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Chemistry Semester 1 Course Review
Name__________________________ Date___________ Per___
Italics indicate required honors extensions
Unit 1: Nature of Science
Essential questions:
 What rules must be obeyed to safely conduct an experiment?
 What are the components of a good scientific experiment?
 What rules must be obeyed to safely conduct an experiment?
 Why are significant figures important to chemists?
 What is the best method/graph to represent specific data?
 How would a scientist organize data collected from an experiment into a graph?
 How does a standard notation number compare to a scientific notation number?
 How do you differentiate between accuracy and precision of data?
 Can you calculate % error of the results from an experiment?
 What are the SI base units?
 How is dimensional analysis used to convert units within the metric system?
 Why does the Big Bang Theory exist?
 Explain scientifically behind how the universe was created according to the Big Bang Theory?
Key vocabulary:
 Scientific method
 Significant figures
 Metric system
 Base units
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Big Bang Theory
Experiment
Hypothesis
Variable
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Control
Accuracy
Precision
Percent error
Practice:
1. Using the picture above, list lab safety rules that are being
ignored.
2. What are the steps to the scientific method?
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Chemistry Semester 1 Course Review
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3. Convert the following into scientific notation:
a) 1500 = ____________
c) 0.001012 = ____________
b) 123 = ____________
d) 1.52 = ____________
4. Convert the following into decimal notation (ordinary notation):
a) 4.59 x 103 = ____________
c) 280 x 10-4 = ____________
b) 5 x 102 = ____________
d) 1.4 x 10-5 = ____________
5. Indicate the number of significant figures:
a) 34 g ___
b) 564 L ___
c) 19.3 mm ___
d) 23.45 mg ___
e) 101 km ___
f) 3400 g ___
6. Round the following numbers to 2 significant figures:
a) 0.826mg ___
b) 19.88mL ___
c) 19250cm ___
d)950L ___
7. What is the SI unit for the following measurements:
a) length _____
b) mass _____
c) time _____
d) volume _____
8) What measurement in millimeters is indicated on the ruler below? _____
9) Convert the following:
a) 850cm to _________ mm
b) 2500 mg to __________ kg
c) 0.2598kL to _________ L
10) Suppose a lab refrigerator should hold a constant temperature of 38.0 F. A temperature sensor is tested 10 times in the
refrigerator. The temperatures from the test yield the temperatures of: 37.8, 38.3, 38.1, 38.0, 37.6, 38.2, 38.0, 38.0, 37.4, 38.3. Is the
distribution of values from the test: accurate (yes or no), precise (yes or no).
11) Explain the Big Bang Theory. _________________________________________________________________________________
____________________________________________________________________________________________________________
____________________________________________________________________________________________________________
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Chemistry Semester 1 Course Review
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12) What is the equation for percent error?
13) A researcher measures the mass of a sample to be 5.51 g. The actual mass of the sample is known to be 5.80g. Calculate the
percent error.
Unit 2: Atomic Structure
Essential questions:
 What is an atom?
 What are the early models of the atom and how has scientific exploration lead to the current model?
 What is a theory?
 How do you identify the relative mass, relative charge, and location of the three smaller subatomic particles of an atom?
 What is the overall charge of an atom?
 What is the relationship between the subatomic particles and their charges, masses, and locations?
 What is an isotope?
 What is the average atomic mass compared to atomic mass?
 How do you calculate the average atomic mass of elements?
 Identity an element when supplied with the natural abundance and mass of each isotope.
Key vocabulary:
 Atom
 Compound
 Subatomic Particles
 Proton
 Neutron
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Electron
Nucleus
Atomic Number
Average Atomic Mass
Mass Number
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Isotope
Neutral
Theory
Practice:
1. What is an atom? _________________________________________________________________________________________
2.
What is the overall charge of an atom? Why? ___________________________________________________________________
________________________________________________________________________________________________________
3.
4.
Isotopes are atoms of the same element, which have the same number of (protons / neutrons) but a different number (protons
/ neutrons).
How do isotopes C-12 and C-14 differ from each other? ___________________________________________________________
How are they similar? ______________________________________________________________________________________
5.
What parts of Dalton’s atomic theory are now known to be incorrect? _______________________________________________
________________________________________________________________________________________________________
________________________________________________________________________________________________________
________________________________________________________________________________________________________
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Physical Science Semester 1 Course Review
6.
Name__________________________ Date___________ Per___
Progression of the atomic model:
Dalton
Thompson
Rutherford
Description of
supporting
evidence &
experiments
Major findings
Drawing and
description of
atomic model
7.
Identify the three basic particles in the atom. Give their location, charge, and mass. (Fill in the table)
Particle
Location
Charge
Mass (amu)
8.
Fill in the blanks for the elements using the periodic table
Element
Symbol
Atomic #
#Protons
Carbon-14
# Electrons
#Neutrons
Mass #
29
16
131
53 I
47
9.
Silicon has three naturally occurring isotopes. Calculate the average atomic mass of Silicon with the information provided. Check
your answer using the periodic table.
Isotope name
Isotope mass (amu)
Relative Abundance
Silicon-28
27.98
92.21
Silicon-29
28.98
4.70
Silicon-30
29.97
3.09
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Physical Science Semester 1 Course Review
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Unit 3: Electron
Essential questions:
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How does an electron act according to de Broglie’s wave-particle duality?
What is a quantum?
In what ways do the Bohr model and quantum mechanical model differ?
How does the quantum mechanical model describe the arrangement of the electrons in atoms and their orbitals?
What happens when electrons in atoms absorb or release energy?
How is the electron arrangement of an atom indicated by electron configurations and orbital diagrams?
How are electron configurations different from noble gas configurations?
How is atomic emission spectra of elements used to identify elements?
What is the relationship between a wave’s frequency, energy, and wavelength?
Key vocabulary:
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Wave-Particle Duality
Electron
Excited State
Ground State
Quantum
Valance Electron
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Frequency
Energy
Wavelength
Electron Configuration
Noble Gas Configuration
Electromagnetic Spectrum
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Orbitals
Period
Energy Level
Cloud Shapes
Practice:
1. Electrons can act as a __________ or ___________ according to the _____ -__________ duality model.
2.
What is a quantum of energy? __________________________________________________________________
3.
When an electron absorbs energy, it jumps to the ___________ state.
4.
How does an electron produce light in things such as fireworks or neon signs? ______________________________
________________________________________________________________________________________
Use the diagram to answer questions 5-7:
5.
Which region is referred to as the s-block on the diagram? ____
How many electrons can each level of the s-cloud hold? ____
6.
Which region is referred to as the p-block on the diagram? ____
How many electrons can each level of the p-cloud hold? ____
7.
Which region is referred to as the d-block on the diagram? ___ How many electrons can each level of the d-cloud hold? ___
8.
What energy level are the outer electrons for Potassium found? ______ This is indicated by Potassium’s ______________.
9.
Indentify the number of valance electrons for the following elements and draw the corresponding electron-dot structure:
a)
Calcium
10. Draw the shape of the s-sublevel:
b) Carbon
c) Neon
Draw the shape of the p-sublevel:
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Chemistry Semester 1 Course Review
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11. Write the electron configuration for:
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Lithium:
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Iron:
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Argon:
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Barium:
12. Identify the element with the electron configuration of 1s22s22p63s23p3 ____________________
13. Identify the element with the electron configuration. [Ne]3s23p4 _________
14. According to the quantum mechanical model, the location of an electron is based on ______________.
15. What does the electromagnetic spectrum show? _____________________________________________________________
_____________________________________________________________________________________________________
16. On the electromagnetic spectrum, as the wavelength increases, frequency ___________, and energy ____________.
17. A barium atom (gains / loses) ______ electrons when it forms a barium ion. What is the symbol for a barium ion? ______
A fluorine atom (gains / loses) ______ electrons when it forms a fluorine ion. What is the symbol for a fluorine ion? ______
18. Write the orbital notation for:
 Carbon:
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Iron:
Unit 4: Periodic Table
Essential questions:
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How is the periodic table organized according to Mendeleev and Moseley?
What information does the periodic table provide?
What property of elements is used to organize the periodic table?
How many groups and periods are on the periodic table?
What are the family names for groups 1,2,3-12, 17,18?
Where are the metals, nonmetals, and metalloids located?
How can periodic trends be explained?
Key vocabulary:
 Atomic Number
 Atomic Mass
 Atomic Symbol
 Group (family)
 Period
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Group Names
Metals
Nonmetals
Metalloids
Periodic trend
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Atomic radius
Ionic Radius
Electronegativity
Octet Rule
Practice:
Use the diagram to answer questions 1-3 :
1.
Which region contains group 18? ___ What is the name of this group? ____________
2.
Which region contains the alkaline earth metals? _____
3.
Which region contains elements with an electron configuration that ends with p 5? _____
4.
The 10 short columns in the middle of the periodic table make up the ___________________ _________________.
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Chemistry Semester 1 Course Review
5.
Name__________________________ Date___________ Per___
The order of the elements on the periodic table is based upon the ____________ ______________. What particle of the
atom is represented by this number? __________________
The diagram below lists the information found on the periodic table. Use this diagram to answer question 6.
79
Au
Gold
107.86
7.
How many groups are on the periodic table? ____________
8.
How many periods are on the periodic table? ____________
6.
The element’s _____________ _____________ is 79,
the _______________ is Au, and the Atomic mass is
______________.
Use the diagrams below to answer questions 9-10.
9.
Which diagram correctly depicts the trend in electronegativity? ________
10. Which diagram correctly depicts the trend in atomic radius? ________
11. Elements in the same group have similar ______________. They behave similarly because they have the same number of
________________.
12. Metals are located on the ____________ side of the periodic table and nonmetals are on the ____________ side.
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Chemistry Semester 1 Course Review
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Unit 5: Bonding
Essential questions:
 How do intermolecular forces differ from intramolecular forces?
 How is a positive and negative ion formed?
 Why does an element gain or lose electrons to become and ion?
 How do ionic compounds form?
 How does covalent bonding satisfy the octet rule?
 How is the bonding in Covalent (molecular) compounds different from the bonding in ionic compounds?
 How do the properties of an ionic compound compare to those of covalent (molecular)compounds?
 How many electrons are shared between atoms in a double and triple bond?
 How do electrons affect the shape of a molecule?
 How do lone pairs around a central atom affect the polarity of the molecule?
 Why can carbon form many compounds?
Key vocabulary:
 Ionic bond
 Covalent compound
 Valence Electron
 Intramolecular forces
 Intermolecular forces
 Ions
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Ionic compound
Cation
Anion
Polarity
Single, double triple bond
Physical properties
Covalent bond
VSEPR
Octet rule
Lone pairs
Practice:
1.
How does a intermolecular forces differ from a intramolecular force? ________________________________________
_________________________________________________________________________________________________
2.
Which electrons are involved in bonding? _________________________________
3.
What types of elements are involved in ionic bonding? ________________________ Covalent?___________________
4.
What type of bond is formed by the transferring of electrons? ________________ Sharing electrons? ______________
5.
How does an element form a cation?______________________________ Anion? _____________________________
6.
Why does Magnesium form a +2 charge? _______________________________________________________________
7.
How many valance electrons does each of the following elements have?
a)
8.
Sodium ________
b) Oxygen __________
d) Krypton __________
List the general characteristics of ionic and covalent bonds (i.e. physical properties (hardness, state of matter, boiling
point, melting point, conductivity), strength)
Ionic
Covalent
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Chemistry Semester 1 Course Review
9.
Name__________________________ Date___________ Per___
Using electronegativity values, what type of bond (ionic, polar covalent, nonpolar covalent) is formed between the
following elements:
a) F and F
b) H and I
c) Mg and Br
d) C and Cl
10. How many electrons are shared in a:
a) Single covalent bond ___
b) Double covalent bond ___
c) Triple covalent bond ___
11. Why do some elements form double and triple bonds during bonding? ________________________________________
_________________________________________________________________________________________________
12. Draw Lewis structures for the following. Identify the shape of the molecules (linear, bent, tetrahedral, trigonal planar
and pyramidal). Identify the polarity of the molecules.
Molecule
Lewis Structure
Shape
Polarity
Molecule
H2O
BF3
NFH2
PH3
CF4
CO2
Lewis Structure
Shape
Polarity
Unit 6: Chemical Nomenclature
Essential questions:
 How are ionic compounds identified?
 What is a monatomic ion?
 What are the common polyatomic ions’ names and formulas?
 How are cations and anions combined to make ionic formulas?
 How are Roman numerals used in ionic compound formulas?
 How do you identify ionic compounds versus covalent (molecular) compounds?
 What are the diatomic elements?
 How are molecular formulas translated into binary covalent compound names?
Key vocabulary:
 Ionic compound
 Covalent compound
 Valence Electron
 Ions
 Cation
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Polyatomic ions
Anion
metals
Transition Metals
Octet rule
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Nonmetals
Binary Compound
Diatomic Molecules
Subscript
Superscript
Practice:
1. Identify all seven diatomic elements: ___________________________________________________
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Chemistry Semester 1 Course Review
Name__________________________ Date___________ Per___
2. (Honors only) Write the formulas and charges for the following polyatomic ions:
hydroxide:
sulfate:
phosphate:
ammonium:
nitrate:
carbonate:
3. For the following ionic names, write the missing chemical formula or chemical name.
a. nickel (II) oxide
_____________
b. calcium carbonate
_____________
c. potassium nitrate
_____________
d. ammonium bromide
_____________
e. gold (III) iodide
_____________
f. zinc phosphide
_____________
g. lithium sulfate
_____________
h. Na2SO4
___________________________
i. Al2O3
___________________________ j. SnO
___________________________
k. K2S
___________________________ l. (NH4)3PO4
___________________________
m. ZnCl2
___________________________ n. PbCO3
___________________________
4. For the following molecular (covalent) compounds write the missing formula or name.
a. carbon monoxide
____________
b. xenon tetrafluroide
____________
c. silicon dioxide
____________
d. iodine pentachloride
____________
e. P2O5
___________________________
f. P4O10 ___________________________
g. SF7
___________________________
h. NI3
___________________________
Unit 7: The Mole
Essential questions:
 How is the mole used to count particles of a substance?
 How does the mole relate to other everyday counting units?

Why is the mole used to count atoms instead of a more common unit?
 How are moles of a substance converting to numbers of representative particles?
 How is molar mass of a substance calculated?
 How is the mass of an atom related to the mass of a mole of atoms?
 How are moles of a substance converted to mass?
 How are number of representative particles of a substance determined from its mass?
 What is the difference between an empirical and molecular formula of a compound?
 How is the percent composition of an element within a compound calculated?
 How are empirical and molecular formulas obtained from percent composition data?
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Key vocabulary:
Mole
Atoms
Avogadro’s Number
Representative Particle
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Hydrate
Molar Mass
Empirical Formula
Molecular Formula
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Subscript
Percent composition
Practice:
1.
Show all work for calculations:
Give conversion factors (in fraction form) for converting between the following
A.) FROM moles TO atoms
B.) FROM grams TO moles of Boron
C.) FROM atoms TO moles
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Chemistry Semester 1 Course Review
Name__________________________ Date___________ Per___
2.
What is the molar mass of Rb2Cr2O7?
3.
What is the % by mass of C in Pb(C2H3O2)2 ?
5.
What is the number of molecules in 16.75 g of H2O ?
6.
What is the mass of 1.75 x 1024 molecules of NH3 ?
7.
How many molecules are in 0.26 mol of CO2 ?
What is the molar mass of CoCl2  6H2O ?
4.
Which of the following are empirical formulas?
8.
H2O
H 2 O2
C6H12O6
N2H5
Unit 8: Chemical Reactions
Essential questions:
 What is evidence of a chemical change?
 How is a chemical change different from a physical change?
 How does a balanced chemical equation show relationships between the reactants and the product of a chemical reaction?
 How can chemical reactions be classified?
 What are the defining characteristics of each classification of chemical reaction?
 What are aqueous reactions?
 How are word equations translated into formula equations and vise versa?
Key vocabulary:
 Chemical Change
 Physical Change
 Coefficient
 Molar Ratio
 Conservation of Matter
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Balanced Reaction
Chemical Reaction
Synthesis
Decomposition
Single Replacement
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Double Replacement
Combustion
Aqueous Solution
Precipitate
Practice:
1.
Differentiate between a precipitate and an aqueous solution. ______________________________________________
____________________________________________________________________________________________
____________________________________________________________________________________________
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Chemistry Semester 1 Course Review
Name__________________________ Date___________ Per___
List the 5 types of chemical reactions AND state how you would identify their type simply by looking at the equation:
Type of reaction:
How you would identify the type simply by looking at the form of the equation?
2.
S = Synthesis (Combination)
2 or more reactants and only 1 product
D=
SR =
DR =
C=
3.
Why do chemical equations need to be balanced? _______________________________________________________
_______________________________________________________________________________________________
4.
5.
Balance the following equations and identify the type of reaction occurring:
a. ____Zn + ____HCl ---> _____ZnCl2 + ____H2
type ___________________________
b. _____KClO3 ---> _____KCl + ____O2
type ___________________________
c. ____S8 + ____F2 ---> _____SF6
type ___________________________
d. _____Fe + ____O2 ---> ____Fe2O3
type ___________________________
e. ____C2H6 + ____O2 ---> ____CO2 + ____H2O
type ___________________________
f. _____MgO  _____Mg + _____O2
type ___________________________
For each of the following; predict the products, balance and identify the type of reaction occurring.
_____a. Fe + HCl  (assume that the iron ion formed has a +3 charge)
type ___________________________
_____b. Ca(OH)2 + HCl 
type ___________________________
_____c. NaI + Br2 
type ___________________________
_____d. Pb(NO3)2 + CuSO4 
type ___________________________
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