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Chemistry SOL Review Packet
Modified from a review created by: Mrs. Joyce McAlister and Ms. Shonna Crisden, Menchville High School.
http://mville.nn.k12.va.us/science/SHONNA/SOLReview/SOLTitlepage.html
Topic 1: Elements and the Periodic Table
1.1 All matter is made from about 100 different chemical elements. The Periodic Table of the
Elements shows all of the known elements, arranged by increasing atomic number. Each element
has a symbol. The symbol for many of the elements is one capital letter. In two-letter symbols for
elements, the first letter is always an upper case letter, the second one a lower case. The smallest
particle of an element is an atom. Some common elements that are gases are composed of
molecules containing two atoms of the same element. Example: hydrogen H2(g) and oxygen O2(g).
2.1 Atoms are made of three types of subatomic particles: protons, neutrons and electrons. Each
atom has a nucleus in the center, made of protons and neutrons packed tightly together. An
electron cloud surrounds the atomic nucleus. The atomic number for an element is the same as the
number of protons. All atoms of the same element have the same number of protons. A proton has
a positive charge and a relative mass of one. The number of electrons is the same as the number of
protons in a neutral atom. An electron has a negative charge and a relative mass of zero. A neutron
has no charge and a relative mass of one.
3.1 There are only certain regions in the electron cloud where electrons are likely to be found. These
regions are called energy levels. The lowest energy level is closest to the nucleus; the highest energy
level is farthest away from the nucleus. Electrons will occupy the lowest available energy level(s)
before they fill in higher levels. The outermost electrons in an atom are called valence electrons. The
period (row) number on the periodic table corresponds to the highest energy level occupied by the
valence electrons in an element. Elements in the same group (column) on the periodic table have
the same number of valence electrons. All of the group 1 elements have one valence electron and
group two elements have two. Group 13 elements have three valence electrons, group 14 elements
have 4, group 15 have 5 and so on through group 18 elements, which have eight valence electrons.
An ion is an atom that has a charge because it has gained or lost electrons. Positive ions have lost
electrons; negative ions have gained electrons. The amount of charge is equal to the number of
electrons lost or gained.
4.1 The principal energy levels (n) around the nucleus of an atom identify the specific regions
(distances from the nucleus) where electrons are likely to be found. Principal energy levels are
identified by n=1, 2, 3 . . . with n=1 closest to the nucleus. As the value of n increases, so does the
distance from the nucleus. Using the periodic table, the period (row) where an element is found
indicates the number of occupied energy levels for that element. The energy level of the valence
electrons corresponds to the period number (row) where the element is found. Each principal
energy level is divided into sublevels (s,p,d and f). In a given energy level, the s sublevel holds up to 2
electrons, and always fills before the p sublevel, which can hold up to 6 electrons. Electron
configurations indicate the filling order of all of the electrons in an atom. The coefficients represent
the principal energy level, the letters represent the sublevels and the superscripts represent the
number of electrons in the sublevel.
5.1 Going down a group on the Periodic Table, each element has one more principal energy level
filled with electrons than the element above it, so the outer electrons are farther away from the
nucleus. This means the size of the atoms increases going down a group. Therefore the atomic radius
increases going down a group. Going from left to right across a period of the Periodic Table the
valence electrons are all in the same principal energy level, but the number of protons in the nucleus
increases from one element to the next. This means that the nucleus becomes more positively
charged and attracts the electrons more strongly. Therefore, the atomic radius decreases going from
left to right across a period. Diatomic elements, N2, H2, O2, Cl2, I2, F2, and Br2.
6.1 Electronegativity is the ability of an atom in a bond to attract electrons. The electronegativity of
an element can be judged from its position on the periodic table. The electronegativity increases
across a period of the periodic table (because the atomic radius deceases, which means that the
valence electrons are held more tightly by the nucleus). The electronegativity decreases going
down a group (because the valence electrons are further away from more loosely held the nucleus).
7.1 Ionization energy is the energy needed to remove a valence electron from a atom. Ionization
energy increases going from left to right across a period of the periodic table because the atomic
radius decreases, which means that the valence electrons are held more tightly by the nucleus.
Ionization energy decreases going down a group because the valence electrons are further away
from and more loosely held by the nucleus.
8.1 Metals react by losing electrons (oxidation). Group 1 metals (called the alkali metals) are the
most reactive metals, because they have only one valence electron to lose. Activity of metals
decreases going across each period to groups 2 and 13. Groups 3 through 12 contain the transition
elements. They are all metals, but less reactive than those in group 1 and 2. Their oxidation states
cannot be easily predicted, and so given in their names. Example: iron (III) means Fe+3; copper (II)
means Cu+2
The nonmetals gain electrons (reduction) when they react. Group 17 nonmetals (called the
halogens) are the most reactive nonmetals because they have only one electron to gain to get a
stable valence shell with eight electrons. Activity decreases in groups 16 and then 15.
Elements in group 18 (called the noble gases) have all their principle energy levels filled with
electrons and so have little chemical reactivity.
Check Your Understanding
1. An atom with atomic number 48 and mass number 120 contains:
a. 48 protons, 48 electrons, 72 neutrons
c. 120 protons, 48 electrons, and 72 neutrons
b. 72 protons, 48 electrons, and 48 neutrons
d. 72 protons, 72 electrons, and 48 neutrons
2. An element which has a mass number of 23 and has 12 neutrons is the element:
a. Lithium
b. Potassium
c. Magnesium
d. Sodium
3. The nucleus of the atom has
a. a high density
c. a negative charge
b. a low density
d. no charge
4. An ion always contains
a. unequal number of protons and electrons
c. unequal number of protons and neutrons
b. equal number of protons and electrons
d. equal number of protons and neutrons
5. The whole number that is closest to the atomic mass of an atom is the
a. atomic number
c. Avogadro's number
b. mass number
d. number of neutrons
6. The ion with the charge of +1 and the same electron configuration as argon is
a. potassium
b. sodium
c. neon
d. magnesium
7. The tendency to lose electron_________ as we move across a period on the periodic table
a. increases
c. decreases
b. remains the same
d. no trend exists
8. Atomic radii generally increases in size from __________ in the periodic table
a. up a group and left to right across a period
b. down a group and left to right across a period
c. up a group and right to left across a period
d. down a group and right to left across a period
9. The _____ generally have the lowest ionization energy
a. noble gases
b. metalloids
c. nonmetals
d. metals
10. Examine the following outer electron configuration and choose the correct location of the
element it represents in the periodic table: 5s2
a. row 7 column 4
c. row 5 column 2
b. row 4 column 7
d. row 5 column 9
11. The element from question 10 is in
a. period 5, alkali metal
c. period 5, halogen
b. period 4, halogen
d. period 5, transitional metal
12. Sodium and potassium have similar properties because they have the same:
a. atomic radius
c. ionization energy
b. number of valence electrons
d. electronegativity
13. The likeliest charge an atom with 2 valence electrons would develop
a. 2+
b. 6+
c. 2d. 614. The likeliest charge an atom with 6 valence electrons will develop
a. 2+
b. 6+
c. 2d. 6
15. How many electrons are in the highest occupied energy level of a sodium?
a. 1
b. 2
c. 3
d. 4
16. What is the maximum number of electrons that can occupy the 5p sublevel?
a. 2
b. 4
c. 6
d. 8
17. How many orbitals in the 4p sublevel?
a. 1
b. 2
c. 3
d. 4
18. What is the maximum number of electrons that can occupy the 4th energy level?
a. 9
b. 18
c. 20
d. 32
19. Which is the symbol of the element whose electron configuration is 1s2 2s2 2p6 3s2 3p6 4s2
a. Ca
b. Li
c. Rb
d. Ar
20. Elements tend to gain or lose electrons in order to acquire the electron configuration of a:
a. halogen
c. noble gas
b. transition metal
d. nonmetal
Topic 2: Compounds and Bonding
2.1 Atoms of different elements can join together by chemical bonds to form a compound. A
compound has totally different properties from its elements. Chemical formulas show the ratio or
number of atoms of each element in a compound. Example: 2 hydrogen atoms bonded to one
oxygen atom make a water molecule (H2O).
2.2 Subscripts in a chemical formula represent the relative number of each type of atom. The
subscript always follows the symbol for the element. Example: In a water molecule, H2O, there are 2
hydrogen atoms and one oxygen atom. Parentheses are used when a subscript affects a group of
atoms. Example: the formula for magnesium nitrate is written Mg(NO3)2 to show that there is a ratio
of one magnesium atom, 2 nitrogen atoms and 6 oxygen atoms in the compound.
2.3 Elements in groups 1, 2 and 13 (metals) will lose electrons and form positive ions. Elements in
groups 15, 16 and 17 will gain electrons and form negative ions. Ionic compounds are formed by the
attraction between positive and negative ions. The charges must be balanced, resulting in a
compound with no net charge. The nomenclature for binary ionic compounds as directed by
IUPAC, metal first keeping its name, non-metal second with an ending change of “-ide”.
Polyatomic Ions are a group of atoms bonded together that have a charge.
2.4 An ionic bond is formed when a metal (element from group 1 through 13) transfers electrons to
a nonmetal (element from groups 15, 16 or 17). This is because metals form positive ions (by losing
electrons) and nonmetals form negative ions (by gaining electrons), resulting in a strong attraction
between oppositely charged ions. Formulas for ionic compounds are written by balancing the ion
charges. A covalent bond is formed when a nonmetal shares electrons with another nonmetal.
Formulas for covalent molecules can be predicted from the dot diagrams of the combining
elements. Sometimes more than one pair of electrons is shared between atoms. If two pairs are
shared then there is a double covalent bond, three pairs shared is a triple covalent bond. Electron
dot diagrams for elements show the number of valence electrons. Elements will transfer or share
valence electrons in order to have eight valence electrons (octet rule).
2.5 The Lewis dot diagram for a covalent compound shows that each of the atoms in the molecule
has a filled valence level. Each shared pair of electrons in a Lewis diagram represents a covalent
bond. A covalent molecule can also be represented by a structural formula in which each covalent
bond is shown as a line joining two atoms. In other words, a line in a structural formula represents two
electrons. Covalently bonded compounds (molecular) use prefixes to show how many atoms of
each element are present. Prefixes indicate 1 to 5 atoms are mono-, di-, tri-, tetra-, and penta-.
Mono- is dropped when only one atom of the first element is present. Both ionic and covalent binary
names end in “-ide”.
2.6 A covalent bond consists of electrons shared between atoms, but this sharing is not equal unless
the two atoms involved are identical (like two chlorine atoms in a Cl2 molecule), because different
atoms have different electronegativities. The more electronegative atom in a covalent bond will
attract the electrons more strongly and this will result in it having a slight negative charge. The less
electronegative atom will therefore be slightly deficient in electrons and so will have a slight positive
charge. A covalent bond in which the atoms have slight electrical charges is known as a polar
covalent bond. A non-polar covalent bond has equal sharing of electrons. A molecule is polar if it
contains polar bonds and if its shape puts a positive and negative charge at different ends on the
molecule. The shape of a molecule is determined by counting the areas of high electron density in
the molecule. A linear shape results from two atoms bonded together, or from three atoms bonded
together with no unshared electron pairs on the center atom (N2 and CO2). A tetrahedral shape
results from four atoms bonded to a central atom and no unshared electron pairs (CH4). A
pyramidal shape results three atoms bonded to a central atom and one unshared electron pair
(NH3). A bent or angular shape results from two atoms bonded to a central atom and two unshared
electron pairs (H2O).
2.7 Intermolecular forces are attractions resulting from forces between molecules. The strength of
these forces will effect the vapor pressure (the pressure in a closed container that is due molecules
that have evaporated of the liquid), boiling point (a liquid boils when its vapor pressure equals the
pressure of the atmosphere) and surface tension (the stretching force among particles that produces
a liquid “skin” on the surface) of the molecules. A hydrogen bond is an attraction occurring when a
hydrogen atom bonded to a strongly electronegative atom is also attracted to another
electronegative atom, often of a different molecule. A hydrogen bond is the strongest
intermolecular force. London dispersion forces are the weakest intermolecular attraction between
non-polar molecules. Dipole-dipole attraction is the attractive force resulting when polar covalent
molecules line up so that the positive and negative ends are close to each other.
2.8 The Group 1 metals, the alkali metals, form ionic compounds (salts) with halogens in a ratio of
1:1. (Example: LiCl, NaCl, NaI, NaF, KCl), and with group 16 non-metals in a ratio of 2:1 (Example: Li2O,
Na2O, Na2S, K2S) Group 2 metals, the alkaline earth metals, form ionic compounds with halogens in a
ratio of 1:2 (Example: BeCl2, MgCl2, CaCl2, CaF2, SrI2), and with group 16 non-metals in a ratio of 1:1
(Example: CaO, MgO, CaS, MgS). Covalent bonding occurs between atoms that have relatively high
electronegativities (between non-metals). Compounds can be predicted with dot diagrams. Group
18 elements (noble gases) do NOT naturally form compounds.
Check Your Understanding
1. Elements tend to gain or lose electrons in order to acquire the electron configuration of a
a. halogen
b. transition metal
c. noble gas
d. nonmetal
2. A double bond consists of
a. two pairs of shared electrons
c. six shared electrons
b. two shared electrons
d. unshared electrons
3. When one atom is significantly more electronegative than the other one, a covalent bond between
them is
a. nonpolar
b. hydrated
c. polar
d. very unstable
4. The type of chemical bonding in which electrons are transferred from one atom to another is
a. nonpolar covalent
c. ionic
b. polar covalent
d. all of the above
5. An example of a polar covalent molecule would be
a. NaCl
b. HCl
c. H2
d. O2
6. Unequal sharing of electrons is a characteristic of:
a. covalent bonds
c. metallic bonds
b. ionic bonds
d. polar covalent bonds
7. Which of the following molecules does not have a linear shape?
a. O2
b. H2S
c. HI
d. CO2
8. The structural formula for the nitrogen molecule (N2) contains a
a. single bond
b. double bond
c. triple bond
d. no bonds
9. A substance that is made up of molecules that have a partially positive and a partially negative
end is
a. nonpolar covalent
c. ionic
b. polar covalent
d. nonpolar ionic
10. The predicted geometry of NH3 is
a. pyramidal
b. trigonal planar
c. bent
d. linear
11. Choose the set of molecules which are in correct order of increasing polarity.
a. C-Br , C-Cl , C-F , C-H , C-I
c. C-Cl , C-F , C-I , C-Br , C-H
b. C-F , C-Cl , C-H , C-Br , C-I
d. none of the above
12. Which of the following represents a tetrahedral molecule?
a. CH4
b. CaCl2
c. NH3
d. Br2
13. Which of the following correctly matches the names and formulas of both compounds?
a. AlCl3 . aluminum trichloride and N2O4 , nitrogen oxide
b. AlCl3 , aluminum trichloride and N2O4 dinitrogen tetroxide
c. AlCl3, aluminum chloride and N2O4 nitrogen oxide
d. AlCl3 , aluminum chloride and N2O4 dinitrogen tetroxide
14. The name of K2SO4 is:
a. dipotassium sulfate
c. dipotassium sulfite
b. potassium sulfoxide
d. potassium sulfate
15. The formula for Copper (I) Chloride
a. CuCl2
b. Cu2Cl
c. CuCl
d. CuClO2
16. The formula for ammonium hydroxide is:
a. AmOH
b. AnOH
c. NH4OH
d. NH4(OH)2
17. The formula for iron (II ) oxide is :
a. FeO
b. Fe2O
c. FeO2
d. Fe2O3
18. The formula for calcium nitrate is:
a. CaNO3
b. Ca(NO3)3
c. Ca(NO3)2
d. Ca4NO2
19. The formula for ammonium phosphate is:
a. NH4P
b. (NH4)3PO4
20. The formula for sulfur trioxide is:
a. SO3
b. SO
c. AmPO4
d. (NH4)3(PO3)2
c. SO2
d. SO4
Topic 3: Kinetic Theory
3.1 Atoms and molecules are in constant motion. Particles of a gas move fastest; particles in a
liquid move slower and particles in a solid move slowest. There is a direct relationship between
temperature and speed of the particles. When the temperature increases, particles move faster.
There is an inverse relationship between pressure and volume of gas particles. When the pressure
increases, the volume decreases.
3.2 Pressure and temperature both affect the volume that a gas occupies.
 Pressure and volume are inversely related; if pressure increases, volume decreases.
Mathematically the relationship means that PV=k, i.e. that if all the factors remain constant
then the pressure times the volume is also constant.
 Absolute temperature and volume are directly related; if absolute temperature increases,
volume increases when the temperature remains constant. Mathematically this relationship
means that V/T=k, i.e. that if all other factors remain constant then the volume divided by the
temperature is also constant.
3.3 Phase changes that require heat (like melting or boiling) are endothermic. H is positive for an
endothermic change. This means heat goes in. Phase changes that give off heat (like freezing and
condensing) are exothermic. H is negative. This means heat is released. Heating and cooling curves
are also known as phase diagrams.
3.4
Elements form bonds to become more stable. A filled valence configuration (eight s and p
electrons) s2p6 is very stable. Stability is inversely related to potential energy, therefore when atoms
bond they become lower in potential energy. Potential energy is stored energy. Chemical bonds
contain potential energy. Energy is required to break bonds. Breaking bonds is endothermic. Energy is
released when bonds are formed. Forming bonds is exothermic.
3.5 In chemical reactions bonds are broken and new bonds are formed. The energy absorbed in
breaking the bonds is never exactly equal to the energy released when the new bonds are formed.
Therefore, all reactions are accompanied by a change potential energy that can be measured and
is represented by the symbol H. An energy level diagram can also be used represent the energy
change during a chemical reaction. Activation energy is the minimum amount of energy that must
be supplied to a system to start a chemical change. Heat of reaction is the amount of energy
absorbed or released during a chemical change. A catalyst is a substance added to a chemical
reaction to increase the rate and that can be recovered chemically unchanged after the reaction is
complete.
3.6 Dalton's Law states, the total pressure in a gas mixture is the sum of the partial pressures of the
individual components. The partial pressure of any gas in a mixture can be calculated using the
mole fraction of that gas in a mixture.
Ptotal = P1 + P2 + P3 +.....
The folowing mathematical relationship between the pressure, volume and temperature of a
gas is used to describe the behavior of gases:
P1 V1 = P2 V2
T1
T2
3.7 Hvap (Heat of vaporization) is the amount of energy needed for the particles of a substance to
escape from the attractive forces of the other particles and escape from the surface into the gas
phase. The stronger the forces of attraction between particles, the greater the heat of vaporization.
Heat of fusion is the amount of energy released when the particles of a substance solidify. The
stronger the attractive forces between the particles, the greater the heat of fusion. Heat of fusion
(Hfus) is the amount of energy released when the particles of a substance solidify. The stronger the
attractive forces between the particles, the greater the heat of fusion.
3.8 A reaction rate describes how rapidly a chemical change takes place. Reaction rates are
determined experimentally by measuring a change in some physical property such as volume,
temperature, color, mass or pH. There is a direct relationship between temperature and reaction
rate. Reversible reactions reach equilibrium. At equilibrium, the forward and reverse reactions occur
at the same rate. Le Chatelier’s Principle states, when a system at equilibrium is disturbed by applying
stress, the equilibrium position shifts to relieve the stress. Stresses that can change equilibrium include
changes in concentration, temperature or pressure.
Check Your Understanding
1. Equal molar quantities of gases A, B, C, and D are put into an evacuated flask. The pressure of the
sample is measured as 160. kPa. The pressure of gas C is
a. less than 20.0 kPa
c. between 60.0 and 100. kPa
b. between 20.0 and 60.0 kPa
d. greater than 100.0 kPa.
2. A 30. mL sample of oxygen is at a temperature of 66oC. If the temperature is lowered to 33oC at
constant pressure, the volume of the gas will become
a. 15 mL
c. 45 mL
b. 27 mL
d. none of the above.
3. A gas at 50.0 kPa has a volume of 4.0 dm3. If the temperature is held the same and the pressure
on the gas in reduced to 10.0 kPa, the volume of the gas would become
a. 0.80 dm3
b. 4.0 dm3
c. 20. dm3
d. 40. dm3.
4. If 1.0 moles of carbon dioxide gas is introduced into an empty vessel at a pressure of 1.2 atm and
55oC, the volume of the gas would be
a. 22.4 L
b. 32.3 L
c. 2270 L
d. 3270 L.
5. Compute the number of moles of a gas in a 5000. L tank if the gas temperature is 44C and its
pressure is 101.3 kPa.
a. 0.00520 moles
b. 192 moles
c. 1410 moles
d. 18900 moles.
6. Given the reaction at equilibrium:
H2(g) + Cl2(g)  2HCl(g) + energy
The equilibrium will shift to the right when
a) chlorine gas is removed
c) hydrogen chloride gas is added
b) hydrogen gas is removed
d) temperature is decreased
7. In the reaction:
N2(g) + 3H2(g) ---->2NH3(g) + 22.0 kcal
a) the reaction is both endothermic and exothermic
b) the reaction is endothermic
c) the reaction is exothermic
d) the reaction is neither endothermic or exothermic
8. On the energy diagram, the energy of the activated complex is represented by:
a) a
b) b
c) c
d) d
9. If the phases of matter are arranged in order of increasing disorder, the arrangement would be:
a. solid, liguid, gas
b. gas, solid, liquid
c. gas, liquid, solid
d. liquid, solid, gas
10. How many kilojoules of energy are needed to convert 2.5 moles of water from ice to liquid if the
heat of fusion of water is 6.00 kJ/mol?
a. 2.5 kJ
b. 6.00 kJ
c. 15 kJ
d. 30kJ
11. The enthalpy change for melting a solid, such as ice, is called:
a. heat of fusion
c. heat capacity
b. heat of vaporization
d. specific heat
12. At chemical equilibrium, the rates of the forward reaction and reverse reactions are:
a. equal to 0
c. at a maximum
b. equal to each other
d. at a minimum
13. The rate of a chemical reaction normally:
a. increases as reactant concentration increases.
b. is slowed down by a catalyst.
c. decreases as temperature increases.
d. decreases as surface area increases.
14. Activation energy is:
a. the heat released in a reaction.
b. the minimum energy colliding particles must have to react.
c. the energy given off when reactants collide
d. generally very high for a reaction that takes place rapidly.
15. In general, increasing the temperature causes the rate of most chemical reactions to:
a. increase
c. remain the same
b. decrease
d. vary unpredictably
16. The principle that relates changes imposed on equilibrium systems to equilibrium position is:
a. Haber's Law
c. Le Chatelier's Principle
b. the law of chemical equilibrium
d. Avogadro's Principle
Topic 4: The Mole and Stoichiometry
4.1 Atoms and molecules are too small to count. Mole is the unit used to tell how many particles
are in a certain amount of a substance. A mole is 602,000,000,000,000,000,000,000 particles (atoms or
molecules). Expressed in scientific notation, a mole is 6.02 x 1023 particles. Scientific notation is used
to express very small or very large measurements in powers of ten. It expresses quantities by using a
number between one and ten, which is then multiplied by ten to a power to give the quantity its
proper magnitude.
4.2 The sum of the protons and neutrons in an atom is known as the mass number. The number of
neutrons in an atom can be found by subtracting the atomic number from the mass number.
Isotopes are atoms of the same element that have different numbers of neutrons. Some isotopes are
radioactive, many are not.
4.3 The molar mass of a compound is the mass of one mole of the compound. It is found by taking
the sum of the molar masses of the individual elements that make up the compound. The percent
mass of an element in a compound can be determined:
% by mass of element = total mass of element in compound
total mass of compound
X 100
4.4 Molar masses from the periodic table can be used to calculate the number of moles in a
given mass of an element or compound. This is because the masses on the periodic table represent
the number of grams in one mole. The number of moles can also be used to calculate the number
of particles – atoms or molecules. The number of particles can be determined from the mass of a
compound or element.
4.5 Because matter cannot be created or destroyed, the total mass of the products is equal to
the total mass of the reactants in a chemical reaction. Molar masses from the periodic table and
mole ratios from the balanced equation can be used to calculate the mass of a reactant or product.
4.6 Balanced chemical reactions and numbers of moles of each substance can be used to
predict the masses of reactants or products. At STP (standard temperature and pressure, 0 Celsius
and 1 atmosphere OR 760 mmHg OR 101.2 kPa) the volume of 1 mole of any gas is 22.4 liters.
4.7 The number of moles of a gas (n) can be determined if the pressure (P), temperature (T) and
volume (V) of the gas sample are known, using the constant R according to the following equation:
PV = nRT is the ideal gas law constant and has two values depending on the pressure units. They are
R = 8.314 L.Kpa/mol.K and R = 0.0821 L.atm/mol.atm.
4.8 An empirical formula shows the smallest whole number ratio of elements in a compound. Ionic
solids are composed of oppositely charged ions arranged in a regular, repeating, crystal lattice
structure; the empirical formula always gives the ratio of positive to negative ions. Covalent
compounds are often in the form of individual molecules; the empirical formula gives the ratio of
atoms in one molecule. Example: The molecular formula for glucose is C6H12O6; the empirical
formula is CH2O.
Choose the best answer that either answers the question or completes the statement
1. The number 2 X 101 expressed in common numerical expression is:
a) 200
b) 20
c) 2
2
2. The number 5.10 X 10 expressed in common numerical expression is:
a) 501
b) 510
c) 5100
3. The number 300 expressed in scientific notation is
a) 3 X 10 -1
b) 3 X 10 -3
c) 3 X 10 2
4. The number 0.0006 expressed in scientific notation is:
a) 6 X 10 -4
b) 6 X 10 2
c) 6 X 10 -2
d) 0.02
d) .00501
d) 3 X 10 1
d) 6 X 10 4
5. How many moles are in 8.5 x 10 25 molecules of CO?
a. 1.4 x 10 2
b. 7.1 x 10 -3
6. What is the molar mass of CO2?
a. 36.0 g
b. 11.0 g
7. It is possible to convert moles to atoms by:
c. 5.1 x 10 49
d. 8.5 x 10 25
a. multiplying by 6.02 x10 23
c. multiplying by the molar mass
d. dividing by the molar mass
c. 44.0 g
d. 6.02 x 10 23 g
b. dividing by 6.02 x10 23
8. The percentage composition of ammonia (NH3) is:
a. 78.5 % N and 21.5 % H
c. 82.4 % N and 17.6 % H
b. 21.5 % N and 78.5 % H
d. 17.8 % N and 82.2 % H
9. How many molecules of sulfur dioxide are present in 1.60 moles of sulfur dioxide?
a. 9.63 x 10 23
c. 7.62 x 10 1
b. 102.1 x10 1
d. 3.76 x 10 23
10. Find the number of moles in 3.30 g of (NH4)2SO4
a. 132.1
b. 40.0
c. 0.279
11. What is the volume of 2.50 moles of carbon monoxide at STP?
d. 0.0250
a. 0.112 L
b. 3.10 L
c. 56.0 L
d. 8.96 L
12. The volume of 2.00 moles of any gas at STP is :
a. 11.2 L
b. 22.4 L
c. 44.8 L
d. 3.2 L
13. Which contains more atoms?
a. 1.00 mole H2O2
c. 1.00 mol CO
b. 1.00 mol C2H6
d. 1.00 mol H2O
14. The volume of 1.20 mol of Ar at STP is
a. 170 L
b. 9.9L
c. 27 L
d. 26.9 L
15. The percent composition of H2S is
a. 5.9 % H and 94.1 % S
c. 9.42 % H and 4.1 % S
b. 5.1 % H and 94.1 %S
d. 50 % H and 50 % S
16. How many liters of hydrogen react with 2.00 mol of nitrogen at STP in the following reaction
N2 + 3H2  NH3?
a. 3.0 L
b. 22.4 L
c. 67.2L
d. 134 L
17. How many molecules of NO2 are produced when 2.0 x 10 20 molecules of N2O4 are decomposed
according to the equation:
N2O4 (g)  2NO2 (g)
a. 4
b. 1.0 x 10 20
c. 2.0 x10 20
d. 4.0 x 10 20
18. Given the reaction: 2H2O  2H2 + O2, how many moles of H2O would be required to produce
2.5 moles of O2 ?
a. 2.0 mol
b. 2.5 mol
c. 4.0 mol
d. 5.0 mol
19 Given the reaction: CuO + H2  Cu + H2O, how many moles of H2O are produced when 240
grams of CuO react?
a. 1.0 mol
b. 3.0 mol
c. 18 mol
d. 54 mol
20. Gven the reaction: Zn + H2SO4  ZnSO4 + H2, how many grams of H2SO4 are required to produce
1.0 g of H2?
a. 1.0 g
b. 2.0 g
c. 49.1 g
d. 98.0 g
Topic 5: Chemical Reactions
5.1 A chemical reaction is required to change one substance into another by rearranging its atoms.
The only way to form a compound from elements is by a chemical reaction. Example: In a synthesis
reaction (combination reaction), hydrogen gas and oxygen gas react to form water. 2H2 + O2 
2H2O The only way to separate a compound into its elements is by a chemical reaction that breaks
the chemical bonds, forming new substances. Example: water decomposes to form hydrogen and
oxygen gas.
2H2O  2H2 + O2
5.2 A chemical equation is a record of what happens in a chemical reaction. It shows the formulas
of all the reactants on the left hand side of the arrow, and the formulas for all the products on the
right hand side.
5.3 Combustion reactions are exothermic reactions in which oxygen combines with other elements.
One example is the reaction between methane and oxygen in a Bunsen burner:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + energy
5.4 Because matter cannot be created or destroyed, elements must be conserved in a chemical
reaction (Conservation of Mass). There must be the same number of each kind of atom on both sides
of a balanced equation. The only way to balance a chemical equation is by placing coefficients in
front of each substance until each side has the same number of atoms of each element.
5.5 When two or more substances combine to form a single product, the reaction is called a
synthesis reaction. For example, the formation of water from hydrogen and oxygen gases is a
synthesis reaction:
2H2 (g) + O2 (g)  2H2O
In a decomposition reaction, a compound breaks down into two or more simpler substances. For
example, in electrolysis, water is broken down into hydrogen and oxygen gases:
2H2O (l)  2H2 (g) + O2 (g)
5.6 In a single replacement reaction (review single replacement reactions here) one element takes
the place of another in a compound. In a double replacement reaction the positive portions of two
ionic compounds are interchanged.
5.7 In a sample of pure water a very small number of water molecules dissociate, producing equal
concentrations of both hydrogen ions [H+] and hydroxide ions [OH-].
The pH of pure water is 7.
H2O  H+ + OH-
or
2H2O  H3O+ + OH-
5.8 Neutralization reactions result from the reaction of an acid with a base to form a salt and water.
These reactions are usually double replacement reactions. HCl + NaOH  NaCl + HOH.
Check Your Understanding
1. Which of the following types of reactions results in a single product?
a. combination /synthesis
c. single replacement
b. decomposition
d. double replacement
2. The numbers used to balance a chemical equation are called:
a. superscripts
b. subscripts
c. coefficients
d. formula units
3. A chemical equation is balanced when ____________ .
a. the equation shows an equal number of atoms for each element on both sides.
b. at least one substance in each of the three physical states is present.
c. the total number of moles of the reactants equals the moles of the products.
d. none of the above
4. The general form for a double displacement reaction
a. element + compound  element + compound
b. compound  two or more elements or compounds
c. element or compound + element or compound  compound
d. compound + compound  compound + compound
5. In the reaction 2KClO3  2KCl + 3O2 oxygen is a _____________.
a. reactant
b. product
c. coefficient
d. subscript
6. In the equation 2Fe + 3H2O  Fe2O3 + H2, iron is a ________.
a. subscript
b. reactant
c. product
d. coefficient
7. The ionic compound formed in an acid-base neutralization reaction is a(n) :
a. indicator
b. hydroxide
c. salt
d. hydride
8. The products of the neutralization reaction between HNO3 and Ca(OH)2 are:
a. 2CaNO3 + H20
c. CaNO3 + H2O
b. Ca(NO3)2 + H2O
d. Ca(NO3)2 + 2H2O
9. In what ratio do HCl and Mg(OH)2 react through neutralization?
a. 1:1
b. 2:1
c. 1:2
d. 2:2
10. The coefficients needed to balance the equation: Zn + CuSO4 ZnSO4 + Cu are
a. 1,2,2,1
b. 2,1,2,1
c. 1,1,2,2
d. balanced
11. The coefficients needed to balance the equation: H2 + O2 H2O are
a. 2,1,1
b. 1,1,2
c. 2,1,2
d. balanced
12. The coefficients needed to balance the equation: PCl5 PCl3 + Cl2 are
a. 2,2,1
b. 1,2,1
c. 1,1,2
d. balanced
13. The coefficients needed to balance the equation: PbCl2 + Li2SO4  LiCl + PbSO4 are
a. 1,1,2,1
b. 1,2,2,1
c. 2,2,1,1
d. balanced
14. The coefficients needed to balance the equation: Zn + HCl  ZnCl2 + H2 are
a. 1,1,2,1
b. 1,2,1,1
c. 1,1,1,2
d. balanced
15. The coefficients needed to balance the equation: AgNO3 + CaCl2  AgCl + Ca(NO3)2 are
a. 2,1,2,1
b. 1,2,1,2
c. 1,1,2,2
d. balanced
16. The coefficients needed to balance the equation: Zn + FeCl2  ZnCl2 + Fe are
a. 1,3,1,3
b. 3,1,3,1
c. 1,1,3,1
d. balanced
17. The chemical reaction in which hydrogen ions from an acid react with hydroxide ions from a base
to produce water is:
a. ionization
b. titration
c. concentration
d. neutralization
18. The reaction H2CO3  H2O + CO2 is an example of a reaction that is
a. synthesis
c. single replacement
b. decomposition
d. combustion
19. The equation Cu + Hg(NO3)2  Hg + Cu(NO3)2 exemplifies a reaction that is
a. synthesis
c. neutralization
b. single replacement
d. double replacement
20. Which of the following is a neutralization reaction?
a. K + Cl2  KCl
c. Cu(OH)2  CuO + H2O
b. NaOH + NH4Cl  NH4OH + NaCl
d. LiOH + HBr  LiBr + H2O
Topic 6: Solutions
6.1 If a substance contains different types of particles, then it is called a mixture. (Example:
Soapy water is a mixture of many water molecules and a few soap particles). In a heterogeneous
mixture, the different parts can be easily seen (like salt and pepper mixed together). In a
homogeneous mixture the particles are mixed so well that the separate parts cannot be seen (like
salt dissolved in water).
6.2 A solution is a homogeneous mixture because the separate parts of the mixture cannot be
seen. The solvent (usually water) is the part of the solution that is present in largest amount. The solute
is the substance that is dissolved. A saturated solution has all the dissolved solute that it can hold,
and can be identified by undissolved solute particles on the bottom after mixing. An unsaturated
solution can still hold more solute.
6.3 Dissolving is a physical change that involves heat. Dissolving and dissociation can be
represented by an equation. Example: NaCl(s) + heat  Na+(aq) + Cl-(aq)
H is positive. If a solution gets cooler when a solute dissolves, it is an endothermic change and H is
positive, and heat is written to the left of the arrow. If a solution gets warmer when a solute dissolves,
it is an exothermic change and H is negative, and heat is written to the right of the arrow.
6.4 Solutions that contain ions are called electrolytes because they can conduct an electric
current. Therefore, solutions of ionic compounds (salts) in water (aqueous solutions) are electrolytes,
because ionic compounds dissociate as they dissolve. Conductivity is directly related to the number
of ions in the solution. Strong electrolytes are good conductors while weak electrolytes are poor
conductors. Solutions that do not contain ions are called non-electrolytes because they cannot
conduct an electric current.
6.5 The concentration of a solution is the amount of solute contained in a certain volume of
solution. If a solution contains a small amount of solute it is called dilute, and if it contains a large
amount of solute it is called concentrated. In chemistry, concentration is given as molarity, the
number of moles of the solute in one liter of solution and expressed as mol/L or just M.
6.6 The general rule for predicting solubility is "like dissolves like". Water is a polar substance, so it
can dissolve ionic and polar solutes. Oil is non-polar, so oil will not dissolve in water. Oil and water
don’t mix but different oils do because a non-polar solute will dissolve in a non-polar solvent.
6.7 pH (0-14) measures the hydrogen ion concentration in water. Each pH unit involves a tenfold
change in hydrogen ion concentration. The pH number and the [H+] are inversely related because
pH = - log [H+]
o An increase of 1 pH unit means that the hydrogen ion concentration has decreased 10 times.
o A decrease of 1 pH unit means that the hydrogen ion concentration has increased 10 times.
Acids are compounds that increase the concentration of hydrogen ions [H+] when they dissolve in
water. Acid solutions have a pH below 7, taste sour and turn litmus paper red. Bases are compounds
that increase the concentration of hydroxide [OH-] when they dissolve in water. Bases have a pH
greater than 7, taste bitter, feel slippery and turn litmus paper blue.
6.8 Both strong acids and strong bases dissociate completely in water, therefore are strong
electrolytes. In a solution of a strong acid like hydrochloric acid, almost all of the HCl molecules
dissociate according to the following equation:
HCl(aq) --> H+(aq) + Cl-(aq)
Weak acids and weak bases are weak electrolytes. In a solution of a weak acid like acetic acid,
only a few of the CH3COOH molecules dissociate:
CH3COOH <---> H+(aq) + CH3COO-(aq)
Check Your Understanding
1. Molarity is expressed as:
a. moles of solvent/liter of solution
c. moles of solute/mole of solvent
b. moles of solute/liter of solution
d. moles of solute/kilogram of solution
2. A saturated solution:
a. contains the maximum amount of solute that can be dissolved
b. is concentrated
c. is diluted
d. none of the above
3. The molarity of a solution that contains 14 g KOH per 150 mL of solution is:
a. 93 M
b. 1.7M
c. 0.093 M
d. 11 M
4. Another name for a solution is a:
a. heterogeneous mixture
c. compound
b. homogeneous mixture
d. element
5. The dissolved substance in a solution is called the :
a. solute
b. solvent
c. hydrate
d. tincture
6. How many moles of solute are present in 1.25 L of a 0.75M NaNO3 solution?
a. 1.7
b. 0.60
c. 0.75
d. 0.94
7. If salt is dissolved in water, water serves as the:
a. solute
c. dissolved medium
b. solvent
d. none of these
8. What is the molarity of a solution that contains 8 moles of solute in 2L of solution?
a. 4M
b. 8M
c. 6M
d. 0.25M
9. In the reaction: N2(g) + 3H2(g) <----> 3NH3(g) + heat
a) the reaction is both endothermic and exothermic
b) the reaction is endothermic
c) the reaction is exothermic
d) the reaction is neither endothermic or exothermic
10. Which of the following compounds will not dissolve in water?
a. KCl
b. CCl4
c. CaBr2
11. Which of the following is true?
a. wood is a good conductor of electricity
b. a solution that contains electrolytes conducts electricity
c. a water solution of an acid will not conduct electricity
d. a water solution of a base will not conduct electricity
d. MgCl2
12. A solution with a pH of 9 has an [OH-] concentration of :
a. 1.0 x 10 -14
b. 1.0 x 10 -9
c. 1.0 x 10 -5
d. 1.0 x 10 -7
13. Among the following, which is the strongest acid?
b. pH = 3
d. pH = 10
a. [ H+] = 1 x 10 -5
c. [ OH -] = 1 x 10 -7
14. Which of the following is true about bases?
a. have a bitter taste
c. react with acids to form a salt and water
b. feel slippery
d. all of these
15. A solution in which the hydroxide - ion concentration is 0.00001 is:
a. acidic
b. basic
c. neutral
d. none of these
16. In a neutral solution, the [H+] is:
b. zero
a. 10 -14
c. 1 x 10 -4
d. equal to [ OH-]
17. Because acids and bases are conductors of electricity, they are referred to as:
a. indicators
b. electrolytes
c. insulators
d. nonelectrolytes
18. A solution with a pH of 5.0:
a. is acidic
c. is neutral
b. is basic
d. has a hydrogen ion concentration of 5.0 M
19. Of the following, the best conductor of electricity is:
a. solid salt
b. solid sugar
c. aqueous salt
d. distilled water
20. A solution which has a pH of 12 would be:
a. acidic
b. basic
c. neutral
d. none of these
Topic 7:
Experimentation
7.1 Safety equipment is used to protect the eyes and skin from contact with laboratory chemicals
and flames: goggles, aprons, gloves, safety shower, eyewash, broken glass container, fume hood
and fire blanket. You must understand and follow the laboratory safety rules and procedures that
are described in your Safety Contract in order to work in the chemistry laboratory.
7.2 Historically, scientists have been known to stand on the "shoulders of giants", meaning that they
build on previous knowledge to make new discoveries. This is especially true as scientists began
to develop theories concerning the atom and its structure. From John Dalton, to J.J Thompson to
Ernest Rutherford to Neils Bohr to the modern day electron cloud (quantum mechanical) model
of the atom.
7.3 Significant figures are used in making calculations with measurements made in the lab. Measure
volume of a liquid in milliliters (mL) using a graduated cylinder and stating measured digits plus
the estimated digit. Measure mass in grams (g) using an electronic balance and identifying the
estimated digit. Significant figures are used in making calculations with measurements made in
the lab. Determine the mean (average) of a set of volume or mass measurements using the rules
for significant digits.
7.4 Percent error is the ratio of absolute value of the difference between the experimental value
and the theoretical value to the theoretical value, multiplied by 100
|Theoretical value - Experimental value | X 100
Theoretical value
7.5 Describe and demonstrate safe techniques for lighting and using gas burners. Understand and
use Material Safety Data Sheet (MSDS) warnings including: handling chemicals, lethal dose (LD)
hazards, disposal and chemical spill clean-up.
7.5
Percent yield is the ratio of actual yield to theoretical yield, multiplied by 100.
Actual yield X 100
Theoretical yield
7.6 Accuracy is how close a measurement is to the true value. An accurate measurement has very
little error. Precision is how exact and repeatable a measurement is.
7.7 To safely dilute an acid, add acid to water. Do not add water to a acid concentrated acid.
Always perform this task in a fume hood. When preparing a more dilute solution from a stock solution
of known concentration, M1 X V1 = M2 X V2. The molarity of the first solution multiplied by its volume
will be equal to the molarity of the new solution multiplied by its volume. Phenolphthalein is an
indicator that is colorless in the presence of an acid or neutral substance but is pink in the presence
of a base. Litmus is also an acid/base indicator. Base turn red litmus blue and acids turn blue litmus
red.
7.8 Titration uses a buret to dispense precise amounts of solution of known concentration to
determine the concentration of another solution.
Check For Understanding
1. The closeness of a measurement to it's true value is a measure of it's:
a) accuracy
b) repeatability
c) precision
d) nearness
2. According to the method of significant figures, the number of digits that are estimated in a
measurement is:
a) none
b) one
c) two
d) three
3. What two components does a measurement always contain?
a) a number and a decimal point
c) a number and a unit
b) a power of ten and a unit
d) a number only
4. Safety goggles should be worn in the chemistry laboratory
a) while performing an experiment
b) during clean-up
c) during all parts of the lab and during clean-up
d) only when using a Bunsen Burner
5. If your lab partner's clothing should catch on fire you should:
a) tell him/her to run until they find a security guard.
b) smother the flames with a fire blanket
c) do nothing, because the flames will eventually die out
d) show him/her the door and hope for the best
6. What should you do if you burn your hand?
a) run hot water on it
b) run cold water on it and notify the teacher immediately
c) put your hand in your pocket and continue with your experiment
d) don't do anything, you should have known better
7. When heating a liquid in a test tube you should:
a) slant the test tube away from yourself and others
b) hold the test tube in the vertical position
c) point the test tube toward your lab partner
d) point the test tube toward the teacher
8. How many significant figures in the measurement 2103.2g?
a) 2
b) 3
c) 4
d) 5
9. How many of the zeros in the measurement 0.000040300 are significant?
a) 8
b) 6
c) 5
d) 3
10. Two pieces of equipment that should be worn for every lab are:
a) heavy winter coat and hat
c) goggles
b) gym shorts and a t-shirt
d) goggles and an apron
11. The number of significant figures in the measurement 0.070g is:
a) 1
b) 2
c) 3
d) 4
12. If the mass of a dry beaker is 19.02 grams and increases to 22.40 grams when a sample is added,
what is the mass of the sample?
a) 22.40 g
b) 41.42 g
c) 3.38 g
d) 1.10 g
13. If the volume of water in a cylinder is 8.0mL, but changes to 10.0mL when a solid is carefully
lowered into it, the volume of the solid is:
a) 2.0mL
b) 10.0mL
c) 18.0mL
d) 80. 0mL
14. A student estimated a mass to be 250g, but upon carefully measuring it, found it to be 240g. What
is the percent error of the estimated mass?
a) 4.0%
b) 4.2%
c) -4.0%
d) -4.2%
15. A beaker has a mass of 89.67g. When a solid is added, the beaker plus the solid have a mass of
92.25g. What is the mass of the solid?
a) 1.92g
b) 2.82g
c) 2.58g
d) 12.58g
16. A particular reaction is expected to produce 2.6 L of oxygen gas. In reality, the reaction only
produces 1.9 L of oxygen gas. The percent yield of the reaction:
a) 27 %
b) 42 %
c) 73 %
d) 85 %
17. What type of instrument is used mainly in a titration experiment?
a) buret
b) bunsen burner
c) beaker
d) test tube
18. Phenolphthalein is an indicator for what type of solution?
a) acidic solution
c) NaCl aqueous solution
b) basic solution
d) none of the above
Additional Resources
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Practice SOL Tests: http://education.jlab.org/solquiz/
Interactive Periodic Table: http://site.ifrance.com/okapi/periodic3.htm
Periodic Table Trends: http://courses.chem.psu.edu/chem12/spring/transparencies/pdfs/pmlec12(4).pdf
Bonding, naming and properties of common compounds:
http://users.senet.com.au/%7Erowanb/chem/chembond.htm
Intermolecular bonds: http://www.cs.stedwards.edu/%7Ewright/text/chembond.html
Lewis Structures: http://www.chem.uncc.edu/faculty/murphy/1251/slides/C18a/
Gas Law Formulas: http://www.pmel.org/HandBook/HBpage20.htm
Kinetic Energy: http://plabpc.csustan.edu/general/tutorials/temperature/temperature.htm
Solutions Review: http://members.aol.com/profchm/solintro.html
Polarity: http://library.thinkquest.org/15567/lessons/14.html
Ionization of Water:
http://www.biology.arizona.edu/biochemistry/tutorials/chemistry/page3.html
Percent error: http://www.ric.edu/bgilbert/s3pcerr.htm
Percent Yield: http://www.pathcom.com/~ngjdw/laws.htm
Significant figures: http://dbhs.wvusd.k12.ca.us/SigFigs/SigFigRules.html
Titration: http://www.dartmouth.edu/~chemlab/techniques/titration.html
History of Atom: http://perso.club-internet.fr/molaire1/e_histoire.html
General SOL Review: http://www.wise.k12.va.us/alted/SOL/sol.htm
General SOL Review: http://www.quia.com/pages/sol12.html
Some General and SOL-specific Test Taking Hints and Strategies
1. Get a good solid night's sleep before the test. Being well rested will sharpen the mind and aid your
memory. Eat a good breakfast the morning of the test.
2. Try to make the morning of the test a pleasant one. Avoid stress. Have your things ready and organized
and get to school on time so that you’re not rushing around.
3. Listen carefully to the instructions from the teacher, and read the directions to each question.
4. There are 60 questions on the SOL. Only 50 count, though. The other ten are questions that are under
review. There’s no way to tell which is which, though, so do your best on all of them!
5. You do NOT have to answer all the questions correctly to pass. It is not expected that you can answer
every question correctly. Currently on the Chemistry SOL, you only need to get 27 out of 50 correct. That
comes out to be a 54%. Other SOL tests have other percentages.
 45 out of 50 is advanced!
6. Attempt to answer ALL of the questions and DO NOT leave any blanks. There is no penalty for guessing.
7. The computer has some things that will help in taking the test - i.e., underlining important words, etc. You
should always highlight words like EXCEPT, NOT and other words that limit the answers.
8. Read each question carefully and think before you answer. Be sure that you understand the question
before you start to compare answer choices. This is especially important on the questions containing
graphs and charts or any of these words: CHOOSE, DESCRIBE, EXPLAIN, COMPARE, IDENTIFY, SIMILAR, EXCEPT, NOT, and
BUT.
9. Look at ALL of the answer choices and choose the best and most complete answer.
10. If you're not sure which answer is correct, eliminate choices that you know are incorrect. Then focus on the
remaining choices.
 Elimination ideas:
 Eliminate the answers you know are wrong.
 Eliminate the answers that you know don’t have anything to do with the question.
 You can usually eliminate answers with ALWAYS, NEVER, EVERYONE and words like
that in them.
 If it’s a math question, you can usually eliminate any numbers that appear in the
question.
 If the questions are all numbers, usually the lowest and the highest are wrong.
 If you’ve already used a vocabulary word once on the test and you’re sure it’s right,
then don’t use it again.
11. A word of caution about changing answers - usually your first choice is correct. Only change answers if
you’re certain the first choice was wrong.
12. If you get stuck on a question, use the Review feature and go back to it after finishing the test. You might
want to jot down a note or two about the question on your scrap paper. That way, if you come to another
questions on the same subject, you can use it to help you with the first question, maybe.
13. Stay focused on the test, even if other students finish early. Don't get distracted. No one really cares if you
finish first. There’s no medal or prize or extra credit and no one really thinks that you’re super-smart if you
finish quickly.
14. If you need to take a break or you’re starting to feel overwhelmed, then rest a minute or two. Close your
eyes, take a deep breath, and think happy thoughts. Draw a picture of your teacher on the scrap paper.
Draw a large bus coming towards him. Do something to relax a bit.
15. Here’s a hint to help you double-check each and every answer.
a) You get a piece of scrap paper with each test. Number one edge of the paper 1-60.
b) Take the test, but instead of marking the answer on the computer screen, write it on the scrap
paper after the question number.
c) Double check to make sure you marked all sixty questions.
d) Flip the scrap paper over.
e) Take the test again, but this time mark the correct answer on the computer screen. Do NOT look
at your paper answers!
f) Once you’re finished with the test the second time, compare the answers on the screen to the
answers on your scrap paper. If you came up with the same answer twice, you’re great. If not,
figure out why. If you can’t figure out why you changed your mind, stick with your first answer!
g) If you’re really, really a Type-A personality (you know who you are), you can do the whole scrappaper thing twice if you want and then you’ll have three versions of your answer to compare.
h) Once all your answers match or you’ve figured out why you changed your mind, raise your
hand to submit your answers.
The Actual Chemistry SOLs
CH.1 The student will investigate and understand that experiments in which variables are measured,
analyzed, and evaluated produce observations and verifiable data. Key concepts include
 designated laboratory techniques;
 safe use of chemicals and equipment;
 proper response to emergency situations;
 manipulation of multiple variables, using repeated trials;
 accurate recording, organization, and analysis of data through repeated trials;
 mathematical and procedural error analysis;
 mathematical manipulations (SI units, scientific notation, linear equations, graphing, ratio and
proportion, significant digits, dimensional analysis);
 use of appropriate technology including computers, graphing calculators, and probeware, for
gathering data and communicating results; and
 construction and defense of a scientific viewpoint (the nature of science).
CH.2 The student will investigate and understand that the placement of elements on the periodic table is a
function of their atomic structure. The periodic table is a tool used for the investigations of
 average atomic mass, mass number, and atomic number;
 isotopes, half lives, and radioactive decay;
 mass and charge characteristics of subatomic particles;
 families or groups;
 series and periods;
 trends including atomic radii, electronegativity, shielding effect, and ionization energy;
 electron configurations, valence electrons, and oxidation numbers;
 chemical and physical properties; and
 historical and quantum models.
CH.3 The student will investigate and understand how conservation of energy and matter is expressed in
chemical formulas and balanced equations. Key concepts include
 nomenclature;
 balancing chemical equations;
 writing chemical formulas (molecular, structural, and empirical; and Lewis diagrams);
 bonding types (ionic and covalent);
 reaction types (synthesis, decomposition, single and double replacement, oxidation-reduction,
neutralization, exothermic, and endothermic); and
 reaction rates and kinetics (activation energy, catalysis, and degree of randomness).
CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar
relationships. Key concepts include
 Avogadro’s principle and molar volume;
 stoichiometric relationships;
 partial pressure;
 gas laws;
 solution concentrations;
 chemical equilibrium; and
 acid/base theory: strong electrolytes, weak electrolytes, and nonelectrolytes; dissociation and
ionization; pH and pOH; and the titration process.
CH.5 The student will investigate and understand that the phases of matter are explained by kinetic theory
and forces of attraction between particles. Key concepts include
 pressure, temperature, and volume;
 vapor pressure;
 phase changes;
 molar heats of fusion and vaporization;
 specific heat capacity; and
 colligative properties.