Name: ________________________________________________ Date: __________________________ Per: _______
College Preparatory Chemistry
Level 1
Worksheet Book
1
Name: ________________________________________________ Date: __________________________ Per: _______
Scientific Method
Put the following steps of the scientific method in the proper order.
_____ Research the problem.
_____ Observe and record.
_____ Make a hypothesis.
_____ Identify the problem.
_____ Arrive at a conclusion.
_____ Test the hypothesis.
Match the following terms with the correct definition:
a) organized process used to test a hypothesis
________ 1. hypothesis
b) an educated guess about a solution to a problem
________ 2. control
c) observations and measurements recorded during
an experiment
________ 3. variable
d) a judgement based on the results of an
experiment
________ 4. experiment
________ 5. conclusion
e) a logical explanation for events that occur in
nature
________ 6. theory
f) used to show that the result of an experiment is
really due to the condition being tested
________ 7. data
g) factor that changes in an experiment
2
Name: ________________________________________________ Date: __________________________ Per: _______
Matter Tree
All matter can be classified as either a substance (element or compound) or a mixture (homogeneous or
heterogeneous).
Pure substances are homogeneous and you can write their chemical formula. There are two types of pure
substances, elements and compounds. Elements are one type of atom while compounds are two or more
different atoms chemically bonded together. Mixtures can have variable ratios. There are two types of
mixtures, homogeneous and heterogeneous. A homogeneous mixture have the same composition all the way
through. A heterogeneous mixture does not have the same composition all the way through.
Classify the following as a substance or a mixture. If it is a substance choose either element or compound. If it
is a mixture choose either heterogeneous or homogeneous. Choose 1 of the following:
Substance
Mixture
Type of Matter
Element
Compound
Homogeneous
Heterogeneous
1. Chlorine
2. Water
3. Soil
4. Sugar Water
5. Oxygen
6. Carbon Dioxide
7. Mint Chocolate chip Ice Cream
8. Rubbing Alcohol
9. Pure Air
10. Iron
3
Name: ________________________________________________ Date: __________________________ Per: _______
Element Symbols
Elements are given names and symbols. When writing the formulas of compounds, symbols are much easier to
use than the names. We need to be familiar with the names and symbols of common elements. The symbol of
an element is usually taken from the first one or two letters of the name of the element. However, sometimes
the symbol comes from the Latin name of the element!
Write the symbol for the following elements:
1. oxygen
__________
2. hydrogen
__________
3. chlorine
__________
4. mercury
__________
5. fluorine
__________
6. barium
__________
7. helium
__________
8. uranium
__________
9. radon
__________
10. sulfur
__________
Write the name of the element that corresponds to each of the following symbols:
11. Kr
_________________________
12. K
_________________________
13. C
_________________________
14. Ne
_________________________
15. Si
_________________________
16. Zr
_________________________
17. Sn
_________________________
18. Pt
_________________________
19. Na
_________________________
20. Al
_________________________
4
Name: ________________________________________________ Date: __________________________ Per: _______
Physical vs. Chemical Properties
A physical properties is observed with the senses and can be determined without destroying the object.
For example, color, shape, mass, length and odor are all examples of physical properties.
A chemical property indicates how a substance reacts with something else. The original substance is
fundamentally changed in observing a chemical property. For example, the ability of iron to rust is a chemical
property. The iron has reacted with oxygen, and the original iron metal is changed. It now exists as iron oxide,
a different substance.
Classify the following properties as either chemical or physical by putting a check in the appropriate
column.
Physical Property
1. Blue Color
2. Density
3. Flammability
4. Solubility
5. Reacts with acid to form H2.
6. Supports combustion
7. Sour Taste
8. Melting Point
9. Reacts with water to form a gas
10. Reacts with a base to form water
11. Hardness
12. Boiling Point
13. Can neutralize a base
14. Luster
15. Odor
5
Chemical Property
Name: ________________________________________________ Date: __________________________ Per: _______
Physical vs. Chemical Changes
In a physical change, the original substance still exists, it has only changed form. In a chemical change,
a new substance is produced. Energy changes always accompany chemical changes.
Classify the following as being a physical or chemical change.
1. Sodium hydroxide dissolves in water.
2. Hydrochloric acid reacts with potassium hydroxide to produce a salt, water and
heat.
3. A pellet of sodium is sliced in two.
4. Water is heated and changed to steam.
5. Potassium chlorate decomposes to potassium chloride & oxygen gas.
6. Iron rusts.
7. When placed in H2O, a sodium pellet catches on fire as hydrogen gas is
liberated and sodium hydroxide forms.
8. Evaporation.
9. Ice melting.
10. Milk sours.
11. Sugar dissolves in water.
12. Wood rotting.
13. Pancakes cooking on a griddle.
14. Grass growing in a lawn.
15. A tire is inflated with air.
16. Food is digested in the stomach.
17. Water is absorbed by a paper towel.
6
Name: ________________________________________________ Date: __________________________ Per: _______
Metrics and Measurement
Scientists use the metric system of measurement, based on the number 10. It is important to be able to convert
from one unit to another.
kilo
(k)
1000
103
hecto
(h)
100
102
deca
(da)
10
101
BASE UNIT
Gram (g)
Liter (L)
Meter (m)
deci
(d)
.1
10-1
centi
(c)
.01
10-2
milli
(m)
.001
10-3
Using the above chart, we can determine how many places to move the decimal point and in what direction by
counting the places from one unit to another.
Convert the following:
1) 35 mL =
____________________ dL
6) 4,500 mg = ____________________ g
2) 950 g =
____________________ kg
7) 25 cm =
_______________________ mm
3) 275 mm = ____________________ cm
8) 0.005 kg = ____________________ dag
4) 1,000 L = ____________________ kL
9) 0.075 m = ____________________ cm
5) 1,000 mL =____________________ L
10) 15 g =
7
____________________ mg
Name: ________________________________________________ Date: __________________________ Per: _______
Scientific Notation
Scientists very often deal with very small and very large numbers, which can lead to a lot of
confusion when counting zeros! We have learned to express these numbers as powers of ten.
Scientific notation takes the form of M x 10n, where 1 < M < 10 and “n” represents the
number of decimal places to be moved. Positive “n” indicates the standard (decimal) form is larger
than zero whereas a negative “n’ would indicate a number smaller than zero.
Example 1: Convert 1,500,000 to scientific notation.
We move the decimal point so that there is only one digit to its left, a total of 6
places.
1,500,000 = 1.5 x 106 (“n” is positive because original number is larger than 0)
Example 2: Convert 0.000025 to scientific notation.
For this example, we move the decimal point 5 places to the right, again so that
there is only one digit to its left.
0.000025 = 2.5 x 10-5 (“n” is negative because original number is smaller
that 0)
Example 3: Convert the 9.6 x 10-4 into standard notation
Because “n” is negative, this indicates that the number is smaller than 0. This
means that we move the decimal point 4 places to the left.
9.6 x 10-4 = 0.00096
Convert the following into scientific notation:
Convert the following into standard notation:
a. 0.005
= _______________
a. 1.5 x 103
= _______________
b. 5,050
= _______________
b. 1.5 x 10-3
= _______________
c. 0.0008
= _______________
c. 3.75 x 10-2
= _______________
d. 0.250
= _______________
d. 3.75 x 102
= _______________
e. 0.025
= _______________
e. 2.2 x 105
= _______________
f. 0.0025 = _______________
f. 3.35 x 10-1
= _______________
g. 500.0
= _______________
8
Name: ________________________________________________ Date: __________________________ Per: _______
Significant Figures Worksheet
1. Determine the number of significant figures in the following, then write the numbers in scientific
notation:
# of Significant
Scientific Notation
Figures
a.
200000000
b.
0.0000003287
c.
3,200
d.
4850000000000
e.
0.00000000398700
f.
0.000000000000004790
2. Perform the indicated operation using scientific notation. Remember significant figures!!!
a. 4800000 / 12000000000
____________________
b. (300000) (230000000)
____________________
c. 0.0000004 x 0.400
____________________
3. Complete the following and report the answer to the correct number of significant figures.
a. 4.320 / 8.7 x 3.209
____________________
b. 1.962 + 3.14 + 14.7
____________________
c. 3.0006 – 2.8
____________________
d. (3.20 + 7) x 8.0
____________________
9
Name: ________________________________________________ Date: __________________________ Per: _______
Density
Which has a greater mass, air or lead? Most of you would answer lead, but actually this question does
not have an answer. To compare these two things you need to know how much of each you have. A large
amount of air could have a greater mass than a small amount of lead. To compare different things, we have to
compare the masses of each that occupy the same space, or volume. This is called density.
Density = Mass
Volume
You can remember that density is a broken heart (M/V)
Solve the following problems. Remember to use the correct number of significant figures!
1) What is the density of carbon dioxide gas if 0.196 g occupies a volume of 100.0 mL.?
_____________
2) A block of wood 3.0 cm on each side has a mass of 27 g. What is the density of this block?
_____________
3) An irregularly shaped stone was lowered into a graduated cylinder holding a volume of water equal
to 2.0 mL. The height of the water rose to 7.0 mL. If the mass of the stone is 25 g, what is the
density?
_____________
4) A 10.0 cm3 sample of copper has a mass of 89.6 grams. What is the density of copper?
_____________
Accuracy vs. Precision
1. What is meant by the term “accuracy”?
2. What is meant by the term “precision”?
3. Under what circumstances could a series of measurements of the same quantity be precise
but inaccurate?
10
Name: ________________________________________________ Date: __________________________ Per: _______
4. Which of the following synonyms or characteristics apply to the concept of accuracy and which apply to
precision?
a. multiple measurements
b. correct
c. repeatable
d. reproducible
e. single measurement
f. true value
5. Three students made multiple weighings of a copper cylinder, each using a different balance. The
correct mass of the cylinder had been previously determined to be 47.32 g. Describe the accuracy and
precision of each student’s measurements.
Mass of cylinder (g)
Lissa
Lamont
Leigh Anne
Weighing 1
47.13
47.45
47.95
Weighing 2
47.94
47.39
47.91
Weighing 3
46.83
47.42
47.89
Weighing 4
47.47
47.41
47.93
11
Name: ________________________________________________ Date: __________________________ Per: _______
Temperature & its Measurement
Temperature (which measures the average kinetic energy of the molecules) can be measured using three
common scales: Celsius, Kelvin, & Fahrenheit. We use the following formulas to convert from one scale to
another. Celsius is the scale most desirable for laboratory work. Kelvin represents the absolute scale.
Fahrenheit is the old English scale, which is never used in lab!
C = K – 273
K = C + 273
F = 9/5C + 32
C = 5/9(F – 32)
Complete the following table.
C
1.
0C
K
2.
F
212 F
3.
450 K
4.
5.
98.6 F
-273 C
6.
294 K
7.
77F
8.
225 K
9.
-40C
12
Name: ________________________________________________ Date: __________________________ Per: _______
Percentage Error
Percentage error is a way for scientists to express how far off a laboratory value is from the commonly accepted
value.
The formula is:
% error =  Accepted Value – Experimental Value  x 100
Accepted value
 ~ absolute value
Determine the percentage error in the following problems:
1. Experimental value = 1.24 g
Accepted value
= 1.30 g
Answer
2. Experimental value = 1.24 x 10-2 g
Accepted value
= 9.98 x 10-3 g
Answer
3. Experimental value = 252 mL
Accepted value
= 225 mL
Answer
13
Name: ________________________________________________ Date: __________________________ Per: _______
Atomic Structure 1
An atom is made up of protons and neutrons (both found in the nucleus) and electrons (in the surrounding
electron cloud). The atomic number is equal to the number of protons. The mass number is equal to the
number of protons plus neutrons. In a neutral atom, the number of protons is equal to the number of electrons.
The charge on an ion indicates an imbalance between electrons and protons. Too many electrons produces a
negative charge, too few, a positive charge.
This structure can be written as part of a chemical symbol.
Example:
mass number
15 3+ 
charge
7N
atomic number
Complete the following table:
Element/Ion
Atomic #
H
Atomic mass
Mass #
# Protons
# Neutrons
1
H+
0
12C
6
7 Li+
3
35Cl17
39K
19
24Mg2+
12
As3-
74
Ag
60
Ag+
61
S2-
32
U
238
14
# Electrons
Name: ________________________________________________ Date: __________________________ Per: _______
Atomic Structure 2
1. What is Dalton’s Atomic Theory? Are there any parts of this theory that are no longer valid?
2. Describe JJ Thompson’s “Plum Pudding” Model of the atom. Discuss what the limitations of this model
are.
3. Describe Rutherford’s Gold Foil Experiment and discuss what information it gave us about the structure
of an atom. (Include a drawing)
4. What three subatomic particles make up the atom? What are the relative masses of each?
5. What is the charge on each of the three particles listed in number 4?
6. How is an ion formed?
7. What is an isotope?
8. What is needed to determine the average atomic mass of an element?
15
Name: ________________________________________________ Date: __________________________ Per: _______
Isotopes & Average Atomic Mass
Elements come in a variety of isotopes, meaning they are made up of atoms with the same atomic
number but different atomic masses. These atoms differ in the number of neutrons.
The average atomic mass is the weighted average of all of the isotopes of an atom.
133
Example: A sample of cesium is 75% Cs, 20%
What is its average atomic mass?
Answer:
132
Cs, and 5%
134
Cs.
.7500 x 133 = 99.75
.200 x 132
= 26.4
.0500 x 134
= 6.7
total
132.9 amu = average atomic mass
Determine the average atomic masses of the following:
1.
127
80.0% I
126
17.0% I
128
3.0% I
2.
197
50.0% Au
198
50.0% Au
3.
15.0%
85.0%
Answer
Answer
Answer
55
Fe
56
Fe
16
Name: ________________________________________________ Date: __________________________ Per: _______
Electron Configuration Worksheet #1
1. What are valence electrons?
2. How many valence electrons do the following elements have
a. Be
f. Ca
b. C
g. K
c. Br
h. Ar
d. Xe
i. Al
e. V
j. He
3. Write the electron configurations for the
following elements:
a. Cr
b. Cl
c. O
d. Ca
e. Te
f. Kr
17
Name: ________________________________________________ Date: __________________________ Per: _______
x
Types of Reactions Worksheet
1. Balance the following reactions. Then, indicate the type of reaction for each equation:
a. ____ MgBr2 + ____ Cl2 ____ MgCl2 + ____ Br2
_______________

b. ___ Pb(NO3)2 +
____ H2S
c. ____ HgO

____ Hg
d. ____ Sb
+
____ Cl2
e. ____ CuO + ____ H2
 ____ Cu
____ PbS +
+
____ HNO3
_______________
+
____O2
_______________

____ SbCl3
_______________
____ H2O
_______________
1. Write and balance the equations for the following reactions. Be sure to include the state of matter for
example (s) for solid.
a. aqueous sodium chloride reacts with aqueous lead (II) nitrate to yield a lead (II) chloride precipitate
and aqueous sodium nitrate
b. silver nitrate reacts in solution with potassium chromate to yield a silver chromate precipitate and
soluble potassium nitrate
c. solid calcium carbonate reacts with hydrochloric acid [HCl(aq)] to yield aqueous calcium chloride,
carbon dioxide gas, and liquid water
18
Name: ________________________________________________ Date: __________________________ Per: _______
2. Determine the type of reaction, predict the products and balance the following:
a.
N2
+
H2

b.
Sn
+
Cl2

c.
Fe
+
HCl

d.
MgCl2
+
LiI

Mole Worksheet 3
19
Name: ________________________________________________ Date: __________________________ Per: _______
1.
An extra strength aspirin tablet contains 500. mg of aspirin (C9H8O4). How many molecules of aspirin
are in one extra strength tablet?
2.
A block of salt contains 3.45 x 1026 formula units of NaCl. How many grams is this?
__________
__________
24
3.
A sample of barium nitrate contains 6.80 x 10 formula units. How many grams is this?
4.
__________
If you burned 4.0 x 1024 molecules of natural gas, or methane (CH4), during a laboratory experiment,
what mass of methane did you burn?
__________
5.
How many hydrogen atoms are in 3.5 g of NaOH? Oxygen atoms?
__________
20
Name: ________________________________________________ Date: __________________________ Per: _______
Molarity 1
1) Molarity is a measure of the C____________________ of a solution. It is calculated by dividing the
m__________ by the L________________. The units for molarity are M or _____________________.
2) Find the molarity of the following solutions: (Show work!)
a. 2.30 moles dissolved to make 2.0 Liters of solution.
b. 0.250 moles dissolved to make 300.0 mL of solution.
c. 0.450 moles dissolved to make 605.0 mL of solution.
3) Find the volume of the following: (Show work!)
a. A 0.308 M solution with 4.3 moles dissolved.
b. 3.42 moles in a 5.0 M solution.
c. 2.57 M solution with .025 moles dissolved.
21
Name: ________________________________________________ Date: __________________________ Per: _______
Percent Composition Worksheet
Determine the percent composition of each atom in the following compounds:
1. H3PO4
2. NaF
3. CuFeS2
4. C6H12O6
Empirical and Molecular Formulas Worksheet
1. Determine the empirical formulas of the following;
a. 63.0 g Rb and 5.90 g O
d. 42.7% Co, 57.3% Se
e. 48.8% Cd, 20.8% C, 2.62% H, 27.8% O
22
Name: ________________________________________________ Date: __________________________ Per: _______
Stoichiometry Worksheet #1
Remember to balance all equations!!!
1. How many moles of hydrochloric acid are required to completely dissolve 2.6 moles of magnesium
hydroxide?
Mg(OH)2 (s) + HCl (aq)  MgCl2 (aq) + H2O (l)
2. Mrs. Loughman found a container of solid sodium hydroxide in the chemical storeroom at
Framingham High School. She needs to neutralize this with sulfuric acid. If she found 56.8 moles,
how much sulfuric acid will she need to completely neutralize the sodium hydroxide?
NaOH (s) + H2SO4 (aq)  H2O (l) + Na2SO4 (aq)
3. How many moles of O2 are required to burn 1 mol of C3H8 (propane) molecules in a camping stove?
4. How many moles of O2 molecules are required to burn 1 mol of CH4 molecules from a Bunsen
burner?
23
Name: ________________________________________________ Date: __________________________ Per: _______
Freezing and Boiling Point Graph
25
E
Temperature °C
20
D
15
10
5
C
B
A
0
-5
Energy 
Answer the following questions using the chart above:
1) What is the freezing point of the substance?
_____________________
2) What is the boiling point of the substance?
_____________________
3) What is the melting point of the substance?
_____________________
4) What letter represents the range where the solid is warmed?
_____________________
5) What letter represents the range where the liquid is warmed?
_____________________
6) What letter represents the range where the vapor is warmed?
_____________________
7) What letter represents the melting of the solid?
_____________________
8) What letter represents the vaporization of the liquid?
_____________________
9) What letter(s) show a change in the potential energy?
_____________________
10) What letter(s) show a change in the kinetic energy?
_____________________
11) What letter represents condensation?
_____________________
12) What letter represents crystallization?
_____________________
24
Name: ________________________________________________ Date: __________________________ Per: _______
Heat Calculations
Heat is measured in units of joules or calories. The amount of heat given off or absorbed can be calculated by
the following formula.
q = m x C x T
Heat = (mass in grams) (specific heat) (temperature change)
The specific heat for water is 4.18 joules/grams C
Solve the following problems:
1. How many joules are absorbed by a pot of water with a mass of 500. grams in order to raise the
temperature from 20.0 C to 30.0 C?
2. If the specific heat of iron = 0.460 J/gC, how much heat is needed to warm 50.0 g of iron from 20.0C
to 100.0C?
Heat and Phase Changes
During a phase change, the temperature remains the same. For these calculations, we use the following
formulas.
For freezing and melting  Heat energy = (moles) (Molar heat of fusion)
For boiling and condensation  (moles) (Molar heat of vaporization)
Molar Heat of fusion of water = 6.02 kJ/mol
Molar Heat of vaporization of water = 40.7kJ/mol
Solve the following problems
1) How many joules of heat are necessary to melt 500. g of ice at its freezing point?
2) How many kilojoules is this?
25
Name: ________________________________________________ Date: __________________________ Per: _______
3) How much heat is necessary to vaporize 780.5 g of water at its boiling point?
Heat and its Measurement
Heat (or energy) can be measured in units of calories or joules. When there is a temperature change (T), heat
(q) can be calculated using the formula:
q = mass x specific heat capacity x ∆T
T = Temperature (Final) – Temperature (Initial)
During a phase change, we use the formula:
q = moles x molar heat of fusion (or heat of vaporization)
Answer the following:
1) Define:
Specific Heat Capacity:
Heat of Fusion:
Heat of Vaporization:
2) When should you use the formula q=mCT and when should you use the formula q=nH?
3) How many joules of heat are given off when 5.00 grams of water cools from 75.0°C to 25.0°C? (Specific
Heat of Water = 4.18 J/g°C)
26
Name: ________________________________________________ Date: __________________________ Per: _______
Nuclear Chemistry Worksheet
Radioactive Decay Processes
Radioactive decay processes involve the emission or capture of a number of different particles. Many
radioactive processes result in the change of either the mass number or the atomic number of the atom
undergoing the change. There are two exceptions to this general statement. In the first exception, alpha particle
emission, both the atomic number and the mass number decrease. When gamma radiation is emitted, neither
the atomic number nor the mass number changes. However the emission of gamma radiation always
accompanies another radioactive process.
Summary of Typical Radioactive Processes
Change in nucleus
Example
Emission
Alpha
Mass Number (-4)
Atomic Number (-2)
Beta
Positron
Neutron
Atomic Number (+1)
14C  14N+ 0 e
6
7
-1
Atomic Number (-1)
64Cu  64Ni + 0 e
29
28
+1
Mass Number (-1)
Electron Capture
218Fr  214At + 4 He
87
85
2
Atomic Number (-1)
137I 136I + 1 n
43
53
0
195Au + 0 e  195Pt
79
-1
78
In addition to the reactions shown in the table, atoms can be bombarded with alpha particles, neutrons, or other
atoms to synthesize different radioactive isotopes. These reactions are called transmutation reactions.
1) What is the composition of the nucleus of the following isotopes?
Isotope
Protons
64Ni
28
Neutrons
136I
53
195Au
79
2) Complete the following equations of radioactive decay.
4
a. 241
95 Am  2 He + ________
0
b. 81
37Rb  ________ + +1e
50
c. 50
26Fe  27Co + ________
(Next page for More!)
27
Name: ________________________________________________ Date: __________________________ Per: _______
3) Predict the products of the following nuclear reactions. It helps to determine what type of nuclear decay
takes place in each reaction. Remember to keep track of the total mass number and atomic number on both
sides of the reaction. Also, keep in mind that coefficients require the atomic numbers and mass numbers to
be multiplied by that numerical value (of the coefficient). If there are any missing atomic or mass numbers,
you may obtain them from the periodic table.
0
a. 42K

_______
-1e +
b. 239Pu

4 He
2
+
_______
c. 235
92 U

231Th +
90
_______
d. 11 H
+
3H
1

_______
28
Name: ________________________________________________ Date: __________________________ Per: _______
Gas Laws Worksheet #2
We have now talked about two major gas laws, the combined gas law and the new ideal gas law, PV = nRT.
You need to know when to use each. When do you use the combined gas law?
1. 9.00 moles of a gas at 347C and 102.7 kPa occupy what volume?
Which gas law should you use and why?
2. 7.65 moles of a gas have a volume of 32.5 L at 560 mmHg, what is the temperature?
3. If I have a balloon with a volume of 4.5 liters, and a pressure of 567 mmHg, what will
the pressure be when the volume is 2.5 liters?
4. 15 mL of a gas is collected at 25C and 650 mmHg. What is the volume at STP?
6. 2.5 L of a gas at 64C and 700 mmHg is now subjected to a temperature of -33C and a
pressure of 900 mmHg. What is the new volume?
29
Name: ________________________________________________ Date: __________________________ Per: _______
Bonding/Solutions Worksheet #2
1. Classify the following compounds as ionic or covalent:
a. MgCl2
b. Na2S
c. H2O
d. H2S
2. Draw the Lewis Dot Structure and then use the VSEPR theory to predict the shape of the following
molecules:
a. CO2
b. H2Se
c. SiCl4
d. SCl2
30
Name: ________________________________________________ Date: __________________________ Per: _______
Bonding/Solutions Worksheet
Use the solubility curve to answer the following questions:
1. Which salt shows the smallest change in solubility with
increasing temperature? _________________
2. What mass of ammonium chloride can be dissolved in 200 g
of water at 70C? ___________________
3. If 90 grams of potassium nitrate are dissolved in 200 grams
of water at 50C, what type of solution is this? (circle one)
unsaturated, saturated, supersaturated
4. What happens to the solubility of ammonia as the
temperature increases? ______________________
5. If a saturated solution (in 100 g of water) of KClO3 is
prepared at 90C and then cooled to 40C, what mass of KClO3
will precipitate?
6. What mass of potassium chloride can be dissolved in 175 g
of water at 40C?
7. Which salt is the most soluble at 60C? ________
8. What affect does stirring have on the rate at which a
solution forms? Explain your answer.
9. What affect does the size of the particle have on the rate at which a solution forms? Explain your answer.
31
Name: ________________________________________________ Date: __________________________ Per: _______
Molarity Worksheet 2
Calculate the molarity of the following:
1. There are 2.3 moles of sodium chloride dissolved in 0.45 L of solution.
2. You have 0.0905 moles of sodium sulfate dissolved in 12.0 mL of solution.
3. You dissolve 0.885 moles of magnesium acetate in 5.67 liters of solution.
Use the information given to complete the calculations below:
4. 45.4 g of ammonia are dissolved in enough water to make 0.750 L of solution.
5. What mass of NaOH is required to prepare 1.57 L of a 2.34 M solution?
6. What mass of Na2SO4 is required to prepare 0.750 L of a 0.250 M solution?
32
Name: ________________________________________________ Date: __________________________ Per: _______
Rate Worksheet #1
1. Define the following terms:
a. Rate of reaction
b. Activation energy
c. Activated complex
d. Transition state
e. Catalyst
f. Inhibitor
g. Entropy
2. What is the collision theory and why is it important?
3. Label the following graph:
Reactants, Products, Activation
energy, Activated
complex/transition state, Heat of
reaction, Energy, Reaction
coordinate (progress)
33
Name: ________________________________________________ Date: __________________________ Per: _______
Rate Worksheet #2
1. Know the definitions of the following terms;
Reversible reaction
Le Chatelier’s Principle
Equilibrium
Equilibrium constant
2. Write the equilibrium constant expression for the following unbalanced reactions:
a.
Br2 (g) +
H2S (g)

HBr (g)
+
S (s)
b.
CaCO3 (s)

CaO (s)
+
c.
Cl2 (g)
+
CS2 (g)
CCl4 (g)
CO2 (g)
+
S2Cl2 (g)
3. Le Chatelier’s Principle states that when a system at equilibrium is subjected to a stress (a change), the
system will shift its equilibrium point in order to relieve the stress.
Complete the following chart by writing LEFT, RIGHT, or NONE for equilibrium shift, and DECREASE,
INCREASE, or REMAINS THE SAME for the concentrations of reactants and products, and for the value of K.
(Remember, pure liquids and solids do not affect equilibrium values)
H2 (g)
Stress
1.
add H2
2.
add I2
3.
add HI
4.
remove H2
5.
remove I2
6.
remove HI
7.
increase temp
8.
decrease temp
9.
increase pressure
10.
decrease pressure
+
I2 (g)
Equilibrium Shift
Right

2HI (g)
[H2]
---
[I2]
decrease
∆H= +52.7 kJ
[HI]
increase
-----------
34
Keq
same
Name: ________________________________________________ Date: __________________________ Per: _______
4. Using the following equation:
I2 (g)  2I (g)
A. If the equilibrium concentrations are as follows:
[I2] = 1.2 M, [I] = 0.640 M
write the Keq expression for the reaction.
B. Calculate the value of Keq.
6.
Find the Keq for the reaction if [H2S] = 0.015 M, [H2] = 0.010 M, and the [S2] = 0.051 M.
2 H2S(g)  2 H2 (g) + S2 (g)
35
Name: ________________________________________________ Date: __________________________ Per: _______
[H+] x [OH-] = 1.0 x 10-14
pH + pOH = 14
[H+] = 10-pH
[OH-] = 10-pOH
pH = - log[H+]
pOH = - log[OH-]
Acid Base Worksheet
2.
3.
Write the names for the following acids and bases:
a)
KOH
____________________________________
b)
H2SO4
____________________________________
c)
C2H3O2H
____________________________________
d)
Fe(OH)2
____________________________________
If 0.001 moles of KOH are added to 0.1 L of solution, what is the:
a. pH
b. pOH
4. When a student performs a titration in lab, he finds that it takes 23.2 mL of .340 M HCl to neutralize 14.1
mL of NaOH. What is the concentration of the NaOH?
1. What volume of 0.500 M HCl is required to neutralize 54.6 mL of 1.2 M KOH?
2. What is the molarity of a solution of NaOH if it takes 45.7 mL of a .750 M solution of HNO3 to
neutralize 43.9 mL?
36
Name: ________________________________________________ Date: __________________________ Per: _______
3. If 0.0025 moles of H2SO4 are added to 1.0 L of solution, what is the:
c. pH
d. pOH
4. Determine the Bronsted-Lowry acid-base pairs for the following reactions:
a. HSO3+
H2O 
SO32+
H3O+
b. HNO3
+
H2O

NO3-
c. HC4H7O2
+
H2O

C4H7O2-
+
H3O+
+
3. If 0.0025 moles of H2SO4 are added to 1.0 L of solution, what is the:
a. pH
b. pOH
37
H3O+