodic Table, specify the number of protons in the nucleus of that particle and the number of electrons around the nucleus. Given an isotope of an element, find the number of protons, neutrons and electrons used. From the position of an element in the Periodic Table, indicate the number of electrons in the outer or valence shell. Conversely, knowing the number of electrons in the valence shell, identify the chemical family (group) to which the element belongs. Name the orbitals that contain the valence electrons for the elements located in any given section of the Periodic Table. Given a Periodic Table with all of the symbols and numbers erased, name the orbitals (s, p, d, f) that contain the valence electrons for a particular element of group of elements just by looking at the position that element would occupy in the table. Write electron configurations for elements and ions. Identify the number of valence e− from the electron configuration. 7. Chemical bonding and related phenomena. Given the symbol, name or atomic number for any element, write the Lewis structure to represent that element. Write the Lewis (electron-dot) structure for any compound containing no more than three elements. Identify the shapes of molecules; use bond polarity and shape to describe the polarity of a molecule. From the chemical and physical properties of a compound (e.g., electrical conductivity, melting point, heat of fusion, etc.) predict whether the compound is ionic or covalent. If you are given a sample of some unknown compounds, describe some tests that you could perform that would help you decide whether the compound is ionic or covalent. Use the Periodic Table to predict which of two elements would have (a) the larger radius, (b) the greater electronegativity, (c) the higher ionization energy. Use the Periodic Table to predict which of two ions would have the larger radius. Use a table of electronegativities to predict whether the bond between two given elements will be ionic or covalent. 5 6 CHEMISTRY 11100 - SAMPLE EXAM This Sample Exam consists of 30 questions. Choose the one best response for each question. _______ 1. When fluorine reacts with calcium to produce CaF 2 , the bond formed is: (a) (b) (c) _______ van der Waals none of these. 3 (b) 4 caloric. spontaneous. 14 + 1.096 x 10 (c) 5 there are ___ significant figures. (d) 6 (e) 17 4. The temperature on a cold day at Purdue was given as −12.0̊F. The temperature in Celsius was: −12.0°C 0.0°C −79.2°C 64 (d) (e) 331.5°C none of these 68 5. 30 Zn and 30 Zn are examples of: (a) (b) (c) _______ (d) (e) 16 (a) (b) (c) _______ exothermic. isothermal. endothermic. 3. In the sum of 2.82 x 10 (a) _______ (d) (e) 2. A chemical reaction that absorbs heat from its surroundings is termed: (a) (b) (c) _______ ionic. covalent. metallic. isotopes. allotropes. phenemones. (d) (e) cations. anions. 6. Which one of the following combinations of names and formulas is incorrect? (a) H 3 PO 4 phosphoric acid (b) HNO 3 nitric acid (c) NaHCO 3 sodium carbonate (d) H 2 CO 3 carbonic acid (e) KOH potassium hydroxide 7 _______ 7. When the temperature of a gas is raised, the molecules must: (a) (b) (c) (d) (e) _______ 8. When NaCl dissolves in water, it is present as: (a) (b) (c) _______ decrease their distance from each other. move about more rapidly. remain stationary. turn colors. none of these. NaCl molecules. Na atoms and Cl atoms. + – Na cations and Cl anions. 9. The name for the ion, S (a) (b) (c) 2– NaOH and HCl. (NaCl) 2 dimers. is: sulfate. sulfite. sodium. _______ 10. The name for the ion, Cr 2 O 7 (a) (b) (c) (d) (e) (d) (e) 2– saccharin. sulfide. is: chromate. chromium. chromium oxidate. (d) (e) dichromate. none of these _______ 11. The orbital that is shaped like a sphere is the: (a) (b) (c) n orbital. s orbital. p orbital. (d) (e) d orbital. f orbital. _______ 12. The proper Lewis structure (electron-dot structure) for N 2 is: (a)  :: N  : :N (d) (b) (c)  : N  : N   :N:::N: (e) 8  :N ⋅ ⋅N   :N::N: _______ 13. Which is not a property of metallic solids? (a) (b) (c) (d) (e) conduct electricity may be molded under pressure usually high melting atoms pack as hard spheres They are all properties of metallic solids. _______ 14. A given liquid occupies 445 mL and weighs 540 g. The density of this liquid is (units missing): (a) 0.73 (b) 0.82 (c) 0.85 (d) 1.21 2 (e) 1.24 6 _______ 15. The ion whose outer shell electronic configuration is 2s 2p is: (a) Ne (b) Li + (c) – F (d) S 2– (e) Ca 2+ _______ 16. The elements of Group VII are known as the: (a) (b) (c) lanthanides. noble gases. alkali metals. (d) (e) ideal gases. halogens. _______ 17. Torr is a unit employed in the measurement of: (a) volume. (b) density. (c) mass. (d) pressure. (e) none of these. _______ 18. To what volume must you dilute 200.0 mL of a 6.00 M solution of NaCl to obtain a 0.30 M solution of NaCl? (a) 1.0 L (b) 2.0 L (c) 4.0 L (d) 0.5 L _______ 19. Calculate the molar mass of (NH 4 ) 3 AsO 4 . (a) (b) (c) (d) (e) 417.80 g/mol 193.03 g/mol 165.02 g/mol 156.96 g/mol None of the above are close to the correct answer. 9 (e) 7.8 L _______ 20. The empirical formula of a compound is CH 2 . If the molecular weight of the compound is 56, how many carbon atoms does a molecule contain? (a) 4 (b) 1 (c) 2 (d) 6 (e) 3 _______ 21. One mole of KClO 3 decomposes to yield: (a) (b) (c) 1 mole of oxygen atoms. 1.5 moles of O 2 . 2 moles of KCl. (d) (e) 3 moles of O 2 . none of these. _______ 22. The oxides of metals: (a) (b) (c) (d) (e) dissolve in water to form acids. dissolve in water to form bases. are salts. dissolve in water to form oxygen. both (a) and (d). _______ 23. A compound is subjected to chemical analysis and you receive the following report from the analysis laboratory. Determine the simplest formula for the compound. % P = 31.6 % O = 65.4 (a) PH 3 O 4 (d) P 2 H 3 O 2 (b) P 2 H 6 O 8 (e) PH 2 O 5 % H = 3.06 (c) PH 3 O 2 _______ 24. Calculate the number of oxygen atoms in 29.34 g of sodium sulfate, Na 2 SO 4 . (a) 1.244 x 1023 O atoms (b) 4.976 x 1023 O atoms (c) 2.915 x 1024 O atoms (d) 1.166 x 1025 O atoms (e) None of the above are close to the correct answer. 10 _______ 25. Balance the following equation: ___Ca 3 (PO 4 ) 2 (s) + ___SiO 2 (s) + ___C(s) → ___CaSiO 3 (s) + ___CO(g) + ___P 4 (s) (a) Ca 3 (PO 4 ) 2 (s) + 3 SiO 2 (s) + 8 C(s) → 3 CaSiO 3 (s) + 8 CO(g) + P 4 (s) (b) Ca 3 (PO 4 ) 2 (s) + 3 SiO 2 (s) + 14 C(s) → 3 CaSiO 3 (s) + 14 CO(g) + P 4 (s) (c) Ca 3 (PO 4 ) 2 (s) + 3 SiO 2 (s) + 8 C(s) → 3 CaSiO 3 (s) + 8 CO(g) + 2 P 4 (s) (d) 2 Ca 3 (PO 4 ) 2 (s) + 6 SiO 2 (s) + 10 C(s) → 6 CaSiO 3 (s) + 10 CO(g) + P 4 (s) _______ 26. How much iron can be made from 495 g of aluminum by the following process: Fe 2 O 3 + 2 Al (a) (b) (c) ➝ 683 g 743 g 1020 g (d) (e) Al 2 O 3 + 3 Fe 1540 g 2310 g _______ 27. Determine the empirical formula for the compound whose percentage composition is 42.1% C, 51.5% O, and 6.4% H. (a) (b) (c) CH 2 O C 2 H2 O2 CH 3 O (d) (e) C 2 H4 O2 CHO 5 _______ 28. How many grams of water will be formed when a mixture containing 1.93 g C 2 H 4 and 5.92 g O 2 is ignited in a closed vessel? (a) (b) (c) 2.22 g 5.89 g 10.21 g (d) (e 11 12.81 g 14.50 g ANSWERS TO 11100 SAMPLE EXAM 1. A 2. C 3. A 4. E 5. A 6. C 7. B 8. C 9. E 10. D 11. B 12. C 13. E 14. D 15. C 16. E 17. D 18. C 19. B 20. A 21. B 22. B 23. A 24. B 25. D 26. D 27. A 28. A 12