Review Sheet for Chemistry* First Semester Final
Refer to your class notes, worksheets, and the textbook to complete this review sheet. Study early so that
you will have time to ask questions about what you don’t understand.
(* The study of matter and the changes it undergoes)
Objective 1:

Students will be able to apply common safety rules.
Know how to handle chemicals and laboratory equipment safely.
Answers May Vary: Wear goggles, tie back long hair, never leave a Bunsen burner unattended, hot glass
looks the same as cold glass , do not put solid materials down the sink, always follow lab
and/or teacher’s directions
Measurement
Objective 1:
Students will be able to gather, interpret, and analyze data to draw conclusions.
 Significant Figures
How many significant figures are in the following measurements?
0.012 g _______2______
1,003 mL _____4_______
0.12300 g ______5______
7.54 x 10–3 m ____3_________
5,000 cm3 _____1_______
5,000.00 nm ____6______

Read a measuring device (ruler, balance, graduated cylinder) to the correct number of significant
figures.
3.58 cm
18.0 mL
-1-

Common SI Units of Measurement
Measurement
Unit
Mass
Kg or ___grams____
Volume
Liters, milliliters,
centimeters cubed
Length
Meters, centimeters,
millimeters
Temperature

K or __Celsius______
Convert from one metric unit to another metric unit
100 meters = ____10,000____cm
230 mL = ____0.23____ Liters
0.062 kilograms = __62,000___mg
7 cm = _0.00007_______ km
Matter: Anything that takes up space and has mass
Objective 2:
Objective 3:
Students will be able to identify physical and chemical properties.
Students will be able to differentiate between physical and chemical changes.
 Physical Changes and Chemical Changes
Define each. How can you tell the difference between the two?
Physical change—a change in the substance that does not alter the chemical identity. Changes include
state changes, shape changes. These changes can be reversible.
Chemical change—a change in which the chemical properties of a substance are altered. Chemical
changes are not reversible. A new substance is formed because of chemical changes.
Classify the following as physical or chemical changes:
a. spoiling of milk ___chemical________________
b. bending wire ____physical_________________
c. cutting paper ___physical__________________
d. rusting of a nail ____chemical_______________
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Objective 1: Students will be able to define matter, differentiate between its different states, and
understand how it remains constant with a system.
 Put the following into a graphic organizer/flowchart and define each:
 Matter—anything that has mass and takes up space
 Pure substances–elements and compounds—can not be separated further without chemical means
 Mixtures – homogeneous (solutions) (same throughout) and heterogeneous (different throughout)
Matter
Mixtures
Homogeneous
Pure Substances
Heterogeneous
Elements Compounds
Identify the following as pure substances, homogeneous mixtures or heterogeneous mixtures:
a. copper __pure substance_______
b. sweetened tea _homogeneous mixture____
c. sand and water _heterogeneous mixture___
d. calcium carbonate (CaCO3) __pure substance________
 Sketch particles in the three states of matter. How close are the particles and how much do they
move?
Solid
Liquid
ooo
ooo
Little movement
o o o
o o o o
More movement
Gas
o
o
o
o
o
o
o
o
o
o
o
o Most movement
State of Mater
Solid
Shape
Definite shape
Volume
Definite volume
Liquid
No definite shape
Definite volume
Gas
No definite shape
No definite volume

State the Law of Conservation of Mass
Matter (mass) cannot be created or destroyed during a physical or chemical change.
 Calculate the number of grams of oxygen needed based upon the Law of Conservation of Mass.
4g H2 + 32 g O2 → 36g H2O

Density
Define density.
Amount of matter in a given amount of volume
-3-
d = mass/volume
Calculate density.
If a given object has a mass 100.0 g and a volume of 25.0 mL. What is this object’s density?
100.0 g/25.0 mL = 4.00 g/mL
Given the following 3 rings:
Ring 1
Ring 2
Ring 3
And given that each ring is made out of pure gold, answer the following questions:
What can you say about the relationship between the mass and volume as you compare the 3 rings? As mass increases
so does volume in the rame ratio.
Which ring would have the greatest mass? __Ring 1__________________________________________________
Which ring would have the greatest density? ___All the rings would have the same density____________________
Atom
Objective 3:

Students will be able to describe isotopes.
Define the term isotope. Atoms of the same element (same number of protons) but have different
number of neutrons, thus differing in their mass number.
 Fill in the following for this Carbon–14 isotope, 146 C
 Atomic number = __6__, Mass number = 14___,
 # of protons = __6__, # of electrons = __6__, # of neutrons = __8__.
 Atomic Masses: What is the difference between the mass number for Carbon–14 and carbon’s
atomic mass of 12.011 amu? Carbon-14 has 2 more neutrons (8) than the most common form of
carbon (which has 6 neutrons). The atomic mass is the weighted average of all of the isotopes for
an element.
 Calculate the atomic mass of lithium is one isotope has a mass of 6.0151 amu and a percent
abundance of 7.59% and a second isotope has a mass of 7.0600 amu and a percent abundance of
92.41%.
6.0151 x 0.0759 = 0.4565 amu
0.4565 + 6.5242 = 6.9807 amu is atomic mass of lithium
7.0600 x 0.9241 = 6.5242 amu
Objective 4:
Students will be able to describe the history of our understanding of the atom.
 Atomic Models:
Philosophers: Democritus (believed in atoms) and Aristotle (didn’t believe in atoms)
Scientists: What was the contribution of each one’s atomic model? Draw a model of each.
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
John Dalton—atoms were tiny solid spheres, like marbles.
List the four postulates of Dalton’s Atomic Theory:
1. Each element is composed of small particles called atoms. 2. All atoms of an element are identical, but different for
different elements. 3. Atoms are neither created nor destroyed in any chemical reaction. 4. A compound always has
the same relative numbers and kinds of atoms
 J.J. Thompson—cathode ray experiments, discovered electrons; negative charges dispersed
uniformly throughout the positively charged interior, like chocolate chip cookie dough or plum
pudding.

Ernest Rutherford—gold foil experiment, determined there must be a nucleus composed of the
atom’s positive charge and most of its mass concentrated in the center of the atom. Electrons
circle around the outside of the nucleus. Planetary model of atom

Niels Bohr—based on the hydrogen atom, quantized energy levels for electrons in atoms.
Didn’t work well for atoms with more than 1 electron

Quantum mechanical model (Werner, Heisenberg, Schrodinger):
 Energy levels (n=1, 2, 3, 4,…) – represented by periods on the periodic table
 Sublevels: (s, p, d, f) – represented by blocks on the periodic table
 Orbitals – region of space where up to 2 electrons may be found
Objective 1: Students will be able to describe the basic structure of the atom.
 Characteristics of subatomic particles
Particle
Mass
Charge
Location in atom
Proton
~1amu
+1
In nucleus
Neutron
~1amu
Neutral
In nucleus
Electron
Virtually 0
-1
Orbiting the nucleus
-5-
Nucleus with:
Entire
structure
In the above atom of carbon, identify the proton, neutron, electron, and the nucleus.
How many electrons are there? ______6_______ Given that this diagram is a carbon atom, how many
protons are there? ___6________ Although you cannot see all of the neutrons, how would you determine
how many neutrons there are in this atom of C-12? I would subtract the number of protons from the
mass number which is 12 – 6 = 6 neutrons.
Objective 2:

Students will be able to differentiate between ions and neutral atoms.
Atoms and Ions
How do atoms and ions differ? Atoms have a neutral charge because they have the same
number of protons and electrons. The number of protons and electrons differ in ions. An ion
must have either gained or lost an electron(s).
Element
Iron
# of
protons
# of
neutrons
# of
electrons
Atomic
number
Mass
number
Hyphen
notation
26
30
24
26
56
Iron-56
Nuclear
symbol
Atom or
Ion?
56Fe+2
ion
26
Iodine
53
74
53
53
127
Iodine-127
127I
53
Atom
32S-2
Sulfur
16
16
18
16
-6-
32
Sulfur-32
16
Ion
Objective 5: Students will be able to describe the changes that occur in the nucleus of an atom
undergoing radioactive decay.

Alpha, Beta, Gamma Decay
Type of Decay
Alpha
Nuclear Symbol
Penetrating Ability
Stopped By
Low
Paper, Clothing, Skin
Medium
Plastics, Thick Metal
Foil
High
Lead and Concrete
Beta

Gamma

If you have a focused beam of radioactive emission channeled through a positive and negatively
charged field, what type of emission would be attracted to:
o The Postive Plate
___________Beta_______________
o The Negative Plate
___________Alpha______________
o Not Effective by Electrical Field
___________Gamma____________
In the diagram to the left, draw and label the path of an alpha
particle, beta particle, and gamma ray.
Beta particle
Gamma radiation
Alpha particle

Write a balanced nuclear equations for the following equations:
o Alpha Decay of Uranium-238
238
92U
→
234 Th
90
+
4 He
2
o Beta Decay of Carbon-14
14 C
6
→
0
–1
e +
14 N
7
-7-
Objective 1:
Students will be able to describe the arrangement of electrons according to the
quantum mechanical model of the atom.
 Electron Configurations

What element has the configuration [Ne]3s23p1? __Al_

What does the 3 mean in 3s2 ? Energy Level 3

What does the s mean? Sublevel s

What does the 2 mean? Number of electrons

How many valence electrons will an atom of this element have? 3

How many electrons will an atom of this element lose to form an ion? Why? This atom will lose 3
electrons. It is easier for the element to lose 3 rather than gain 5 in order to satisfy the Octet rule
 Complete the following table:
SubLevel
s
p
d
f
Number of Electrons
2
6
10
14
Number of Orbials
1
3
5
7
 Write out the electron distribution for the element Aluminum according to Hund’s rule. The
1s2 sublevel is done for you.
1s2 _↑ ↓_
_↑ ↓_
2s2 _____
_  _
2p6 _____ _____ _____
_  _ _  _ _  _
3s2 _____
_  _
3p6 _____ _____ _____
_  _ ____ ____
 Write the electron configuration for the following elements:
Calcium
1s22s22p63s23p64s2
Iodine
1s22s22p63s23p64s23d104p65s24d105p5
Vandium
1s22s22p63s23p64s23d3
-8-
 Emission (or bright-line) Spectrums
 What is needed for an electron to “jump” to a higher energy level? A Specific amount, or Quanta, of
energy
 What happens when an “excited” electron falls back to its ground state? The energy required for the
electron to “jump” to the higher energy level is released and can be observed as photons of light
which if seen through a prism is observed as an emission (line) spectrum.
 What does an emission spectrum allow one to do? Using the measure of the wavelength of the
photon of light, one can pinpoint which element is releasing the energy in order for its electrons to
return to ground state.
 Draw the Bohr Model for Chlorine atom below and give the number of valence electrons.
7 valance electrons
Using the bright-line (or emission) spectra above, identify the two elements making up the unknown spectra.
Hydrogen and Helium
Objective 1:
Objective 2:
Students should be able to summarize the periodic law and explain how relates to
physical and chemical properties.
Students will be able to use the periodic table as a reference tool.
Periodic Table
 Locate or define parts of the periodic table:
 Groups Columns

Periods Rows

Transition metals (d & f blocks) B Groups vs. Representative Elements (s & p blocks) A Groups
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
Alkali metals (IA, 1A, 1), Alkaline Earth metals (IIA, 2A, 2), Halogens (VIIA, 7A, 17), Noble Gases
(VIIIA, 8A, 0, 18)
A
B
C
D
On the table above: (See pages 165 through pages 173 in your textbook for more complete answers.)
1. Label the A groups (i.e. IA, IIA, IIIA, etc.)
2. Label the Alkali Metal family. (Red block) This is also known as group __IA______. All elements in this
family have __1_____valance electron. These elements form an ion with a ____+1_____charge. These
elements have their __s______sublevel partially filled.
3. Label the Alkaline-Earth Metal family (Purple Block). This is also known as group __IIA___. All
elements in this family have ___2____valance electrons. These elements form an ion with a
__+2________charge. These elements have their ___s____sublevel completely filled.
4. Color the transition metals blue (Blue Block). These elements make up the _d__block.
5. Label the Halogen family (Green Block). This is also known as group __VIIA____. All elements in this
family have _7____valence electrons. These elements form an ion with a ____–1_____ charge. These
elements have their ___p____sublevel partially filled.
6. Label the Noble Gas family (Orange Block) . This is also known as group __VIIIA____. All elements in
this family have __8___valence electrons. These elements have their __p___sublevel completely filled.
7. Elements found in the f block of the periodic table are called _inner transition____________elements.
- 10 -
Periodic Trends
 Periodic Trends: Increasing or Decreasing from top to bottom or left to right?
Electronegativity
Top to Bottom in a Group
Decreases
Left to Right across a Period
Increases
Ionization energy
Decreases
Increases
Atomic size
Increases
Decreases
Using the periodic table on page 10 of this study guide, answer the following questions:
1. Which element stands alone in its family? ____A (hydrogen)__________
2. Which element has a larger atomic radius A or C? ____C_________
3. Which element has a larger atomic radius C or D? ____C_________
4. Which element has a higher electronegativity? A or B? ___B_________
5. Which element has a higher ionization energy? C or D? ____D_______
 Elements in the same ___Group__ have similar physical and chemical characteristics because the
(group, period)
they have the same number of _____Valence Electrons___.
(atoms, protons, neutrons, electrons, valence electrons)
 Draw a electron dot diagram (or Lewis Dot structure) for Be and for N
correct number of valence electrons.
showing the
 From their positions on the periodic table, what charges would the ions of Be and N have?
Gains or loses electrons?
Be
Loses
Symbol for ion
Be+2
Gains or loses electrons?
N
- 11 -
Gains
Symbol for ion
N-3
 Properties of Metals vs. Nonmetals vs. Metalloids
Metals
Nonmetals
Yes
No
Luster
Metalloids
Maybe
Malleable vs. Brittle
Malleable
Brittle
Maybe either
Conducts electricity
& heat
Typical state(s) at
room temperature
Yes
No
Maybe
Solid
Solid, Liquid, or gas
Solids
Ionic vs. Molecular Compounds
Objective 1:
Objective 2:
Students will be able to describe characteristics of a molecular compound and how a
covalent bond forms at an atomic level.
Students will be able to describe characteristics of an ionic compound and how an
ionic bond forms at an atomic level.

Ionic bonds are formed when a __metal (cation)__ and a __nonmetal (anion)____ combine.

Metals lose electrons and form ____cations____ while nonmetals gain and electrons form
___anions__.

Molecular compounds form when a ____nonmetal____ and a _____nonmetal__ combine as they
share electrons.
 Identify the following pairs of atoms as potentially forming an ionic or molecular compound:
Mg and Cl ______ionic__
Ag and S ___ionic_____
I and F ____molecular_____ P and Cl __molecular_____
K and Br ____ionic_______ Sn and O __ionic_________
Covalent Bonding in Molecules
 Draw Lewis Structures (dot diagrams) for N2, HCl, H2S, CH2Cl2, and O3.
 Use Lewis Structures to predict molecular shapes and polarity of molecules.
 Identify shapes of the molecules as: Linear, Bent, Pyramidal, Trigonal Planar, or Tetrahedral.
 Use electronegativity values to determine if the individual bonds in the molecules above are polar.
 Look at the polarity of the bonds and the symmetry of the molecules above to determine if the
molecules are polar (if one side of the molecule will be more negative than another).
- 12 -
H
2.1
Li
1.0
Na
0.9
K
0.8
He
Be
1.5
Mg
1.2
Ca
1.0
Lewis
Structures
& Total #
of Valence
Electrons
B
2.0
Al
1.5
Ga
1.6
C
2.5
Si
1.8
Ge
1.8
N
3.0
P
2.1
As
2.0
O
3.5
S
2.5
Se
2.4
F
4.0
Cl
3.0
Br
2.8
Ne
Ar
Kr
Electronegativity Type of Bond
Difference (x)
Non–polar
0.0  x  0.4
covalent
0.4 < x < 2.0
Polar covalent
2.0  x
Ionic
N2
HCl
H2 S
Total # of
Valence
electrons = 10
Total # of
Valence
electrons = 8
Total # of
Valence
electrons = 8
Contains
polar
bonds?
Is a polar
molecule?
O3
Total # of
Valence
electrons = 18
Tetrahedral (4
single bonds)
Bent (1 lone pair
on central atom)
C-Cl (polar)
C-H (nonpolar)
Electronegativity
difference = 0
Non polar
Polar
Polar
or
Structural
Formula
Shape of
Molecule
CH2Cl2
Total # of
Valence
electrons = 20
Linear (only 2
atoms)
Bent (2 lone
pairs on central
atom)
Electronegativity Electronegativity Electronegativity
difference = 0
difference = .9
difference = 0.4
Non- polar
Polar
Non- Polar
Non- polar
(perfectly
symmetrical)
Linear (only two
atoms)
Polar
Polar
 How do you know if a bond is polar or nonpolar?
You find if a specific bond is polar or nonpolar by taking the Electronegativity difference between the
two atoms, and use the chart at the top to see which category it fits into
 How do you know if a molecule is polar or nonpolar?
You look at the overall shape of the molecule in 3D. If it is perfectly symmetrical (all bonds around
central atom are the SAME), it is nonpolar. If the “things” around the central atom are asymmetrical,
it is polar.
- 13 -
Properties of Ionic and Molecular Compounds
Molecular Compounds
Ionic Compounds
Combination of elements involved
(metals? nonmetals?)
Nonmetals (maybe
semimetals)
Metal + nonmetal or
Metal + polyatomic ion
How is bond formed?
Electrons are shared
between atoms
Electrons are transferred
between atoms
Typical state(s) at room temperature
Liquid or Gas
Solid
Melting and boiling points
(relatively high or low?)
Low Melting & Boiling
points
High Melting & Boiling
points
Conduct electricity if dissolved in
water?
NO
YES
Naming Molecular and Ionic Compounds
Objective 3:

Students will be able to name and write formulas for compounds.
Naming molecular compounds
Name: N2O: _____Dinitrogen monoxide__ and NO2 _____Nitrogen Dioxide_______
 Naming Ionic Compounds
Name: Li2O _______Lithium Oxide__ and (NH4)2SO4 ___Ammonium Sulfate___
Name: FeO ___Iron (II) Oxide____ and Sn3(PO4)4 ____Tin (IV) Phosphate_____
Name: NaHCO3 ___Sodium bicarbonate or Sodium hydrogen carbonate and
CuCl2 ____Copper (II) chloride______
Formulas of Molecular and Ionic Compounds

Write formulas for the following compounds:
Water _________H2O____________________ silicon dioxide __SiO2______________
Phosphorous trihydride _____PH3______________ dioxygen difluoride __O2F2_______________
Lead (II) hydroxide ____Pb(OH)2______________ chromium (III) sulfate _______Cr2(SO4)3______
- 14 -

Write formulas for the following ionic compounds:
Ba2+ with OH– _Ba(OH)2___
iron (III) sulfide __Fe2S3___________
Na+ with OH– ___NaOH____ NH4+ with PO43– __(NH4)3PO4__ magnesium oxide___MgO______
Chemical Reactions
Objective 1:
Objective 2:
Students will be able to identify types of reactions.
Students will be able to apply the Law of Conservation of Matter to balance equations.
 Define what is meant by the term chemical reaction.
A reaction in which a chemical change has occurred, forming new products.
 In the following chemical equation, identify the reactants and the products.
3Ba(C2H3O2)2(aq) + 2Na3PO4(aq)  Ba3(PO4)2(s) + 6NaC2H3O2(aq)
REACTANTS
PRODUCTS
 In the above chemical equation, what do the symbols (aq) and (s) stand for? What would the symbols
(l) and (g) stand for in a chemical equation? (aq)=aqueous, (s)=solid, (g)=gas, (l)=liquid
 Chemical reactions can often be classified as one of five types. Write the general form for each type
of reaction.
 Direct Combination (or synthesis):
A + B  AB
 Decomposition
 Single-Replacement
AB – A + B
A + BC  B + AC
 Double-Replacement
AB + CD  CB + AD
hydrocarbon + O2  CO2 + H2O
 Combustion
 Using the five types of reactions listed above, classify the following reactions:
2Na(s) + Br2(l)  2NaBr(s)
________Combination_____________________
heat
CH4(g) + 2O2(g) 
CO2(g) + 2H2O(g)
_______Combustion_______________________________
K2CrO4(aq) + Ba(NO3)2(aq)  BaCrO4(s) + 2KNO3(aq) ___Double replacement__________
electric
2H2O(l) 
2H2(g) + O2(g
_____________Decomposition_____________
current
- 15 -
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
H2(g) + O2(g)  2H2O(l)
____________Single replacement___________
_____________Combination________________
heat
2C8H18(l) + 25O2(g)  16CO2(g) + 18H2O(g) __________Combustion_________________
Predicting Products
 Using the activity series, will the following reactions occur?
a. 2Al (s) + Fe2O3 (aq)  Al2O3 (aq) + 2Fe (s)
y or n ______Y____________
b. Zn (s) + 2HNO3 (aq)  Zn(NO3)2 (aq) + H2 (g) y or n _______Y___________
c. Pb (s) + MgO (aq)  PbO (aq) + Mg (s)
y or n _____________N________
 For the above three equations, write out the correct word equation. The first one is started for
you.
a. Solid Aluminum reacts with Iron (III) Oxide to yield aqueous Aluminum Oxide plus solid Iron
b. _Solid Zinc reacts with Nitric Acid to yield aqueous Zinc Nitrate plus Hydrogen gas
c. _Solid Lead reacts with aqueous Magnesium Oxide to yield aqueous Lead (II) Oxide plus solid
Magnesium
 Predict the products of the following reactions. Using the solubility table, predict if there will
be a precipitate. If there is a precipitate, what is it?
Reaction
Products
Precipitate (Y/N)
What is the
Precipitate?
AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
Y
AgCl
2(NH4)3PO4 (aq) +3CaCl2 (aq)  6NH4Cl (aq) + Ca3(PO4)2 (s)
Y
Ca3(PO4)2
NH4NO3 (aq) + NaCl (aq)  No Reaction (both products are aqueous) N
- 16 -
Objective 2:
Students will be able to apply the Law of Conservation of Matter to balance equations.
 When balancing equations it it necessary to have the same kind and number of __atoms______
on each side of the equation.
 Balance the following chemical equations and indicate what type of reaction it represents:
1. __4___Fe(s) + _6____O2(g) 
__2___Fe2O3(s)
Type: ___Combination_____________
2. __3__Fe(s) + ___2__H2SO4(aq)  _____Fe3(SO4)2(aq) + ___2__H2(g) Type: _Single replacement_
3. _____Zn(s) + ___2__HNO3(aq)  _____Zn(NO3)2(aq) + _____H2(g) Type: __Single replacement_
4. _____Pb(NO3)2 (aq) + _____Na2CO3(aq)  ____PbCO3(s) + __2___NaNO3(aq)
Type: ______Double replacement______________________________
5. ___2__C2H2(g) + __5__O2(g)  __4__CO2(g) + __2__H2O(g) Type: __Combustion____________
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