Download Chapter 1-3 Exam Review

AP Chemistry
Unit I Targets
Chapters 1-3
By the end of this unit you should be able to . . .
chapter 1
1. define and provide examples for each of the following terms: physical property, chemical
property, physical change, chemical change, intensive property, extensive property element,
compound, mixture
2. differentiate between the three states of matter.
3. list the commonly used metric prefixes and their meanings.
4. determine the number of significant digits in a measured quantity and determine the appropriate
number significant digits in a calculation.
5. convert temperatures between Celsius and Kelvin.
6. perform calculations involving density.
7. convert between units by using dimensional analysis.
chapter 2
1. differentiate between protons, neutrons, and electrons in terms of charge, mass and location in an
atom.
2. I can determine the number of protons, neutrons and electrons in isotopes and in ions.
3. describe the works of John Dalton, J.J. Thomson (cathode ray tube), Robert Millikan (Oil Drop
Experiment) and Ernst Rutherford (Gold Foil Experiment).
4. use the periodic table to predict the charges of monatomic ions.
5. use the periodic table to predict whether an element is a metal, nonmetal or a metalloid.
6. write the names and formulas of ionic compounds, molecular/covalent compounds and acids.
7. calculate the atomic weight of an element given the abundances and masses of its isotopes.
8. distinguish between empirical formulas, molecular formulas and structural formulas.
chapter 3
1. predict the products for and write balanced equations for the following types of reactions:
combustion, decomposition, sysnthesis (called combination reactions), single displacement and
double displacement reactions.
2. interconvert between the number of moles and mass of a substance. I can also use Avogadro’s
number and molar mass to calculate the number of particles (atoms, molecules or formula units)
making up a substance.
3. calculate the percentage composition of a compound by mass.
4. calculate the empirical formula of a compound, having been given either:
a) mass or % composition, or
b) the mass of CO2 and H2O produced by combustion.
5. calculate the molecular formula, having been given the empirical formula and the molecular
weight.
6. use stoichiometry to solve problems involving chemical reactions.
7. determine the limiting reactant in a reaction and determine the amount of excess reactant left over
from a reaction.
8. calculate the theoretical and actual yields of a chemical reaction when given the appropriate data.
1
Target 1. I can define and provide examples for each of the following terms: physical property,
chemical property, physical change, chemical change, intensive property, extensive property
element, compound, mixture
physical property –
chemical property –
physical change –
chemical change –
intensive property –
extensive property element –
compound –
mixture –
Target 2: I can differentiate between the three states of matter.
State
Picture
Movement
Shape
Gas
Liquid
Solid
2
Volume
Compression
Target 3: I can list the commonly used metric prefixes and their meanings.
Unit
Symbol
Giga Mega Kilo Deci Centi Milli Micro Nano Angstrom Pico Femto
Meaning
Metric approximations - name 2 objects with a . . .
a) mass of 1 gram
d) length of a 1 mm
b) mass of 1 kg
e) volume of 1 mL
c) length of 1000 cm
f) volume of 1 liter
List the best metric unit that would be used to measure each of the following:
a)
b)
c)
d)
e)
The distance from Chicago to Detroit. __________
The mass of your body. ________
The volume of coffee held in a cup of coffee. __________
The length of a soccer field. _________
The length of this sheet of paper. ________
Target 4: I can determine the number of significant digits in a measured quantity and determine the
appropriate number significant digits in a calculation.
Rules For Significant Digits
1)
2)
3)
4)
All non-zero numbers are significant.
Exact numbers are assumed to have an infinite amount of significant digits.
Exact numbers are often seen in conversion factors, such as, 3 feet = 1 yard.
SANDWICH RULE - any zero between 2 significant digits is significant
RIGHT/RIGHT RULE - zero’s to the right of a decimal point AND to the right of
another significant digit is significant
Note: Placing a number in scientific notation removes any ambiguity about whether or not a zero is
significant. If a zero is used in scientific notation, then it must be significant.
3
Examples: Underline all of the significant digits in the following measurements:
a) 0.0090 mm
d) 10.00 tons
b) 102 L
e) 2.50 X 10-3 grams
c) 30,400 miles
f) 50 angstroms
Multiplication and Division Rule
The result will have the same number of significant digits as the measurement with the fewest
number of significant digits.
Example: Calculate the density (in g/cm3)of a rectangular solid which has the following
measurements:
height:
2.00 cm
width:
150 mm
length:
0.0500 m
mass:
300.0 grams
Addition and Subtraction Rule
The result cannot have more significant digits to the right of the decimal point than any of the
original numbers.
Example: Add the following numbers (first convert all of the measurements to the
same unit!)
1.000 m + 175 cm + 0.0010 km = ________________
4
Note: When performing a multi-step calculation, always retain at least one additional significant
digit for each of the intermediate answers.
Target 5: I can convert temperatures between Celsius and Kelvin.
Special note about units of temperature: Kelvin and Celsius are metric units of temperature and
one degree of Kelvin is equal to one degree centigrade.
The Kelvin scale is based upon absolute zero. -273C = 0 Kelvin
The Celsius scale is based upon the boiling and freezing point of water. 0C = freezing point of
water and 100C = boiling point of water.
Formula: K = oC + 273
Example 1: convert 195 Kelvin to Celsius.
Example 2: convert 55C to Kelvin.
Target 6: I can perform calculations involving density.
Density is a property of matter that is widely used to characterize a substance. Density is defined as
the amount of mass in a unit of volume of the substance:
D=
mass
volume
Example 1: A rectangular solid has the following dimensions:
length:
2.00 cm
width:
75.0 mm
height:
4.00 cm
mass:
436 grams
Using the given data, is the solid most likely silicon, tin, or gold?
5
Example 2: Assume you had a silver sphere with a mass of 1.50 kg. Calculate the diameter of the
sphere (in cm). The density of silver is 10.5 g/cm3. The formula for the volume of a sphere is . . . V
= 4/3 π r3.
Target 7: I can convert between units by using dimensional analysis.
Example 1: The diameter of metal wire is often referred to by its American wire gauge number. A
16-gauge wire has a diameter of 0.05082 inches. What length of wire, in meters, is there in a 1.00
pound spool of 16-gauge copper wire? The density of copper is 8.92 g/cm3. (There are 2.2 pounds
per kilogram. The volume of a cylinder is found by using the formula V = r2h).
Example 2: A typical rate of deposit of dust from unpolluted air is 10.0 tons per square mile per
month (30 days).
a- What is this dust fall, in milligrams per square meter per hour?
b- If dust has an average density of 2.0 g/cm3, how many hours would it take to accumulate a layer
of dust 1.0 mm thick?
6
Problems
1. A bug travels at the rate of 3.0 miles/hour. How fast is this in m/nsec? Hint:
2.54 cm = 1 inch and 1 mile = 5,280 feet
2.
A cylinder has
Calculate its density in g/cm3.
Diameter = 25.0 mm
Height = 59.0 cm
Mass = 38.0 g
3. The concentration of CO in a room is 48 g/m3. What mass in grams is present in a room
which measures 8.0 x 12.0 x 22 feet?
4.
Ben Franklin showed that 1 teaspoon of oil would cover about 0.50 acre of still
water. If you know that 1.0 x 104 m2 = 2.47 acres, and that there are 5.0 cm3 in a
teaspoon, what is the thickness (in cm) of a layer of oil?
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5.
An empty 3.00 L bottle weighs 1.70 kg. Filled with wine it weighs 4.55 kg. The
wine contains 11.0% ethyl alcohol by mass. How many ounces of ethyl alcohol are
present in a 400.0 mL glass of wine? (1 lb = 16 oz = 453.6 g)
6.
An average human male breathes about 8.50 X 103 liters of air per day. The
concentration of lead in highly polluted air is 7.0 X 10-6 g Pb/m3 of air. Assume
that 75% of the lead is present as particles less than 1.0 X 10-6 m in diameter and
that 50% of the particles below that size are retained in the lungs. Calculate the
grams of lead absorbed in this manner in one year by an average male living in this
environment. (Assume 365 days per year.)
8
AP Chem Unit I
Quick Check For Understanding
Directions: Label each of the following statements as either true or false. If false, change the
statement to make it true.
1.
Air, gaseous carbon dioxide, and light are all examples of matter.
2.
Gases can be compressed quite considerably, liquids can be compressed fairly
easily, and solids are nearly incompressible.
3.
The particles in a solid can move.
4.
Pure substances are either elements, compounds, or homogeneous mixtures.
5.
Melting point, boiling point, and density . . . all three of these properties could be
correctly classified as being both intensive and physical properties of matter.
6.
The gram, kilometer, and hour are the SI units for mass, length, and time
respectively.
7.
150 femtometers is a longer distance than 1.5 picometers.
8.
All of the following represent absolute zero: -273oC, 0 Kelvin, and -419.4oF
9.
Your AP chem teacher gave you a choice in the lab . . . your results could be
either accurate or precise, but not both. The best choice would be for results to
precise.
10.
If you were to round 62,421 km to 3 significant digits, it would be 624 km.
11.
[23.05 - (14.0000 + 6.050)] X 1159 = ________ ; The correct answer to this
problem is 1.802 X 104.
9
be
AP Chem Notes
Chapter 2
Target 1: I can differentiate between protons, neutrons, and electrons in terms of charge, mass and
location in an atom.





The nucleus of the atom is centrally located and is comprised of protons and neutrons.
The electrons are found in regions of space outside the nucleus called orbitals.
Protons and neutrons have a relative mass of 1 amu (1.67 x 10-27 kg).
The proton has a charge of +1 and the neutron has no charge.
The electron has a relative mass of zero (9.11 x 10-31 kg) and a charge of -1.
Target 2: I can determine the number of protons, neutrons and electrons in isotopes and in ions.
Isotopes - atoms of a the same element (so they have the same atomic number) which have different
masses, mass numbers, number of neutrons, and physical properties
# neutrons =
mass # - atomic #
The number of protons determines the identity of the element, but it’s an atom’s electrons (valence)
which determine the chemical properties of the element.
Isotopic Notation
Type I
Carbon-14
Carbon-13
Oxygen-15
Protons
6
6
8
Neutrons
8
7
7
Electrons
6
6
8
Type II
14 -4
C
6
6
8
10
8
7
10
15
8
O-2
10
Target 3: I can describe the works of John Dalton, J.J. Thomson (cathode ray tube), Robert
Millikan (Oil Drop Experiment) and Ernst Rutherford (Gold Foil Experiment).
Scientists that you should be familiar with:
A- Democritus (400 B.C.) -vs- Aristotle
B- John Dalton (early 1800’s) 4 postulates of his atomic theory
1)
2)
3)
4)
C- Henri Becquerel (1896) discovered radioactivity
D- J.J. Thomson (1897) discovery of electron through use of cathode ray tube
E- Robert Millikan (1909) discovered the charge & mass of electron through his “Oil
Drop Experiment”
F- Ernest Rutherford (1910)
(i) discovered 3 types of radiation
(ii) discovered the the nucleus is small, dense, and has a positive charge through
the “Gold Foil Experiment”
(iii) discovered protons in 1919
11
G- James Chadwick (1932) discovered the neutron
Target 4: I can use the periodic table to predict the charges of monatomic ions.
Cations = ions with a positive charge formed by a metal atom losing one or more electrons
Na (atom)
11 protons
11 electrons
Na+ (ion)
11 protons
10 electrons
Anions = ion with a negative charge formed by a nonmetallic atom gaining 1 or more electrons
Charges of common monatomic ions
1A ions
2A ions
3A ions
5A ions
6A ions
7A ions
+1 charge
+2 charge
+3 charge (usually just aluminum)
-3 charge (usually just nitrogen, sometimes P)
-2 charge
-1 charge
Target 5: I can use the periodic table to predict whether an element is a metal, nonmetal or a
metalloid.
Periodic Table Notes
Groups & families
Periods
Group IA
Group IIA
Group VIA
Group VIIA
Group VIIIA
= vertical columns
= horizontal rows
Alkali metals
Alkaline Earth metals
Chalcogens
Halogens
Noble gases / rare gases / inert gases
Metals - elements found on the left side of the “stair case” on the periodic table as well as the
Lanthanoids and Actinides on the bottom, good conductors of heat & electricity, ductile, malleable,
solids at room temperature (except Hg)
Nonmetals - elements found on the right side of the staircase, gases, liquid, & solid; usually poor
conductors and are brittle
Metalloids - elements that lie along staircase which have properties of both metals and nonmetals
(except Al, which is usually considered a metal)
--------------------------------------------------------------------------------------------------------------------Elements found as diatomic molecules in nature include:
hydrogen, oxygen, fluorine, bromine, iodine, nitrogen, chlorine
H2
O2
F2
Br2
I2
N2
Cl2
12
These were discovered by Prof. HOFBrINCl or was his name BrINClHOF?
--------------------------------------------------------------------------------------------------------------------Target 6: I can write the names and formulas of ionic compounds, molecular/covalent compounds
and acids.
Ionic Compounds - compound made up of cations and anions
Naming ionic compounds:
1) Name cations first - cation has the same name as metal; transition metals often need
roman numerals to indicate their charge
2) Name monatomic anions by dropping their ending and adding suffix of “-ide”.
Polyatomic anions most commonly end in “-ate” but can end in “-ite”; do not change
the ending of a polyatomic ion.
The prefix “per-” indicates the anion contains one more oxygen than the common ion. The prefix
“hypo-” indicates that the oxyanion contains one less oxygen atom than the anion which ends in “ite”.
NaClO4
NaClO3
=
=
NaClO2
NaClO
=
=
Naming Acids
prefix
suffix
example
Contains 1 more oxygen
than the common ion
HClO4
Contains the common ion
HClO3
Contains 1 less oxygen
than the common ion
HClO2
Contains 2 less oxygen
than the common ion
HClO
Contains no oxygen
(called binary acids)
HCl
13
Naming Binary Molecular Compounds
1) Name the element farthest to the left in the periodic table first. If both elements are in the same
group, name the lower one first.
2) Use suffix “-ide” on the second element.
3) Use the prefixes listed in table 2.6 in front of both elements (except if you have just one of the
first element).
Name the following binary molecular compounds:
N2O3 ____________________________
SO2 ________________________________
Name or write the formulas for each of the following:
_________
1. phosphoric acid
________________________
11. K2SO4
_________
2. dinitrogen heptoxide
________________________
12. Cu2CO3
_________
3. calcium acetate
________________________
13. (NH4)2S
_________
4. iron (II) chromate
________________________
14. P2O5
_________
5. magnesium hydride
________________________
15. H2CO2
_________
6. hydrofluoric acid
________________________
16. Na3AsO4
_________
7. potassium dichromate
________________________
17. N2O
_________
8. ammonium sulfite
________________________
18. LiMnO4
_________
9. hyposulfurous acid
________________________
19. KOH
_________
10. tin (IV) phosphate
________________________
20. CoCO3
Name each of the following:
NO3- __________________________
NO3 _________________________
14
Target 7: I can calculate the atomic weight of an element given the abundances and masses of its
isotopes.
Atomic mass unit (amu or u) = a unit used to express small masses;
1 amu = 1.66054 X 10-24 g. The amu is defined by assigning a mass of exactly 12 amu to the C-12
isotope.
ATOMIC MASSES: The atomic mass of an element is the weighted average of the masses of the
isotopes of that element. A weighted average reflects both the mass and the relative abundance of
the isotopes as they occur in nature. Most elements occur as two or three isotopes in nature. For
example, chlorine has two isotopes, both of which have 17 protons in their atomic nuclei. One
isotope has 18 neutrons (Cl-35) and the other isotope has 20 neutrons (Cl-37). In order to calculate
the atomic mass (also known as the average atomic mass) you use the following formula:
(mass)(% abundance) + (mass)(% abundance) . . . . = atomic mass
100
example:
isotope
Cl-35
Cl-37
mass (amu)
34.969
36.966
% ab
75.53
24.47
atomic mass = (34.969 amu)(0.7553) + (36.966 amu)(0.2447) = 35.46 amu
Example 1: Magnesium has three naturally occurring isotopes. The data for these isotopes are
given below. Calculate the atomic mass of magnesium.
Isotope Percent
abundance
Mg-24 78.70
Mg-25 10.13
Mg-26 11.17
Mass
(amu)
23.985
24.985
25.983
Example 4: Gallium has two isotopes: Ga-69 (mass of 68.9255 amu) and Ga-70 (mass of 70.9247).
If the average atomic mass of gallium is 69.723 amu, what is the relative abundance (in %) for each
of the two isotopes of gallium?
15
Target 8: I can distinguish between empirical formulas, molecular formulas and structural
formulas.
Empirical Formula - the simplest whole number ratio of atoms of each element present in a
compound.
Lets look at ethene, C2H4…
Ethene can undergo polymerization and become a long strand of linked ethene molecules. The
empirical formula would be C2H4 but the polymer may have a molecular formula several hundred
times larger than the “building block” of the substance.
Molecular Formula - indicates the number of atoms of each element found in each discrete
molecule of that compound.
An example of polyethene: C2H4 x 150 = C300H600
Structural Formula - a graphical representation of the molecular structure, showing how the atoms
are arranged.
Ethene would look like this:
Polyethene would look like this
16
AP Chemistry
Review Chapter 1 and Chapter 2
Name ________________________
1.
_______ is the state of matter which has a definite volume but no definite shape.
2.
Another name for a homogeneous mixture is a(n) ________________.
3.
Can the elements in a compound be separated by chemical means? by physical
means?
4.
List 2 intensive properties and 2 extensive properties.
5.
The SI unit for mass is the _______________; the SI unit for length is the _________.
6.
Which of the following is the longest distance? 1 km, 106 m, or 10-6 Mm?
7.
32°F = ________________°C = ___________________ K
8.
17.2 cm + 204.8 mm = ________________ mm
9.
Which of John Dalton’s postulates was incorrect about atomic theory? Why?
10. 1 meter = _________________ cm
11. Name an element which is:
a. a gas and also a halogen (assume room T)
b. a metalloid
c. an alkaline earth metal
12. Write the empirical formula for glucose.
13. Write the formula for:
a. a polyatomic anion
b. a monatomic cation
c. any molecular compound
d. magnesium cyanide
e. nickel (II) nitride
f. hyponitrous acid
14. Name each of the following:
a. MnO4b. Cr(OH)3
c. NBr3
d. NH4Cl
17
Target 1: I can predict the products for and write balanced equations for the following types of
reactions: combustion, decomposition, sysnthesis (called combination reactions), single
displacement and double displacement reactions.
1- ALKALI METALS WITH WATER
M (s)
+
H2O (l) ---------->
Na(s)
K(s)
+ H2O (l) ---------->
+ H2O (l) ---------->
2- COMBUSTION IN AIR
Hydrocarbon (fuel) +
CH4(g)
+
C3H8(g)
+
CH3OH(l) +
MOH (aq)
NaOH (aq)
KOH (aq)
oxygen ------->
O2(g) ----------->
O2(g) ----------->
O2(g) ----------->
CO2
CO2(g)
CO2(g)
CO2(g)
+
H2(g)
+
+
H2(g)
H2(g)
+
+
+
+
H2O
H2O(l)
H2O(l)
H2O(l)
3- COMBINATION REACTIONS - 2 or more reactants form one product
A + B -------> C (generic equation)
2 Mg(s) +
N2(g)
+
O2(g) --------------> 2 MgO(s)
3 H2(g) -----------> 2 NH3(g)
4- DECOMPOSITION REACTIONS - one reactant forms 2 or more products
A ----------> B + C (generic equation)
2 KClO3(s) ----------->
2 KCl(s)
+
3 O2(g) ; chlorates decompose to form
chloride salts and oxygen gas
2 NaN3(s) ----------> 2 Na(s) + 3 N2(g)
CaCO3(s) ---------> CaO(s) + CO2(g) ; metal carbonates decompose to form metal
oxides and carbon dioxide gas
Target 2: I can interconvert between the number of moles and mass of a substance. I can also use
Avogadro’s number and molar mass to calculate the number of particles (atoms, molecules or
formula units) making up a substance.
Avogadro’s Number: A number equal to the number of atoms in exactly 12 grams of C-12.
Experimentally we have found this number to be 6.022 X 1023.
Molar mass
Avogadro’s #
GRAMS <--------------------> MOLES <--------------------> PARTICLES (atoms, ions, molecules)
18
Example: How many moles of CO2 are in 10.0 grams of CO2?
Practice 1: How many molecules of H2O are contained in 150. grams of water? SYW!
Practice 2: How many hydrogen atoms are there in 36.0 grams of sulfuric acid? SYW!
Formula weight = sum of atomic weights of the atoms in a formula; we can use the term molecular
weight if the substance is a molecular formula or molar mass for either.
Example: Find the formula weight of calcium nitrate.
calcium nitrate is Ca(NO3)2
1 Ca atom = 1(40.1 amu)
2 N atoms = 2(14.0 amu)
6 O atoms = 6(16.0 amu)
=
=
=
40.1 amu
28.0 amu
96.0 amu
___________
164.1 amu
Practice: Find the formula weight of ammonium phosphate.
Target 3: I can calculate the percentage composition of a compound by mass.
Percent Composition: % composition is the % by mass contributed by each element in the
substance
(# of atoms of element)(atomic weight) X 100
molar mass of the compound
19
Example: Find the percentage of nitrogen in calcium nitrate. (Answer is 17.1% N)
Practice: Find the percentage of water in copper sulfate pentahydrate.
Target 4: I can calculate the empirical formula of a compound, having been given either:
a) mass or % composition, or
b) the mass of CO2 and H2O produced by combustion.
mass % ------------> grams of each ---------------> moles of each --------------> empirical
element
element
formula
Example: A compound is 62.58% C, 9.63 % H, and 27.79 % O by mass. Calculate its empirical
formula.
Practice: What is the empirical formula of a compound which contains 25.9% nitrogen and 74.1%
oxygen? (N2O5)
20
Determining the empirical formula of a compound from combustion analysis:
Example: Combustion of a 0.2000 gram sample of vitamin C (which contains C, H, and O) yields
0.2998 g of CO2 and 0.0819 g of H2O. Calculate the empirical formula of vitamin C.
Practice: Combustion analysis of toluene, a common organic solvent, gives 3.53mg of CO2 and
0.822mg of H2O. If the compound contains only carbon and hydrogen, what is its empirical formula?
21
Target 5: I can calculate the molecular formula, having been given the empirical formula and the
molecular weight.
Example: A compound is found to contain 16.66 grams of carbon and 3.49 grams of hydrogen.
a. Find the empirical formula. (C2H5)
b. The molar mass is 87. Find the molecular formula: (C6H15)
Target 6: I can use stoichiometry to solve problems involving chemical reactions.
Stoichiometry is a process which uses balanced equations to solve problems.
Example: How many grams of O2 can be prepared from the decomposition of 4.50 grams of
potassium chlorate?
22
Target 7: I can determine the limiting reactant in a reaction and determine the amount of excess
reactant left over from a reaction.
Example: In order to bake 1 cake you need the following ingredients:
1 box of cake mix
1 egg
2 cups of flour
1/2 cup of sugar
3 tablespoons of water
Assume that you had 1 dozen eggs, 6 boxes of cake mix, 10 cups of flour, 10 cups of sugar, one
quart of oil, and an unlimited supply of water. How many cakes can you make? What is the
limiting “reagent”?
Example: A 2.00 gram piece of zinc is placed in 2.50 grams of aqueous silver nitrate.
a) Write the balanced equation.
b) Determine the limiting reagent.
c) How many grams of silver will be formed?
d) How many grams of the excess reactant will be left over?
23
Target 8: I can calculate the theoretical and actual yields of a chemical reaction when given the
appropriate data.
Percent Yield Problems
Definitions:
theoretical yield = the quantity of product that should “theoretically” be produced (if all
of the reactant is used up)
actual yield = the quantity of product that is actually produced when the reaction is done
in the lab
percent yield = the relationship between the actual yield and the theoretical yield
percent yield = actual yield
X 100
theoretical yield
Example:
Fe2O3(s) + 3 CO(g) -------> 2 Fe(s) + 3 CO2(g)
If 150.0 grams of Fe2O3 reacts with excess CO, 87.9 grams of Fe is produced. Calculate the %
yield.
24
AP Chemistry Ch. 3 Review Sheet
Name_____________________
Directions: Answer the following questions on a separate sheet of paper. Show all of your work,
label your final answers with correct units, and circle your final answers!
1. Predict the products for the following reactions and write a balanced equation for each:
a)
The combustion of glucose.
b)
The synthesis reaction between potassium and chlorine gas.
c)
The decomposition of magnesium carbonate.
d)
Reacting magnesium oxide and water.
e)
Reacting sulfur trioxide and water.
f)
The decomposition of sodium chlorate.
2. Below is a chart containing data for the three naturally occurring isotopes of Mg:
Isotope abundance (%)
mass (u)
Mg-24
78.70
23.98504
Mg-25
10.13
24.98584
Mg-26
11.17
25.98259
Calculate the atomic mass of magnesium.
3. Calculate the percentage of oxygen (by mass) in nickel (II) acetate.
4. Assume you have 5.0 liters of water. Calculate the number of each of the following:
a) the number of grams of water.
b) the number of moles of water.
c) the number of molecules of water.
d) the number of hydrogen atoms in this sample of water.
5. Antifreeze is composed of 51.6 % oxygen, 9.70% hydrogen, and 38.7% carbon by mass. The
molar mass of antifreeze is 62.1 g/mol. Calculate its empirical and molecular formulas.
6. Menthol, the substance we can smell in mentholated cough drops, is composed of C, H, and O.
A 1.005-g sample of menthol is combusted, producing 0.2829 g CO2 and 0.1159 g of H2O. What
is the empirical formula of menthol? If the compound has a molecular mass of 156 g/mol, what is
its molecular formula?
7. When a mixture of 10.0 g of acetylene, C2H2, and 10.0 g of oxygen, O2, is ignited, the resultant
combustion produces CO2 and H2O.
a) Write the balanced equation for this reaction.
b) Which reactant is the limiting reactant?
c) How many grams of C2H2, O2, CO2, and H2O are present after the reaction is complete?
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AP Chemistry Unit I Test Hints
What’s on your test? Anything from chapters 1-3 is fair game . . . even if we have not discussed it
in class. Please look at the objectives for each chapter from the beginning of this packet. Test
questions are taken from this list. Every unit test after this unit will contain material from previous
units. This is to prepare you for the AP exam. Many questions on the AP exam require you to have
knowledge of several topics in order to answer the question.
What is the format of the test? Part I will be multiple choice (no calculators) and part II will be
free response/problems (you can use calculators). When working out a problem, try to do it in as
organized a fashion as possible. You may not get all of the questions correct and the only way I can
give you some partial credit is to try to follow your thought process. Always remember to pay
attention to significant digits and PLEASE put units on your final answers and circle them.
How hard is the test going to be? The test will be challenging. Please do not get discouraged
during the test. Just keep going and answer the questions that you do know. Come back to more
difficult questions.
How should you prepare for this test?
1. Reread the chapters . . . it should make more sense now!
2. Check out the practice problems within the chapter.
3. Look over your notes and any handouts.
4. Make sure you can do all of the assigned problems in the book.
5. Get together with a friend who is also taking AP and study
together.
6. Come in and see me if you need help.
7. Get plenty of rest and exercise.
8. Memorize those persnickity polyatomic ions and metric prefixes!
Study Skills: Learning good study skills is one of the most important things that you should master
before you go to college. Most students taking AP chemistry are also taking other AP courses and
are likely involved in lots of school activities. In order to survive, you will need to become an expert
at time management. The key is not how long you study, but how effectively you are making use of
your time. Develop a daily routine where you have a specific time and place to do your homework.
Keep up on your readings and homework, try to find a group of students that are interested in
forming study groups. Be sure to ask lots of questions during class.
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