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CHEMISTRY Final Review
Name_________________________
Chemistry is defined as the study of the
it undergoes.
of matter and the
Matter is defined as anything that has
and takes up
Mass is a measure of the amount of
an object contains.
Matter exists in three states:
1)
which has a definite shape and definite volume.
2)
which has an indefinite shape but definite volume.
3)
which has an indefinite shape and indefinite volume.
.
Matter can be classified as either:
or


which cannot be separated into simpler substances
by physical methods.
which can be separated into simpler substances by
physical methods such as filtering, settling, evaporating, boiling, etc..
Pure Substances can be classified as either:

which are the simplest forms of matter and act as
the building blocks for all other substances. The smallest particles with
all the properties of this material are called
. EXAMPLE:

which are chemical combinations of two or more
of the building blocks and have properties much different than the building
blocks from which they are formed. The smallest particles with all the
properties of this material are called
. EXAMPLE:
Mixtures are physical blends of materials which can be classified as either:

___________________ which means that the composition is uniform throughout. All samples of
this type of mixture would have the same ratio of components. A mixture of this type is also
known as a
. EXAMPLE:

__________________ which means that the composition is not uniform. Every sample of this
type of mixture would have a different ratio of components. EXAMPLE:
Identify the following as homogeneous mixture (HOM), heterogeneous mixture (HET), cmpd (C), or elem (E).
A. Copper __________
B. Air __________
C. Sugar _________
D. Salt water _________
E. Water ________
F. Apple juice _______
Properties of Matter
All matter has two types of properties:

which are descriptive qualities or conditions of a substance
which can be measured without changing the composition of the substance.
EXAMPLES:

which are a measure of a substance’s ability to change its
composition or its identity. EXAMPLES:
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Identify the following as Physical properties ( P ) or chemical properties ( C ).
A. Color ________
B. Size _________
C. Flammability __________
D. Reactivity _________
E. Shape _________
F. Explosiveness __________
Changes of Matter
All matter can undergo two types of changes:
which are typically a change in a substance’s appearance

but not its identity.
EXAMPLES:
and identity.
EXAMPLES:

Indicators of a chemical change:
3.
which result in the change of a substance’s composition
________________
1.
4.
2.
5.
Identify each of the following as either a physical change ( P ) or chemical change ( C ).
A. Iron Rusts _______
B. Alcohol evaporates ______
C. Milk sours ____
D. Coal burns _____
E. Ice melts ______
F. Sugar dissolves in water_____
The changes that matter undergoes can be written in the form of a chemical equation which is a short
hand representation of the reaction using chemical and other symbols.
Define the following symbols used in writing a chemical equation:
“+”
“”
(s)
(g)
(aq)
“Δ” ______________
(l) _____________
A chemical reaction that releases heat is said to be: ______________________
A chemical reaction that absorbs heat is said to be: ______________________
The scientific method is a systematic way of solving problems.
List and define the steps in the scientific method.
1.________________________________________________________________________________
2.________________________________________________________________________________
3.________________________________________________________________________________
4.________________________________________________________________________________
5.________________________________________________________________________________
Chapter 2: Scientific Measurement
Information (data) which is collected about a substance or a reaction can be of two types:
- __________________ which is information that is descriptive in nature and uses relative terms in
its descriptions. For example: tall/short, high/low, old/young, large/small
- __________________ which is information that is specific in nature and uses numbers in its
descriptions. For example: 29 oC, 34.6 mm, 17 years old
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Indicate whether each of the following descriptions is Qualitative (Qual) of Quantitative (Quan)
A. Hot ________
B. 35.8 ft __________
C. Soft __________
D. 90 g _________
E. 2.76 hrs __________
F. Dark __________
Different tools are used to collect data in four major categories (mass, length, volume, temperature).
Identify which type of data (Mass, Length, Volume, Temp) can be collected using each of the
following tools
A. 12 inch ruler _________
B. Thermometer __________
C. Triple-beam balance _________
D. Meter Stick _________
D. Graduated Cylinder _________
E. Measuring Cup _________
The quality of measurements can be evaluated in two ways:
- ________________ which is a measure of how close an experimental answer is to a true answer
- ________________ which is a measure of how repeatable a series of measurements are.
Metric System
Base Units:
Length = __________
m
Mass = __________
g
Volume = ____________
L
Measuring tools:
Perfixes:
Equalities:
Kilo =
(K)
A)
B)
C)
D)
deci =
(d)
1m =
1m =
1m =
1m =
Metric Conversions
DRUL Rule:
D R U
centi =
(c)
Km
dm
cm
mm
E)
F)
G)
H)
K (H) (D) b d c m
L
1Km =
1dm =
1cm =
1mm =
milli =
(m)
m
m
m
m
(King Henry Drinks basically delicious chocolate milk.)
If going down to a smaller unit, move the decimal to the right.
If going up to a larger unit, move the decimal to the left.
A) 35.3 cm = ____________ km
B) 14,000 dm = __________ mm
C) 0.0007 m = ___________ cm
D)
E)
F)
0.077 mg = __________ g
22.7 L = _____________ mL
92.1 dL = ____________ kL
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G)
H)
I)
140 L = __________ dL
25.9 kg = ___________ mg
7200 g =
Kg
Chapter 4: Atomic Structure
The idea of an indivisible particle in all matter, the “atomus”, was first suggested in the 4 th century B.C. by
the philosopher, ________________.
This concept was lost for many centuries until the work of several scientists made key discoveries and
presented increasingly sophisticated models for the atom.
Match the scientist with his key contribution to the model of the atom. Some may be used more
than once.
1. Dalton
2. Thomson
3. Rutherford
4. Bohr
A. Proposed the “Plum-pudding” Model________D. Stated electrons must travel in defined orbits ________
B. Proposed the 1st Atomic Theory________
C. Discovered the nucleus________
E. Discovered the electron________
F. Determined that most of the atom is empty space________
Atoms
The atom consists of three major, subatomic particles. They are:

______________ (symbol =
) These are negatively charged particles which are found
orbiting around the atom’s nucleus.

______________ (symbol =

______________ (symbol =
) These are positively charged particles which are located
in the nucleus of the atom.
) These are neutral particles located in the nucleus of the atom.
Atomic Structure
Atoms are composed of mostly
.
_________________________ tells the number of protons in the nucleus of an atom.
The number of protons in an atom of an element never changes and the # p+ identifies the element.
___________________________ tells the number of protons + neutrons in the nucleus of an atom of an element.
The number of neutrons is found by subtracting the ______________________ from the ____________________.
The number of electrons in a neutral atom is equal to the number of ___________________.
The mass of one proton or one neutron is equal to
(atomic mass unit).
Isotopes - atoms of the same element with the same # of protons but a different # of
(same Atomic #, different Mass # and different # no, same #p+ and e-)
is the number of protons + neutrons for one atom of an element.
_________________________ is the weighted average of all the naturally occurring isotopes of an element.
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Isotopes
Symbols of isotopes
X – 10
symbol
mass #
10
atomic#
5
X
OR
mass #
To calculate the atomic mass we multiply the mass of each isotope by its abundance
and then add the masses together to get the atomic mass.
EXAMPLE: Chlorine consists of two isotopes 75% Cl - 35 & 25% Cl – 37.
35 amu x .75 =
26.25 amu
37 amu x .25 = +
9.25 amu
Atomic mass =
35.50 amu
Solve the following problem
Silver has two naturally occurring isotopes:
Calculate its average atomic mass.
Silver -107 (52%) & Silver -108 (48%).
Ions - atoms that have lost or gained electrons and therefore are
(same Atomic #, same Mass # and same # no, different # p+ and e-)
Positively charged ions have
Negatively charged ions have
electrons are called
electrons are called
(107.48 amu)
particles.
.
.
Compare an oxygen atom (O) to an oxide ion (O-2).
# p+ =
# e- =
#n =
What would be the ionic charge on atom, X, with 18 protons, 40 neutrons, & 16 electrons?
Complete the following table of atoms, isotopes and ions using the periodic table as needed.
Element
Symbol
Atomic #
Mass #
# of p+
Sodium
Ca
80
Zn - 67
Carbon -14
S - 33
Li+1
Fluoride
Barium ion
P-3
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# of e-
# of n0
Density
Mass
Density =
M
or
_______________
D =
________
Volume
V
g
=
M
________
cm3 or mL
D
V
Problems
A) A sample of copper has a mass of 44.5 g and occupies a volume of 5.00 cm 3. What is its density?
(8.9 g/cm3)
B) A sample of a liquid with a density of 4.75 g/ml has a volume of 23.1 mL. What is the mass of the sample?
(17.33g)
C) The following data was collected by a group of students who were trying to determine the density of
a metal bar. Calculate the density and identify the metal.
(7.70 g/cm3, Fe)
length = 10.0 cm
width = 4.0 cm
height = 3.0 cm
Mass of bar = 924.0 g
Known densities:
Aluminum = 2.7 g/cm3
Iron= 7.87 g/cm3
Lead = 11.3 g/cm3
What is the % error of the student’s density determination?
(2.16%)
|Experimental Value – True Value|
% Error =
x 100
True Value
Chapter 5: Electron Configurations & Light
A. Identify the sublevel ( s,p,d,f ) associated with each of the following on the periodic table:
alkali &
alkaline metals
nonmetals
& noble gases
transition
metals
inner transition
metals
B. Which sublevels are present in each energy level? 1 _______ 2 ______ 3 _______ 4 _______
C. How many orbitals and electrons are are in each sublevel?
s
p
d
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f
.
D. Aufbau Principle –
E. Hund’s Rule –
F. Pauli Exclusion Principle –
G. Orbital shapes: s
p
d
H. Write the complete electron configuration for a Magnesium (#12) ion.
_____________________________________________________________
I. What noble gas is the magnesium ion isoelectronic to? _____________________________________________
J. Write the complete electron configuration for Chlorine (Cl) #17.
K. For each of the 2 elements listed below:
a. Give the full / longhand e- configuration
b. Give the orbital notation
c. Give the shorthand e- configuration
K (#19)
Na (#11)
L. ionization energy –
Trend: ↓down a group ___________
→across a period ___________
M. atomic size –
Trend: ↓down a group __________
→across a period ___________
N. electronegativity –
Trend: ↓down a group __________
O. Metals want to
P. Nonmetals want to
→across a period __________
electrons so
atoms are more metallic in nature.
electrons so
(size)
atoms are more nonmetallic in nature.
Q. Why do metals have low ionization energies compared to nonmetals?
R. Why do nonmetals have high electronegativies compared to metals?
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S. Why is a Sulfide ion so much larger than a Magnesium ion even though they are in the same period?
The Periodic Table, the Elements, and Compounds
A) The stair-step divider on the periodic table divides the elements known as the ________ from the
elements known as the _________________.
B) The elements that fall just to either side of the divider have properties of both metals and
nonmetals and are known as ____________________.
C) The rows on the periodic table are also called _______________
D) The columns on the periodic table are called ____________ or _____________
E) The elements in a column on the periodic table have similar chemical and physical properties because
all the members of that column have the same # of outermost electrons known as its __________________
electrons
F) Elements in the last column on the table are very stable and tend to not form compounds with
other elements and are known as the __________________
Writing Chemical Formulas
IONIC Compounds (metal + nonmetal)
- The symbol for the positive ion (metal) is written first
- The symbol for the negative ion (non-metal) is written second.
- The charges on each ion are criss-crossed and used as subscripts for the formula but
the positive and negative signs are never used in the subscripts.
- If the subscripts can be divided by the same whole number they must be reduced
to their lowest whole number ratio (empirical formula)
COVALENT Compounds (nonmetal + nonmetal)
- The symbol for the 1st element in the name is written 1st.
- The symbol for the 2nd element is written second.
- Prefixes are used to determine the number to be used as the subscript for each element
- Subscripts are not reduced
IONIC Compounds with 3 or more elements (metal + polyatomic ion)
- The symbol for the positive ion (metal) is written first
- The symbol for the negative ion (polyatomic ion) is written second.
- The charges on each ion are criss-crossed and used as subscripts for the formula
but the positive and negative signs are never used in the subscripts.
- If a subscript other than “1” is used for the polyatomic ion, parentheses must be used
- If the subscripts can be divided by the same whole number they must be reduced
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Writing Names for Compounds
IONIC Compounds (metal + nonmetal)
- The name of the metal is written 1st.
Group A metals - name of the metal element only is written
Group B metals - name of metal followed by a Roman numeral showing the ionic charge
- The name of the nonmetal is written second but the ending of its name is changed to “ide”
COVALENT Compounds (nonmetal + nonmetal)
- The name of the 1st element is written 1st with a prefix to indicate the number of atoms
- The name of the 2nd element is written 2nd with a prefix to indicate the number of atoms
and the ending of its name is changed to “ide”
IONIC Compounds with 3 or more elements (metal + polyatomic ion)
- The name of the metal is written 1st.
Group A metals - name of the metal element only is written
Group B metals - name of metal followed by a Roman numeral showing the ionic charge
- The name of the polyatomic ion is written second (from the list of ions).
Chemical Names and Formulas
A) The elements to either side of the stair-step divider on the periodic table have properties of both metals
and nonmetals and are known as ________________________ or
.
B) The elements known as the _____________ are to the left of the stair-step divider and elements on the
right of the divider are known as the ___________________________ (in the “triangle”).
C) The horizontal rows on the periodic table are also called _______________.
D) The vertical columns on the periodic table are called _______________ or ___________________.
E) The number above a column on the table tells the # of
F) Elements in the same
e- or
e-.
(col) are more similar than elements in the same
(row).
G) Elements in the last column on the table are very stable and tend to not form compounds with other
elements and are known as the ___________________________
H) Metals tend to lose electrons to form ions with a ____ charge known as __________________.
I) Nonmetals tend to gain electrons to form ions with a ____ charge known as__________________.
J) The chemical formula of a compound tells two important things:
___________________________________________________________________in the smallest unit of a compound.
K) An
unit of a compound.
shows the lowest ratio of atoms in the smallest
L) A binary compound is made up of ____ different elements. The 2nd element in a binary compound
always gets an
ending.
M) A ternary compound is made up of ____ different elements and often contains a
. These ions often have
or
N) The group B elements are also known as the
endings.
.
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O) When naming compounds containing group B metals,
to designate the ionic
of the metal.
P) Molecular compounds are composed of
must be used
and are named using
.
Practice for writing compound names and formulas
1) Magnesium cyanide
__________________
9) Sn(OH)2
___________________________________
2) Mercury (II) bromide
__________________
10) PI3
___________________________________
3) Silicon dioxide
__________________
11) N2O3
____________________________________
4) Antimony (V) oxide
__________________
12) KMnO4
____________________________________
5) Diiodine decachloride
__________________
13) Na2CO3
____________________________________
6) Aluminum sulfide
__________________
14) SF6
____________________________________
7) Strontium nitride
__________________
15) CuCl
____________________________________
8) Lead IV sulfate
__________________
16) CaSO4
____________________________________
The ionic charge of a Group A metal is
and equal to
The ionic charge of a Group A non-metal is
.
and equal to
.
For the following formulas, identify the cation and anion, and number of each present in one compound.
cation #
anion
#
cation #
anion
#
A) NaCl
___
___
C) Ca3(PO4)2
B) MgSO4
___
___
D) Fe(OH)3
Chemical Quantities
% element =
___
______
___
___
_______
___
g element x 100
g compound
A. % Composition
1) Find the percent composition of Barium Nitrate.
2) Calculate the percent composition of Ammonium Sulfate?
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(Ba = 52.5%; N = 10.7%;O = 36.8%)
(N = 21.2%, H = 6.1%, S = 24.2%, O= 48.5%)
B. Grams of an Element in a Sample of a Compound
1) How many grams of Copper are in a 75 gram sample of Copper II Chloride?
(35.8 g)
2) How many grams of Nitrogen are in a 124 g sample of Aluminum Nitrate?
(24.5 g)
C. Empirical Formula (Lowest whole number ratio of elements in a compound)
STEPS:
- Divide the amount of each element by its atomic mass (to get moles)
- Divide all of the answers by the lowest answer
- If needed, multiply to get a whole number
1) Calculate the empirical formula given the following percent composition:
52.8% Sn, 12.4% Fe, 16.0% C and 18.8% N
(Sn2FeC6N6)
2) Find the empirical formula of a compound that is composed of 65.2% Sc and 34.8% O.
D. Mole Problems
(Sc2O3)
1 mol = 6.02 x 1023 particles = Molar Mass ( g/mol )
1) How many moles are in 4.6 x 1018 molecules of Diphosphorous Pentoxide?
(7.6 x 10-5 mol)
2) How many moles are in 200 g of Chromium (III) Chlorate?
3) How many molecules are in 58 g of Magnesium Phosphate?
4) How many grams are in .72 moles of Calcium Acetate?
5) How many hydrogen atoms are in 5 molecules of Ammonium Sulfate?
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(.66 mol)
(1.33 x 1023 molecules)
(113.76 g)
(40 H atoms)
Chapter #12: STOICHIOMETRY
1. If 27.0 g of magnesium react with hydrochloric acid, how many grams of magnesium chloride are
produced?
Mg + HCl  MgCl2 + H2
(78.33 g MgCl2)
2. In the following reaction:
Al + H2SO4  Al2(SO4)3 + H2
What would be the Molar Ratio used to convert from moles of Aluminum to moles of Sulfuric Acid (H 2SO4)?
( 3 mol H2SO4 / 2 mol Al )
Chapter # 16 Solutions
1. Name and distinguish between the 2 components of a solution.
2. Define the following terms: solubility, saturated solution, unsaturated solution, supersaturated solution,
miscible, and immiscible.
3. What is the effect of pressure on the solubility of gases in liquids?
4. Calculate the molarity of a solution that contains .50 g of NaCl dissolved in 100mL of solution.
5. How many moles of KCl are needed to create a 1.3 M solution when mixed with 750mL of solution?
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Chapter # 17: The Flow of Energy
1. Explain the Law of the Conservation of Energy.
2. How do endothermic processes differ from exothermic processes?
3. Draw a reaction progression vs. energy graph to show an endothermic reaction and an exothermic
reaction.
4. What is calorimetry?
5. What is specific heat? What is the specific heat of water? What formula can be used to calculate
specific heat?
6. The temperature of a 95.4 gram piece of copper increases from 25 oC to 48oC when the copper absorbs
849 J of heat. What is the specific heat of copper?
Chapter # 19: Acids and Bases
1. List and describe the 3 Acid-Base definitions. Give examples.
2.
a.
b.
c.
d.
Classify the following compounds as Arrhenius acids or bases
Ca(OH)2
HNO3
KOH
H2SO4
3. Identify each reactant in the following equations as a hydrogen ion donor (acid) or a hydrogen ion
acceptor (base). Also label the conjugate acid- base pairs.
a. HNO3 + H2O  H3O+ + NO3b. NH3 + H2O  NH4+ OH-
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4. Calculate the pH for the following solutions and indicate whether each solution is acidic or basic
a. [OH-] = 1.0 X 10-2 M
b. [H+] = 1.0 X 10-2 M
5. What is a neutralization reaction? Give an example
6. How is it possible to recognize the endpoint of a reaction?
7. What are the hydroxide ion concentrations of solutions with the following pH values?
a. 4.00
b. 8.00
c. 12.00
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