Name:______________________________________
Date:______________________ Per. #:_______
Chemistry
Semester 1 Final Exam Study Guide – Units 1-3
Vocabulary
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Conversion factor
Dimensional analysis
Atom
Nucleus
Electron cloud
Proton
Neutron
Electron
Charge
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Ion
Molecule
Formula Unit
Element
Substance
Mixture
Solution
Mole
Double displacement
Actual yield
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Periodic table
Element
Period
Group
Atomic number
Atomic mass
Symbol
Subatomic particle
Atomic mass units
Valence electron
Isotope
Homogeneous
Heterogeneous
Chemical Bond
Covalent Bond
Ionic Bond
Cation
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Avogadro’s Number
Stoichiometry
Theoretical yield
Percent yield
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Metals
Semimetals (metalloids)
Nonmetals
Alkali metals
Alkaline earth metals
Transition metals
Halogens
Noble gases
Atomic radius
Ionization energy
Electronegativity
Anion
Nomenclature
Polyatomic Ion
Physical property
Chemical property
Physical change
Chemical change
Molar Mass
Density
TOPICS
Unit 1 – States of Matter, SI Units, Atoms, and the Periodic Table
Notes 1.1 – States of Matter (pgs. 6 – 9)
o Know the major differences between the 3 states of matter: solid, liquid, and gas.
o Know whether each state has definite or indefinite volume or shape.
Notes 1.2 – SI Units (pgs. 33 – 38)
o Know the SI base units for length, volume, and mass (meter, liter, and gram, respectively)
o Understand that prefixes are used to multiply base units by powers of 10.
Notes 1.3 – SI Prefixes and Scientific Notation (pgs. 33 – 38, pgs. 50 – 54)
o Know the definitions and be able to make conversion factors for 5 SI prefixes: Kilo, centi, milli, micro, and nano.
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Example: Definition: centi means 1/100th
Conversion Factor: 1 m = 100 cm
o Know how to convert numbers from standard notation to scientific notation, and vice versa.
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Example:
Standard notation: 6,200,000
Scientific notation: 6.2 x 106
o Know how to solve multiplication and division problems using scientific notation.
Notes 1.4 – Intro to Dimensional Analysis (pgs. 40 – 42)
o Know how to solve dimensional analysis problems using SI base units and prefixes as conversion factors.
o The 5 Steps
1. Identify the given
2. What unit are you trying to convert to (the “arrow”)
3. Determine appropriate conversion factor(s)
4. Set up the equation. Start with the given and multiply by conversion factor(s) in fraction form. Cancel units.
5. Solve by multiplying tops and dividing by bottoms.
Notes 1.5 – Atoms and the Periodic Table (pgs. 67 – 69)
o Know Dalton’s 4 rules of Atomic Theory
o Know the basic model of the atom (nucleus made of protons and neutrons with electrons surrounding)
o Know how to read the periodic table and how to find the # of protons, neutrons, and electrons in an atom.
o Know what isotopes are.
Notes 1.6 – Reading the Periodic Table (pgs. 135 – 136)
o Know the difference between the atomic number and atomic mass. Know how to determine the number of
protons, neutrons, and electrons using the periodic table.
Name:______________________________________
Date:______________________ Per. #:_______
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Understand why the atomic mass number has a string of numbers after the decimal based on the weighted
average of isotopes.
Notes 1.7 – Atomic Experiments (pgs. 72 – 76)
o Know how J. J. Thomson used cathode ray tubes to discover the electron and deduce the existence of the proton.
o Know how Rutherford’s gold foil experiment worked and what feature of the atom he discovered with it.
o Know the difference between Thomson’s “plum-pudding” model and Rutherford’s model of the atom.
o Know how Mendeleev organized the periodic table and what he discovered about elements in the same group.
Notes 1.8 – History of the Periodic Table (pgs. 133 – 135)
o Know how Mendeleev originally grouped elements in his periodic table.
o Know why elements in the same family have similar reactivity.
o Know the special family/section names and where they are on the PT: alkali metals, alkaline earth metals,
transition metals, boron-nitrogen groups, halogens, noble gases, lanthanides, actinides.
Notes 1.9 – Periodic Trends (pgs. 150 – 162)
o Know how the size of an atom changes when it becomes a positive (+) or negative (-) ion.
o Know the trends of increasing/decreasing as you move down a group or to the right across a period for the
following characteristics.
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Atomic radius
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Ionization energy
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Electronegativity
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Metallic properties
o Know what conductivity, luster, malleability, and ductility mean in relation to metals.
Unit 2 – Types of Matter, Covalent and Ionic Bonding, Lewis Structures, Molecular Shape, Polarity,
IMFs
Notes 2.1 – Types of Matter – pgs. 11 - 14
o Know the major differences between the 4 types of particle: Atom, ion, molecule, and formula unit
o Know the differences between the 3 types of matter (or substance): Elements, compounds, and mixtures
o Know the difference between homogenous and heterogeneous mixtures and be able to identify them.
Notes 2.4 – Electron Configuration – pgs. 111 - 122
o Know how to use the Periodic Table to determine the electron configuration, Noble gas notation, or orbital
diagram of any element.
o Know what is meant by the Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule.
Notes 2.5 – Lewis Dot Structures and Covalent Bonds – pgs. 178 - 185
o Know what valence electrons are and how to find the number of valence electron for any main group element
using the periodic table.
o Know what the Octet Rule means.
o Understand how a covalent bond is made of at least two shared electrons between two nonmetals.
Notes 2.6 – Ionic Bonding and Nomenclature – pgs. 222 - 226
o Understand that an Ionic bond forms from the attraction of at least two oppositely charged ions.
o Know how to name covalent compounds using prefixes.
o Know how to name ionic compounds not using prefixes.
o Know how to get the formula of an ionic compound from the name.
Unit 3 – The Mole, Stoichiometry, Balancing Chemical Reaction Equations
Notes 3.1 – The Mole (pgs. 82 – 87)
o Know the definition of the mole as the number of atoms in exactly 12 grams of carbon-12.
o Know that a mole of ANY substance is equal in amount to Avogadro’s number, or 6.022  1023 particles.
o Know that the molar mass is the sum of atomic masses of every atom in a molecule and know how to find the
molar mass of any chemical using the periodic table.
o Know that the units of molar mass are in grams/mole.
o Know how to use Avogadro’s number, and the molar mass to convert between units of particles, moles, and grams.
Notes 3.3 – Chemical Reaction Notation (pgs. 261 – 269)
o Know how to get the formula of a covalent or ionic compound from the name.
o Know how to identify the reactants and products in a chemical reaction equation.
o Understand the different uses of coefficients and subscripts in a chemical reaction equation.
o Know how to identify a symbolic equation from the word form of an equation.
Notes 3.4 – Balancing Chemical Reaction Equations (pgs. 270 – 274)
o Know how to balance chemical reaction equations.
o Know how to deal with subscripts that are outside of parentheses, as in the 2 in Pb(NO 3)2.
Notes 3.5 – The Five Types of Chemical Reaction (pgs. 276 – 283)
o Know how to identify double displacement reactions
o Know how to find the products of double displacement reactions if given the reactants.
Name:______________________________________
Date:______________________ Per. #:_______
Notes 3.6 – Stoichiometry (pgs. 299 – 301, 304 - 311)
o Know how to use the mole ration from a balanced chemical reaction equation as a conversion factor.
o Know how to convert from moles or grams of a reactant to moles or grams of a product.
Notes 3.7 – Theoretical Yield (pgs. 312 – 317)
o Know how to find the theoretical yield of a product when given the mass of both reactants.
Know the difference between substances and mixtures and chemical vs. physical properties
o Know the difference between homogeneous and heterogeneous mixtures.
Practice Problems
Unit 1
1. Convert each base unit to the following prefixed units.
Base Unit
KiloCenti-
Milli-
5 meters
5000 mm
0.005 km
500 cm
6 meters
6 grams
5.5 liters
100 grams
2. A rock is weighed on a balance and is found to have a mass of 125 grams. When put in a graduated
cylinder, the water level rises from 25 mL to 50 mL. What is the density of the rock?
Mass
Volume (change)
Calculation
Density
3. Which would you expect to have a larger atomic radius: Magnesium (Mg) or Sulfur (S)? Explain.
4. Which would you expect to have a larger electronegativity: Lithium (Li) or Fluorine (F)? Explain.
5. How many liters are in 15.5 milliliters?
Arrow:
Equation:
_________________  _________________
Conversion Factors:
Solution:
Name:______________________________________
Date:______________________ Per. #:_______
6. How many meters are in 12,334 centimeters?
Arrow:
Equation:
_________________  _________________
Conversion Factors:
Solution:
Unit 2
1. What type of particle is each of the four examples below?
a) Bi_____________ b) Fe2+______________ c) H2O_______________ d) NaCl_________
2. Label each of the following as an element, compound, or mixture. For any example that is a
mixture, tell whether it is homogenous or heterogeneous.
Sample
Type of Substance
Type of Mixture
a) air
b) chlorine gas (Cl2)
c) rust (Fe2O3)
3. Name the following covalent compounds
Example: NF3  nitrogen trifluoride
a) SF6 ______________________________________________________________
b) N2O5 ______________________________________________________________
c) C4H8 ______________________________________________________________
4. Name the following Ionic compounds
Example: CaBr2  calcium bromide
a) CuCl2 ______________________________________________________________
b) NaF ______________________________________________________________
c) Al2O3 ______________________________________________________________
5. Determine the correct formula for the following ionic compounds.
Example: Sodium chloride  Na+ and Cl-  NaCl
Iron (III) oxide  Fe3+ and O2-  Fe2O3
a) calcium oxide ____________________________________________________
b) copper (II) bromide ___________________________________________________
c) sodium nitrate ____________________________________________________
Name:______________________________________
Date:______________________ Per. #:_______
6. Write the electron configuration, Noble gas configuration and orbital diagram for Cl and Fe.
Name
Electron Configuration
Noble Gas Configuration
Chlorine, Cl
Orbital Diagram
Iron, Fe
Orbital Diagram
Unit 3
1. Identify the number of moles of each substance.
# grams
# moles
# atoms/molecules
a) 32 grams O2
d) 1.204  1024 molecules NH3
b) 128 grams SO2
e) 3.011  1023 atoms Si
c) 22 grams He
f) 9.033  1023 molecules Cl2
2. Find the molar mass of the following substances.
a) Ga ______________________
b) MgSO4 ______________________
c) CH3OH ______________________
d) Ca(OH)2 ______________________
3. How many atoms are in 25 grams of MgSO4?
4. What is the mass in grams of 3.011  1024 atoms of sodium (Na)?
# moles
Name:______________________________________
Date:______________________ Per. #:_______
5. Write and balance the following chemical reaction equation in symbolic form.
“Potassium iodide reacts with lead (II) nitrate and yields lead (II) iodide and potassium
nitrate.”
6. Balance the following reaction equations.
a) ____ Cr + ____O2  ____Cr2O3
b) ____ Fe + ____Cl2  ____FeCl3
c) ____ C6H8 + ____O2  ____CO2 + ____H2O
d) _____CaCl2 + _____AgNO3  _____AgCl + _____Ca(NO3)2
e) ____Al + ____CuO  ____Al2O3 + ____Cu
Use the above reaction equations to solve the following stoichiometry problems. Use the GUESS
method.
7. How many moles of carbon dioxide can be produced from 4 moles of C6H8?
8. How many grams of aluminum is needed to produce 150 grams of aluminum oxide (Al2O3)?
Name:______________________________________
Date:______________________ Per. #:_______
9. If 45 grams of potassium oxide (K2O) is reacted with 45 grams of water, what will be the
theoretical yield, in grams, of potassium hydroxide (KOH)?
_____ K2O + _____H2O  _____KOH
Molar mass of K2O:__________________
Molar mass of H2O:__________________
Molar mass of KOH:__________________
#1: 45 g K2O  # g KOH
#2: 45 g H2O  # g KOH