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CP Chemistry – Final Review
Chapter 1
1. What is chemistry?
2. List the 5 areas of study in chemistry.
3. What are some current areas of research in chemistry?
Chapter 2
1. Identifying substances 2.1: Define physical property, give examples of physical
properties:
2. States of Matter 2.1: List the three states of matter and the properties of each (see
figure 2.3 on page 41)
3. Classifying mixtures 2.2:
a. Define heterogeneous mixture and give examples.
b. Define homogeneous mixture and give examples.
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4. Distinguishing Substances from mixtures 2.3: What defines a substance vs. a
mixture?(hint: Key point).
5. Chemical symbols and formulas 2.3: What is the difference between chemical symbols
and formulas? (hint: key point)
6. Recognizing Chemical Changes 2.4: What 4 clues serve as a guide that a chemical
change has taken place?
7. Conservation of Mass 2.4: Define the Law of Conservation of Mass.
Chapter 3
1. Accuracy, Precision and Error 3.1: Define each, understand Key Point.
2. Understand Scientific Notation 3.1:
a. Express in scientific notation:
0.00012 _____________________
421000000___________________
10000_______________________
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3. Significant Figures 3.1: Understand Key Point
a. How many significant figures is in:
0.05730 meters _____________
.00073 meters ______________
143 grams _________________
4. SI Units 3.2: What are the five SI base units commonly used?
5. SI Units 3.2 : What is Kelvin, how do you convert from Celsius?
6. Density 3.4: What is density, what is the formula for density?
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Chapter 4
1. Early Models of the Atom 4.1: Describe Democritus’s philosophy
2. Early Models of the Atom 4.1: Describe Dalton’s Theory
3. Subatomic Particles 4.2 : Describe the kinds of subatomic particles, including their
charge and location in the atom.
4. Atomic Number 4.3: What is the atomic number represent?
5. Mass Number 4.3: What does the mass number represent?
6. Isotopes 4.3: Define isotope:
7. Atomic Mass 4.3 : What unit is used measure atomic mass?
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Chapter 5
1. The Quantum Mechanical Model 5.1 : define quantum mechanical model
2. Electron Configurations 5.2: understand the location of s, p, d, & f orbitals on the
periodic table
3. Electron Configurations 5.2: Review Table 5.3 and writing electron configurations
4.
Electron Configurations 5.2: review which atoms are most stable in each level, for
example p6 would be the most stable for the p level
5. Quantum Mechanics 5.3: What are photons?
6. Quantum Mechanics 5.3: Understand that Erwin Schrodinger developed the Quantum
Mechanical Model of the atom.
Chapter 6
1. The Periodic Law 6.1: Explain groups and periods on the periodic table:
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2. Metals, Nonmetals, and Metalloids 6.2: Define each of these, which one is in greatest
abundance on the periodic table?
3. Metals, Nonmetals, and Metalloids 6.2: Be able to locate each group on the table and
identify elements in these categories.
4. Squares in the Periodic Table 6.2: Be able to identify the data in a square on the
periodic table (figure 6.8, page 161)
5. Electron Configurations in Groups 6.2 Be able to name an element based on it’s
electron configuration: example: 1s22s22p6 is Neon
6. Trends in Atomic Size 6.3: Describe atomic size as you move across the periodic table
and down the periodic table.
7. Ions 6.3: Explain how ions form.
a. What are cations?
b. What are anions?
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Chapter 7
1. Valence Electrons 7.1: How do you find the number of valence electrons? What is a
valance electron?
2. Formation of Cations, Formation of Anions 7.1: Understand electron configuration of
cations and anions and how they differ from regular elements.
3. Formation of ionic compounds 7.2: Which types of elements form ionic compounds?
4. Properties of ionic compounds 7.2: What are the 3 properties of ionic compounds?
5. Metallic Bonds and Metallic properties7.3: What do valance electrons do in metals?
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6.
Metallic Bonds and Metallic properties 7.3: What are the characteristics of metals?
Why are they good electrical conductors?
Chapter 8
1. The Octet Rule in Covalent Bonding 8.2: Why do atoms share electrons in covalent
bonding?
2. The Octet Rule in Covalent Bonding 8.2: What are diatomic molecules?
3. The Octet Rule in Covalent Bonding 8.2: Describe single covalent bonds. Give an
example and draw the Lewis dot structure for it.
4. Double and Triple Covalent Bonds 8.2: Describe double covalent bonds. Give an
example and draw the Lewis dot structure for it.
5. Double and Triple Covalent Bonds 8.2: Describe triple covalent bonds. Give an example
and draw the Lewis dot structure for it.
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Chapter 9
1. Monatomic ions 9.1: Identify the charges associated with cations and how they relate
to their group on the periodic table. Give the charges for the ions of: Na, Ca, Al
2. Monatomic ions 9.1: Identify the charges associated with anions and how they relate
to their group on the periodic table. Give the charges for the ions of: N, O, F
3. Monatomic ions 9.1: What type of ions end in ide?
4. Ions of Transition Metals 9.1: In the Stock system of naming transition metals, what
types of numbers are used to indicate the charge?
5. Polyatomic Ions 9.1: What do the endings ite and ate represent?
6. Naming binary Ionic Compounds 9.2: How are the chemical formulas of binary ionic
compounds written? (key point)
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7. Naming binary Ionic Compounds 9.2: Ionic compounds are made up of a monatomic
metal cation (Groups 1-3) and a monatomic nonmetal anion (Groups 5-7), Give several
examples of binary ionic compounds:
Chapter 10
1. What is a Mole? 10.1: What is the SI unit used to measure the number of
representative particles in a substance?
2. What is Avogadro’s Number?
3. The Mass of a Mole of an Element 10.1: understand the Key Point – The Atomic Mass
of an element expressed in grams is the mass of a mole of the element.
What is the atomic mass of Na ______________ C ________________ Au ___________
4. The Mass of a Mole of a Compound 10.1: How do you calculate the molar mass of a
compound (key point)?
What is the molar mass of NaCl?
What is the molar mass of NH3?
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5. The Percent Composition of a Compound 10.3: What is the formula for % mass of an
element?
If 20.0 grams of Ca combines completely with 16.0 grams of S to form a compound,
what is the percent composition of Ca in the compound?
Chapter 11
1. Writing Chemical Equations 11.1: What is a skeleton equation?
2. Balancing Chemical Equations 11.1: Balance the following equation:
AgNO3 + H2S
Ag2S + HNO3
3. Classifying reactions 11.2 : What are the five general types of reactions? Define each.
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Chapter 13
1. Kinetic Theory and a Model for Gases 13.1: Define Kinetic theory and list the three
fundamental assumptions about gases (key points)
2. Gas Pressure 13.1: What instrument is used to measure atmospheric pressure?
3. Gas Pressure 13.1: What is the SI unit for pressure? What is normal pressure at STP?
4. The Nature of liquids 13.2: Describe what happens on a particle level during
evaporazation:
5. A Model for Solids 13.3 : Describe the properties of solids:
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Chapter 15
1. Water in the Liquid State 15.1: What is surface tension? Describe how the surface
tension of water compares to the surface temperatures of most other liquids:
2. Solvents and Solutes 15.2: What is a solution? Can it be filtered?
3. Electrolytes and NonElectrolytes 15.2: Ionic compounds are always electrolytes in
solution. Give an example of an ionic electrolyte solution:
4. Electrolytes and NonElectrolytes 15.2: How is a Hydrate written?
Chapter 19
1. Properties of Acids and Bases 19.1: List the properties of Acids:
2. Properties of Acids and Bases 19.1: List the properties of Bases :
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3. Arrhenius Acids and Bases: 19.1 : Define an Arrhenius Acid (key point)
4. Hydrogen Ions from Water 19.2 : At a pH of 7 (neutral) what is the concentration of
hydrogen ions?
5. The pH Concept 19.2: Acids have a ph of __________________ than 7
Bases have a ph of ___________________ than 7
Neutral is a pH of ____________________