Download CHEM 400 - El Camino College

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.