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Transcript
CHEM 4
OUTLINE FOR EXAM 1
TEXT: Chapters 2, 3, 5, 6 (sections 7.1-7.2 of Ch. 7 and section 10. 7 of Ch. 10)
LAB: Experiments 1 and 2
I.
MEASUREMENTS and CALCULATIONS
Metric (SI) System (Ch. 3, p. 49, 61-64; Appendix II at the end of the text; Exp. 2)
Know the metric system's basic units of mass, length, volume, temperature and time.
Know the meaning of and symbols for the common metric prefixes (mega-, kilo-, deci-, centi-, milli-, micro-, nano-).
Given a measurement in the metric system, convert it to any other related unit in the metric system.
Given a measurement in the metric system, convert it to the corresponding unit in the English system and vice versa.
(Know these English-metric equivalents: 2.54 cm = 1 inch exactly; 453.6 g = 1 lb; 3.785 L = 1 gal)
Significant Figures (Ch. 3, p. 67-74; Exp. 2)
Distinguish between accuracy and precision.
Given a measurement, state the number of significant figures it has.
Round off a given number to a stated number of significant figures.
Express the sum or difference or product or result of a division of experimental measurements in the proper number of
significant figures.
Significant figures in calculations involving both multiplication/division and addition/subtraction.
Exponents and Scientific Notation (Ch. 3, p. 49-53)
Express numbers in scientific notation.
Convert decimal numbers to exponential form and vice versa.
Be able to perform all operations involving scientific notation.
Problem Solving and Dimensional Analysis (Ch. 3, p. 53-60, especially the table on p. 60; Ch. 7, p. 174 “The
Plan”; Exp. 2, lecture notes & problem sets)
Set up and solve problems by analyzing units (the unit conversion method or dimensional analysis).
Know what conversion factors and how to use them including problems using multiple conversion factors.
Solve the problems the way we did in class: (1) state the main question (given quantity  desired quantity); (2) identify
conversion factors that are given in the problem and think about additional conversion factors that you may need and
where to get them; (3) outline the complete solution map starting with units of the given quantity; (4) write a given
quantity (number and units) and multiply it by appropriate conversion factors in the order they appear in the solution
map; (5) check the units; (6) perform the calculation remembering about significant digits; (7) check if the final answer
makes sense.
Problems involving squared or cubed units require special attention. Know how to solve such problems.
Solve problems involving conversions of units in both the numerator and the denominator (density, speed, price, car
efficiency).
Density (Ch. 3, p. 61, 80-84; Exp. 1; lecture notes & problem sets)
Be able to explain mass, weight, volume. Units for mass and volume.
Define density. What are its units? What is the SI unit for density?
What do we mean when we say that density is an "identifying" or "characteristic" property of a material?
Know the value of the density of water at 4°C.
Given two sets of the following, calculate the third: mass of a sample; volume occupied by the sample; density of the
sample. Density as a conversion factor: VOLUME   MASS
Why do you think density changes with temperature?
DENSITY
Temperature (Ch. 2, p. 33-35; Ch. 3, p. 72-75; Ch. 10, p. 289-290)
Be able to explain energy, heat, exothermic and endothermic.
Units of energy: joule (J; SI unit), calorie (cal), and Calorie (Cal).
What is the meaning of temperature?
Given a temperature in F or C or K, convert from one scale to another.
Laboratory (Exp. 2; Measurement Supplement)
Given a set of measurements, calculate the average value and deviation of each measurement.
How the density of a solid or liquid can be determined?
Given the true value and an experimental value, calculate error and percent error.
Know how to read measurements from the ruler, electronic balance, hanging pan balance and graduated cylinder.
II. MATTER
Elements and their Symbols (Ch. 2, p. 30; Ch. 5, p. 133-135; Ch.6, Fig. 6.1 on p. 143; Formulas Supplements 1 &
2; handout)
Know the chemical symbols for the most common elements and be able to spell each name.
Know the elements that naturally occur as diatomic molecules: H2, O2, N2, F2, Cl2, Br2, I2.
Building Blocks Derived from Atoms (Ch. 2, p. 30; Ch. 6, p. 142, 146)
Distinguish between:
(a) an atom and an ion
(b) an atom and a molecule
(c) an atom and a formula unit
(d) a molecule and a formula unit
(e) an anion and a cation
(f) an element and a compound
Isotopes (Ch. 5, p. 124-126)
Explain how we can have 114 known elements, but about 1,000 different kinds of atoms.
What is an isotope?
Know the meaning of atomic mass.
Calculate the atomic mass from percent abundances of the naturally occurring isotopes of an element.
Introduction to Atomic Structure (Ch. 2, p. 35; Ch. 5, p. 120-123; Ch. 6, p. 146)
What are the main ideas of Dalton's theory?
What are the main ideas of the nuclear theory of matter?
Identify the three principal components (subatomic particles) of the atom, stating their electrical charges and relative
masses.
Describe the general arrangement of subatomic particles in the atom.
Given the atomic number and mass number of an atom, state the number of protons, electrons, and neutrons.
Given the symbol of a monatomic ion, state the number of protons and electrons in it.
Given the identity of the element from which an ion was derived and the number of electrons, write the symbol for the
ion, including the charge.
Compounds Display Constant Composition (Ch. 2, p. 32; Ch. 7, p. 174-178)
What is the Law of Constant Composition?
Compounds are represented by chemical formulas, elements by chemical symbols.
Know how to read and calculate the number of atoms in a chemical formula.
What are molecular mass and formula mass?
Classification of Matter (Ch. 2, p. 16, 19-20, 24-29; handouts)
Identify and distinguish between the three states of matter.
Explain the difference between gases, liquids, and solids in terms of particle behavior.
Distinguish between homogeneous and heterogeneous samples of matter.
Distinguish between a pure substance and a homogeneous mixture (or as such a mixture is sometimes called, a solution).
Distinguish between miscible and soluble.
Know the meaning of terms solution, solvent, solute and examples.
Distinguish between an element and a compound.
Identify examples of ionic and covalent compounds using chemical formulas.
In what ways do the characteristics of ionic and covalent compounds differ?
General Properties and Changes of Matter (Ch. 2, p. 21-23)
Distinguish between physical and chemical properties.
Distinguish between physical and chemical changes.
Give at least three kinds of evidence that would indicate that a chemical reaction has occurred.
Laws of Conservation of Mass and Conservation of Energy (Ch. 2, 38)
Introduction to the Periodic Table (Ch. 5, p. 129-131)
What attribute/property characteristic of each element is considered when arranging all of the elements in consecutive
order in the periodic table?
In which direction, horizontal or vertical, are elements with similar chemical properties aligned? What are classes or
collections of elements with similar chemical properties called?
Recognize examples of metals and nonmetals using the periodic table.
Discuss the trend of metallic character within the periodic table. What is a metalloid?
Give the characteristic properties of metals and those of nonmetals.
What is a group in the periodic table? What is a period?
Locate within the periodic table: alkali metals; alkaline earth metals; halogens; noble gases; transition metals, main
group (representative) elements
III. CHEMICAL NOMENCLATURE
(Ch. 6, p. 146-164, especially tables; handouts & worksheets)
Naming Compounds/Writing Formulas
(a)
All, binary compounds end in -ide.
(b)
Naming Ionic Compounds:
Type I
Ionic compounds containing metallic ions of fixed charge and ammonium ion.

(c)
(Groups IA, 2A, & 3A cations plus Ag+, Cd2+, Zn2+, Ni2+; and NH 4 )
Type II
Ionic compounds containing metals that form ions of more than one charge
(Most transition-metal cations plus metal cations of Groups 4A, & 5A)
Naming Binary (= two elements) Covalent Compounds (They contain only nonmetals). Almost all covalent
compounds are molecular compounds.
(d)
Given the name of a compound containing a metal and a nonmetal, or a metal and a polyatomic ion, write the
correct formula, or given the formula, write the name.
(Do NOT use Greek prefixes in the name of an ionic compound.)
(e)
Distinguish between the cation suffixes –ous and –ic (old names); the anion suffixes -ite, -ate, per - ate, and
hypo - ite.
(f)
Given the name of a compound containing two nonmetals, write the correct formula, or given the formula,
write the name. (This is a covalent compound. Use Greek prefixes, as needed, to name.)
(g)
Distinguish between oxygen-containing acids and nonoxygen acids and be able to name each of them as pure
covalent compounds and as inorganic acids (i.e. as aqueous solutions) when given their formulas and their
formulas when given their names.
(1)
nonoxygen acids are hydro -ic acid rule
(2)
oxygen-containing acids: use the -ous/-ic acid rule
(e)
What is a hydrate? Name a salt hydrate if given the formula. Write the formula if given the name.
IV. CALCULATIONS INVOLVING ELEMENTS, COMPOUNDS, PERCENT
Atomic Mass (Ch. 5, p. 125, 127-128)
(a) What is the difference between mass number and atomic number?
(b) What is an atomic mass unit (amu)?
(c) Given the name or symbol of an element or the atomic number of an element, by referring to the periodic table,
state the atomic mass of the element.
(d) What are the two numbers that are given for each element in the periodic table?
Molecular Mass/Formula Mass (Ch. 7, p. 174-178)
(a) Distinguish between a molecule and a formula unit.
(b) What is the difference between atomic mass & molecular mass;
formula mass & molecular mass?
Percent (Ch. 7, p. 184; class worksheet; Exp. 2)
(a) What is the meaning of the term percent? How does percent differ from fraction? How does fraction differ
from ratio?
(b) Know how to use percent as a conversion factor (part/total, part A/part B). Think of 100 grams (or 100 kg or
100 tons) of a substance!
(c) Given experimentally measured values and the accepted value, calculate the percent error. Make sure to
properly handle significant figures in the calculations.
HINTS:
1.
2.
3.
4.
5.
6.
Practice writing by writing definitions and explanations of concepts. Take notes when you read your
textbook.
Go over the lab reports, quizzes and homework (making sure that you correct your mistakes).
Get copies of the notes you missed from a reliable source.
Come to the exam rested and knowing that you'll do the best you can.
On the exam read questions very carefully. Reread the question making sure you really understand it.
Be careful with calculations. If you have doubts about your numerical answer, redo the calculation.
(Get to know your calculator before the exam.)