* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download CHEM 400 - El Camino College
Chemical bond wikipedia , lookup
Hypervalent molecule wikipedia , lookup
Atomic nucleus wikipedia , lookup
Electron configuration wikipedia , lookup
Abundance of the chemical elements wikipedia , lookup
Size-exclusion chromatography wikipedia , lookup
Periodic table wikipedia , lookup
Gas chromatography–mass spectrometry wikipedia , lookup
Isotopic labeling wikipedia , lookup
History of chemistry wikipedia , lookup
Chemical element wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
Molecular dynamics wikipedia , lookup
Extended periodic table wikipedia , lookup
Chemistry: A Volatile History wikipedia , lookup
IUPAC nomenclature of inorganic chemistry 2005 wikipedia , lookup
CHEM 60 E X AM 1 I F O R M AT I O N F AL L 2 0 1 5 GENERAL INFORMATION Exam Time, Place, Format, and Rules The exam will take place on Wednesday, 09/30, from 7:00 pm to 9:00 pm in room SCI 130. The exam will have two parts: multiple choice questions and open-end questions. You will need to bring a Scantron form No. 882E. Partial credit will be given for open-end questions only. Open-end questions will be of two categories: quantitative problems and short essay questions. They will be similar to those from your homework assignments, worksheets, and quizzes. You are expected to give a logical, well organized, and complete solution for each problem. The answer only responses to problems will not receive much credit. You do not need to show your work for conversions that involve metric prefixes. Your responses to essay questions should be short, but complete; they should demonstrate your knowledge of specific information (not just your opinion) as it was presented and discussed in lecture and/or the textbook. Your random personal thoughts on a subject will generally get little credit. During the exam, you are not allowed to use any dictionaries and you are not allowed to use any electronic devices except for your own electronic calculator. You cannot share a calculator with a classmate or use your cell phone in place of a calculator. Your cell phone cannot be in your hands at any time during the exam. It must be kept in your pocket or in your bag. In case you have neither, you should place it on the instructor’s desk for the time of the exam. General Tips on Preparation. 1. Review your own lecture notes or, in case you are not a good note taker or missed a lecture, borrow the notes from a classmate. On the exam you should expect questions on any material covered in lecture and lab including instructor’s demonstrations. You will not be tested on any topics that are covered in the current chapters of the textbook, but have not been discussed in lecture or lab. 2. Complete all the worksheets from lecture and lab (even if you have not submitted them for grading) and review your graded quizzes. Blank copies of most of the worksheets and of all quizzes are posted on the course website along with the corresponding answer keys. Problems on the exam will be similar to problems from the worksheets, homework assignments, and quizzes. 3. Review the exam outline below and identify areas that you need to work on. Review the related textbook material and try to solve as many problems from the textbook as possible. Make sure you know where to find the answers to the textbook problems. You may consider a problem solved only if your answer matches closely the answer from the textbook. Textbook Sections to Study Chapter 1: all sections 1.1 – 1.3, 1.5; summary on page 9. Chapter 2: sections 2.2 – 2.10; summary on page 42. Chapter 3: sections 3.2 – 3.6, 3.10 (covered in lecture); 3.8, 3.9, 3.11, 3.12 (covered in lab); summary on pages 81 – 82. Chapter 4: sections 4.2 – 4.6, 4.8, 4.9; summary on pages 116 – 117. Chapter 5: sections 5.1 – 5.3 (omit polyatomic ions). Chapter 6: sections 6.3, 6.4, 6.6, 6.7. Laboratory Work to Know Exp. 2 “Graphing and Density”; Exp. 3 “Temperature”; Exp. 4 “Separation of a Mixture”; Exp. 14 “Calorimetry”; and Exp. 5 “Percentage of Oxygen” EXAM OUTLINE CLASSIFICATION OF MATTER; NAMES AND SYMBOLS OF ELEMENTS Be able to compare and contrast: mixture and (pure) substance, substance and molecule, element (elementary substance) and compound, mixture and compound, atom and molecule, molecule and formula unit. Know physical states for all elementary substances at room temperature: Br2 and Hg are liquids; H2, N2, O2, F2, Cl2, He, Ne, Ar, Kr, Xe, and Rn are gases; all other elementary substances are solids. Know all the elements that exist as diatomic molecules when isolated in pure form: H2, N2, O2, F2, Cl2, Br2, and I2. MEASUREMENTS, UNITS, METRIC SYSTEM, UNIT CONVERSIONS Have an understanding of the process of measurement as a comparison with a standard. Know that a unit of measurement serves as a standard. Know why measured quantities have a certain limited number of significant digits. How many significant figures should be written when you use a digital balance? How many significant figures should be written when you take measurements in centimeters using a metric ruler or in milliliters using a graduated cylinder? Be aware of rules for significant digits in calculations. There are the two different rules: one for multiplication/division and the other for addition/subtraction. You should to follow those rules when performing calculations that involve measured quantities. Know that expressions containing a number and a unit, for example 2 L and 5 dm3, are similar to algebraic entities such as 2x or 5y3. You are expected to always write measured quantities as numbers accompanied by appropriate units when you plug them into any mathematical formula. The result of your calculations must be a number accompanied by a unit that comes as a result of algebraic manipulation of units in the formula. You will lose points for not following these guidelines. Know the SI base units for length, mass, time, and temperature. What is the difference between base units and derived units? What are the SI units for area, volume, velocity (speed), and density? Know the temperature conversions: °C ↔ °F. Memorize as a number line: p = 10−12 n = 10−9 μ = 10−6 m = 10−3 c = 10−2 d = 10−1 k = 103 Practice changing units with metric prefixes by moving the decimal point or by changing the power of 10 without writing conversion factors (as we discussed and practice in class). MASS, VOLUME AND DENSITY Know the relationship between most common metric units of volume: m3, dm3, cm3, cm3. What is one liter? What is one mL? Refer to the cubic meter handout (class laminated handout with cubes). Imagine a cube with a side of one meter, a cube with a side of one decimeter, a cube with a side of on centimeter. How many smaller cubic units are in each of the cubes? Does the following list make sense? 1 m3 = 1000 dm3 = 1000 L 1 dm3 = 1 L 1 cm3 = 1 mL 1 mm3 = 1 μL Memorize: density of pure H2O at 4°C is 1.00 (3 sig. figs.) g/mL. The densities of other substances are not 1.00 g/mL and not need to be memorized. Do the following relationships for pure water make sense? 1 mL H2O = 1.00 g H2O 1 L H2O = 1.00 kg H2O 100 mL H2O = 100. g H2O 1 m3 H2O = 1.00×103 kg H2O What does density show? Does density of a homogeneous material at a fixed temperature and pressure depend on mass or volume? Is density a characteristic (independent of the amount of substance) property of substance? (Hint: What is the density of one ounce of water? Of one pound of water? Of one quart of water? Of one liter of water? Of one gallon of water?) What is a common unit of density we use in the lab? What is the SI unit of density directly constructed from the base units? Does the density of a material depend on temperature or pressure? Why? Density of which materials shows the greatest variation with temperature and pressure: solid, liquid, or gaseous? Why? Remember that conversion of a quantity expressed in any units of mass into a quantity expressed in any units of volume always requires the use of density as a conversion factor. (What does your small pink card say about ounces and fluid ounces?) You should be able to solve any density problems similar to those solved in class or those from your homework assignment. Wikipedia has a good article on density. Look it up! ENERGY, HEAT, HEAT CAPACITY, SPECIFIC HEAT Have a very basic understanding of the concepts of energy, heat, and temperature. Know the units of energy and heat: joule (J; a derived SI unit) and calorie (cal). Know the modern definition of calorie (1 cal is exactly 4.184 J) and the old one (the amount of energy required to increase the temperature of one gram of liquid water by one degree Celsius). Nutritionists used kilocalories, but call them simply calories!!! Be able to understand and use the formula: q = c×m×ΔT. You are recommended to memorize the specific heat of H2O(l): 1.00 cal/(g∙°C) and 4.184 J/(g∙°C). Know how calorimetry can be used to determine specific heats of substances. ATOMS AND MOLECULES; ATOMS AND SUBATOMIC PARTICLES Properties of subatomic particles: the charge in the units of one electron charge, approximate mass in amu, and the location within the atom. (Do not assume that the combined mass of free neutrons, protons, and electrons is the same as the mass of the atom that is made by combining those particles. Atomic physics and common sense are not exactly the same! Sorry.) What is the difference between atom and molecule? What is the total net charge of any atom or any molecule? What are the numbers of protons and electrons (relative to each other) in an atom or in a molecule? Know the definitions of mass number and atomic number. Look them up and memorize. Remember that both are exact whole numbers (unlike any quantity whose name ends with the word mass). Why are masses of atoms, molecules, and subatomic particles expressed in atomic mass units (amu) and not in grams or milligrams? ELEMENTS, ISOTOPES, ATOMS Know the two definitions of the term element as discussed in lecture: (1) type of atom and (2) elementary substance). Remember that neither element nor isotope mean one atom. An element is a collection of atoms (or a type of atom) and an isotope is a collection of atoms (or a type of atom). What is the difference? What is the principal numeric characteristic of an element? What are the two principal numeric characteristics of an isotope? How many different elements are known today? How many different types of atom (isotopes) are known today relative to the number of known elements? Be able to give a symbol for an atom or monatomic ion with specified numbers of subatomic particles. Using the table of naturally occurring isotopes, for a substance of a relatively simple composition (oxygen, water, methane, carbon dioxide), be able to find the number of types of molecule differing by isotope composition. ATOMIC MASSES AND ATOMIC WEIGHTS Know that atomic mass is the relative mass of atoms. The standard for atomic mass scale is an atom of carbon12. This atom by agreement has a mass set up at exactly 12 amu. No other atom has a mass in amu that is an exact number. Know the basic principles of operation of a mass-spectrometer. What information can be obtained in a massspectrometry experiment for a particular element (blue laminated handout with the table of isotopes)? Know what atomic weight is and how it is calculated. Know what molecular weight is and how it can be calculated from a chemical formula of a compound and the atomic weights of constituent elements. THE PERIODIC TABLE OF ELEMENTS Memorize symbols and names of 50 most common elements (the handout from the 1st class meeting). You should be able to correctly spell their names (!) and be able to quickly find them in the periodic table. Use a blank periodic table to write in their symbols into proper cells. Know what periods and groups are. Know positions for metals, nonmetals, and metalloids in the periodic table. Know the locations of main-group elements (also called representative elements) and transition metals in the periodic table. Which elements are liquids and gases and which are solids at room temperature? Memorize the following family names for group IA, group IIA, group VIIA, and group VIIIA elements: alkali metals, alkaline earth metals, halogens, and noble gases. Know how to answer the following questions. Can atoms be created or destroyed? (Can atoms be created or destroyed in a chemical reaction? Can molecules be created or destroyed in a chemical reaction?) What does it mean an artificial or synthetic element? What does it mean radioactive isotope? What does it mean radioactive element? What are the numbers that are given for each element in the periodic table of elements above and below the element’s symbol? Why for some elements the numbers below the symbol are in parentheses and have no decimal point? Moles and Formulas Distinguish atomic weight (=relative atomic mass), molecular weight (=relative molecular mass), formula weight (=relative formula mass), and molar mass. What are the units that are commonly used with values of each of them? The atomic weight of iodine is 126.90447 amu and the atomic weight of bromine is 79.904 amu. In nature, there are atoms of iodine that weigh 126.90447 amu, but there no atoms of bromine that weigh 79.904 amu. Why? The molecular weight of fluorine is 37.9968 amu and the molecular weight of chlorine is 70.906 amu. In nature, there are F2 molecules that weigh 37.9968 amu, but there no Cl2 molecules that weigh 70.906 amu. Why? Have a good understanding of the mole concept. What is the Avogadro’s number? Memorize the value of the Avogadro’s number with 4 significant digits: 6.022×1023 atoms, ions, molecules, formula units etc. per one mole of atoms, ions, molecules, formula units etc. Be able to perform calculations for elements and compounds that involve the use the Avogadro’s number and molar mass. Notice: g/mol = kg/kmol = mg/mmol. Be able to calculate the percent composition (% by mass) if the formula of a compound is known.