Download Welcome to AP Chemistry

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Welcome to AP Chemistry!
I hope you are ready for a fun, yet challenging year. You already have a solid background in basic
chemistry from your first year Chem class, and this is critical to success in AP Chem. As the year
progresses and you develop your skills for making connections and problem solving as we delve into great
depth of the study of matter and the changes it undergoes, you will really begin to understand why
chemistry is called the central science, and you will be able to apply your learning to all sorts of different
situations. Students who finish AP Chemistry come out with a much better understanding of the world
around them. They also come out with a sense of great accomplishment. AP Chemistry is a difficult class,
but with determination and perseverance, you will surely succeed!
1. Purchase your own copy of 5 Steps to a 5 on the AP: Chemistry, John T Moore, McGraw Hill. (You
can purchase them at Amazon.com Online, probably the more recent the publication the most
current information however any version will work.)
2. Buy a few color highlighters.
3. Read and study thru the chapter titled ‘How to Approach Each Question Type’. Highlight material
that applies to you. These sections give advice on what to expect and how to study.
4. Take the Diagnostic test. (Go ahead and write in the book, I will
make an additional copy of this test for you to take before the AP Exam.)
5. Take a look at the AP and other websites given in this summer packet. List the three most useful
in the front cover of your book.
6. Read and study (highlight, take notes in the margin, etc) and do all the review questions at the
end of the chapter for these two sections (Chapters vary based on publication year of your 5 Steps
to a 5).
 Basics
 Stoichiometry
AP Chemistry is a difficult course. It is not all about memorization; however, having these items
memorized is essential for success in learning the concepts covered in the course. Make flashcards,
have your friends and family quiz you, take the lists with you on vacation, or do whatever it takes to
get this information firmly planted in your head. Do not wait until the night before school begins.
To give yourself a jump start in the course I recommend seven areas of memorization before the first
day of class:
1. Diatomic molecules (7)
2. Polyatomic Ions (including name, symbol and charge)
3. Variable Valences for Transition Metals
4. Rules for Naming Acids
5. Rules for Naming Ionic Compounds
6. The Solubility Rules
7. Determining Oxidation Numbers
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AP CHEMISTRY SUMMER ASSIGNMENT
The following topics are included in this summer assignment:
- You should memorize the 7 diatomic molecules. When alone and uncharged, the following
substances exist as diatomic molecules: hydrogen, nitrogen, oxygen, fluorine, chlorine,
bromine, and iodine. There are 7 and they form a 7 on the periodic table (with the exception
of hydrogen). Whenever you write the formula for these diatomic molecules it should look
like this: H2, N2, O2, F2, Cl2, Br2, I2. Remember to name them with just the element name.
- Scientific (exponential) notation
- Significant figures
- Naming compounds and Writing formulas from names
- Balancing equations
- Basic chemical conversions
- Basic Stoichiometry
Things to memorize for AP Chemistry:
AP Chemistry is a difficult course. It is not all about memorization; however, having these items
memorized is essential for success in learning the concepts covered in the course. Make flashcards, have
your friends and family quiz you, take the lists with you on vacation, or do whatever it takes to get this
information firmly planted in your head. Do not wait until the night before school begins. Recall of this
information is expected to be automatic throughout the school year. Start studying now! In addition to
these concepts, you will be expected to know all the information in this packet. Stoichiometry in
particular should be mastered prior to beginning the AP Chemistry course.
1. The 7 diatomic molecules
2. Polyatomic Ions (including name, symbol and charge)
3. Variable Valences for Transition Metals
4. Rules for Naming Acids
5. Rules for Naming Ionic Compounds
6. Solubility Rules
7. Determining Oxidation Numbers
Advanced Placement Chemistry is a college level course. You will need to be dedicated and work very
hard if you are to be successful.
Solubility Rules
1. All compounds containing alkali metal cations and the ammonium ion are soluble.
2. All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble.
3. All chlorides, bromides, and iodides are soluble except those containing Ag +, Pb2+, or Hg2+.
4. All sulfates are soluble except those containing Hg2+, Pb2+, Sr2+, Ca2+, or Ba2+.
5. All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+,and Ba2+.
6. All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble except those that also contain
alkali metals or NH4+.
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The 7 Diatomic Molecules
When alone and uncharged, the following substances exist as diatomic molecules: hydrogen, nitrogen, oxygen,
fluorine, chlorine, bromine, and iodine. There are 7 and they form a 7 on the periodic table (with the exception
of hydrogen). Whenever you write the formula for these diatomic molecules it should look like this: H2, N2,
O2, F2, Cl2, Br2, I2. Remember to name them with just the element name.
Rules for Naming an Acid
When the name of the anion ends in –ide, the acid name begins with the prefix hydro-, the stem of the anion has the suffix
–ic and it is followed by the word acid.
-ide becomes hydro _____ic Acid
Cl- is the Chloride ion so HCl = hydrochloric acid
When the anion name ends in –ite, the acid name is the stem of the anion with the suffix –ous, followed by the word acid.
-ite becomes ______ous Acid
ClO2- is the Chlorite ion so HClO2 = Chlorous acid.
When the anion name ends in –ate, the acid name is the stem of the anion with the suffix –ic, followed by the word acid.
-ate becomes ______ic Acid
ClO3- is the Chlorate ion so HClO3 = Chloric acid.
Rules for Naming Ionic Compounds
1. Balance Charges (charges should equal zero)
2. Cation is always written first (in name and in formula)
3. Change the ending of the anion to –ide
4. Roman Numerals are used to identify the charge of the transition metal
Rules for Determining Oxidation Number (this may be the one thing you did not
learn in regular chemistry, but it’s very similar to charges, so it should be easy for
you)
Oxidation Number: A number assigned to an atom in a molecular compound or molecular ion that indicates the
general distribution of electrons among the bonded atoms. (these are a lot like charges, except they are
assigned to covalent compounds in addition to ionic compounds. They just indicate the relative attractions for
electrons.
1. The oxidation number of any uncombined element is O.
2. The oxidation number of a monatomic ion equal the charge on the ion.
3. The more electronegative element in a binary compound is assigned the number equal to the charge it
would have if it were an ion.
4. The oxidation number of fluorine in a compound is always –1
5. Oxygen has an oxidation number of –2 unless it is combined with F, when it is +2, or it is in a peroxide,
when it is –1.
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6. The oxidation state of hydrogen in most of its compounds is+1 unless it combined with a metal, in which
case it is –1.
7. In compounds, the elements of groups 1 and 2 as well as aluminum have oxidation number of +1, +2, and
+3, respectively
8. The sum of the oxidation numbers of all atoms in a neutral compound is O.
9. The sum of the oxidation number of all atoms in a polyatomic ion equals the charge of the ion.
Variable Valences For Transition Metals
Name
Symbol
Chromium Cr
Manganese Mn
Charge
Stock Name
+2
+3
Chromium (II)
Chromium (III)
+2
+3
Manganese (II)
Manganese (III)
+2
+3
Iron (II)
Iron (III)
+2
+3
Cobalt (II)
Cobalt (III)
+1
+2
Copper (I)
Copper (II)
+2
+4
Lead (II)
Lead (IV)
+1
+2
Mercury (I) (this is an odd one. It’s actually Hg22+)
Mercury (II)
+2
+4
Tin (II)
Tin (IV)
+1
+3
Gold (I)
Gold (III)
+1
+2(rarely)
Silver (doesn’t need a roman numeral because it is almost always 1+)
Silver (II)
+3
+5
Bismuth (III)
Bismuth (V)
+3
+5
Antimony (III)
Antimony (V)
Iron
Fe
Cobalt
Co
Copper
Cu
Lead
Pb
Mercury
Hg
Tin
Sn
Gold
Au
Silver
Ag
Bismuth
Bi
Antimony
Sb
Cadmium
Cd
+2
Cadmium (no roman numeral needed because charge doesn’t change)
Zinc
Zn
+2
Zinc (no roman numeral needed because charge doesn’t change)
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Polyatomic Ions – Make Flashcards!!
Name
ammonium
acetate
bromate
chlorate
chlorite
cyanide
dihydrogen phosphate
hypochlorite
hydrogencarbonate(bicarbonate)
hydrogen sulfate (bisulfate)
hydrogen sulfite (bisulfite)
hydroxide
iodate
nitrate
nitrite
perchlorate
permanganate
thiocyanate
carbonate
chromate
dichromate
oxalate
selenate
silicate
sulfate
sulfite
phosphate
phosphite
Symbol
NH4
C2H3O2
BrO3
ClO3
ClO2
CN
H2PO4
ClO
HCO3
HSO4
HSO3
OH
IO3
NO3
NO2
ClO4
MnO4
SCN
CO3
CrO4
Cr2O7
C2O4
SeO4
SiO3
SO4
SO3
PO4
PO3
Charge
+1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-1
-2
-2
-2
-2
-2
-2
-2
-2
-3
-3
Scientific (Exponential) Notation
1. Go to this website for the basics on scientific notation:
http://www.wwnorton.com/college/chemistry/gilbert/tutorials/interface.asp?chapter=chapter_01&folder
=scientific_notation . Make sure you take notes below.
2. Go to this website for information on addition, subtraction, multiplication, and division using scientific
notation. http://www.sparknotes.com/math/algebra1/scientificnotation/section1.html
a. Take notes below.
b. Look to the left corner of the screen to find rules for multiplication and division. Click on this.
c. Go to problems in the upper left corner of the screen. You will find practice problems and
answers. Please write these on your own paper and attach them to this packet.
3. Go to this website for extra practice if desired: http://janus.astro.umd.edu/astro/scinote/
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Scientific Notation Notes:
-
Describe how to put numbers into scientific notation.
-
Describe how to take numbers in scientific notation and write them in regular notation
- Describe how to add and subtract numbers in scientific notation. (you will be required to do
this without a calculator in AP Chemistry, so it is essential you remember how!)
-
Describe how to multiply and divide numbers in scientific notation. (you will be required to
do this without a calculator as well, so make sure you remember how!)
Significant Figures:
You are expected to memorize the rules for determining significant figures and use those rules on every
assignment, test, quiz, and lab report. You will be required to use the rules for significant figures on the AP
exam.
1. Go to this website: http://www.chem.sc.edu/faculty/morgan/resources/sigfigs/index.html
2. Fill out the notes from the website below.
3. Complete the practice on the website. Print out the page. Complete the problems, then check your
work. Attach your work to this packet.
4. Go to this website and take the quiz: http://antoine.frostburg.edu/chem/senese/101/measurement/sigfigquiz.shtml . Be sure that you understand why the answers are correct.
5. Go to this site for practice if needed: http://www.lon-capa.org/~mmp/applist/sigfig/sig.htm or this site:
http://science.widener.edu/svb/tutorial/sigfigures.html
Significant figures notes:
- Explain the concept of significant figures:
Every number that is measured plus 1 that is estimated. Whenever you are using a measuring device,
you should record your measurement so that the last number is estimated. For examples please go to
this site: http://student.ccbcmd.edu/~cyau1/WhatMeaningBehindSigFig.htm
-
-
Define the rules for deciding the number of significant figures in a measurement.
To help define the ambiguity of rule 5, we will also use the following guidelines:
o 10 cm has only 1 significant figure
o 10. cm has 2 significant figures. The decimal point indicates the zero was significant.
o We will usually use scientific notation to avoid confusion about significant figures.
What is an exact number?
What are the rules for determining the number of significant figures in the following types of
mathematical calculations?
o Addition and subtraction
o Multiplication and Division
-
Explain the rules for rounding numbers
-
How do you use significant figures when using a calculator?
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Density
Density is the mass of a substance per unit volume of the substance.
Density = Mass/ volume
The density of an object can be determined through the water displacement method. The object is massed then
submerged in a measured amount of water in a graduated cylinder. The final volume in the graduated cylinder
is read. The volume of water displaced by the object is the volume of the object.
Solve the following problems using correct significant figures:
1. A sample containing 33.42 g of metal pellets is poured into a graduated cylinder containing 12.7 mL of
water, causing the water level in the cylinder to rise to 21.6 mL. Calculate the density of the metal.
2. The density of a piece of metal can be determined from mass and water displacement data. A piece of
metal with a mass of 15.54 g is placed in a flask with a volume of 50.00 cm3. It is found that 40.54 g of
water (d = 0.9971 g/cm3) is needed to fill the flask with the metal in it. What is the density of the metal?
3. The density of copper is 8.96 g/cm3. What is the mass of 18.88 cm3 of pure copper?
4. Describe how you might determine experimentally the density of a solid, such as a sugar cube, which is
water soluble. Indicate what equipment you might use in the process.
Precision vs. Accuracy
Accuracy refers to the agreement of a particular value with a true value
Precision refers to the degree of agreement among several measurements of the same quantity. The degree of
precision refers to the number of digits that a measuring device permits one to measure. In a measuring device,
all except the last digit, which is estimated, are certain. For example, a balance which measures to the nearest
0.0001 g is more precise that one that measures to the nearest 0.01 g.
Percent error = Experimental value = actual value x 100%
Actual value
Types of error:
Random error (indeterminate error) means that a measurement has an equal probability of being high or
low.
Systematic error (determinate error) occurs in the same direction each time; it is either always high or
always low. For example: A balance could have a defect that caused it to read 1.000 g too high every time. A
thermometer could have a defect that cased it to read 1.00 degrees C too low every time.
1. Would using the same balance every time in an experiment be important? Why or why not?
2. In an experiment, the density of aluminum is to be determined. Two students perform the experiment 3
times and obtain the following results:
Trial
Student A
Student B
1
2.45 g/mL
2.68 g/mL
2
2.43 g/mL
2.70 g/mL
3
2.44 g/mL
2.71 g/mL0
a. Describe the accuracy and precision of each student’s results. For student A, calculate the mean
and the percent error if the actual value is 2.70 g/mL.
Temperature:
Remember that science uses temperature in Kelvin almost exclusively. The Kelvin temperature scale is used
because it does not have any negative numbers and the position of 0 K matches the motion of no particle
motion. Temperature is a measure of the average kinetic energy (energy of motion) of a substance. When
kinetic energy is zero, temperature is as low as it can possibly go. This is why the Kelvin temperature scale
makes sense. At zero Kelvin particle motion stops.
To convert from Celsius to Kelvin: K = oC + 273.15
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Classification of Matter
Matter commonly exists in 3 states: solid, liquid, and gas. (plasma also exists, but since we are discussing earth
chemistry, we will ignore it). Throughout this course, you will need to be thinking in terms of particle motion
and particle attraction. These two factors should be the first you think of when considering most problems
throughout the year.
State of matter
Shape
Volume
Particle Motion
Solid
Definite
Indefinite
Slow- vibrate in
place
Liquid
Indefinite
Definite
Particles can move
past each other but
are moving slow
enough to
experience some
attraction.
Gas
Indefinite
Indefinite
Particles moving
very fast. They
move so fast that
particles cannot be
impacted by
attractions
Particle
attractions
Strongest
attractions of the
3 states OR
particles are
moving slow
enough that they
cannot overcome
attractions
Particles are
attracted to each
other but can
overcome the
attraction to
move past each
other. They
cannot overcome
the attractions
enough to escape
the liquid entirely
Particle
attractions are
irrelevant until
there is a high
pressure or low
temperature.
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Identify the following as substance or mixture, then identify mixtures as heterogeneous or homogeneous AND
identify substances as elements or compounds.
1. Carbon
2. salt water
3. sand in water
4. H2O
5. NaCl
6. Cl2
7. Cu
8. Silver
9. CuSO4
Physical versus Chemical Changes
Physical changes- changes that do not change the original composition of the substance. Changes in state such
as boiling or melting are physical changes. Changes involving an alteration in the form of the substance such as
grinding or tearing are physical changes. Bonds are not broken and no reaction occurs.
Physical properties- properties that can be observed without changing the composition of the substance.
Examples: Density, color, and boiling point.
Chemical changes- change the composition of the original substance by breaking and making bonds between
atoms A new substance is produced when a chemical change occurs.
Evidence a chemical change has occurred includes: Change in color, change in odor, production of gas or solid
(precipitate), change in energy. These indications do not always mean a chemical change has occurred, but they
often do.
Examples of chemical properties include: flammability and reactivity to air.
History of the atom and atomic structure:
1. Go to this website:
http://www.neoam.cc.ok.us/~rjones/Pages/online1014/chemistry/chapter_8/pages/history_of_atom.html
2. Read about the history of the atom. Pay special attention to the people and experiments involved in
finding the electron, proton, neutron, and nucleus. In addition, pay special attention to the periodic table
developments.
3. Study the history of the atom. You may wish to take some notes from this site.
4. Make sure you know what protons, neutrons, and electrons are. Make sure you know what ions and
isotopes are. You can check out this website: http://web.jjay.cuny.edu/~acarpi/NSC/3-atoms.htm if you
forgot.
5. Remember that the atomic number indicates the number of protons in an element. This is the subatomic
particle that identifies the element.
6. Electrons are equal to protons if the charge of the atom is zero.
7. A negative charge means that additional electrons were gained by the atom.
8. A positive charge means that electrons were lost by the atom.
9. neutrons = mass number – atomic number
10. Remember that the symbol for an element is often written like this:
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Practice finding electrons, protons, and neutrons:
Element
Symbol
Mass number Atomic
name
number
Carbon
14
6
# of protons
# of neutrons
# of electrons
C
Fluorine
Magnesium
8
35
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Naming Compounds and Writing Formulas
In AP Chemistry, you will be expected to name compounds and write formulas accurately for almost every
problem that is done the entire year. You are expected to have the rules memorized and will not be given the
use of a cheat sheet or flow chart for any test or quiz.
*The first step in any naming or formula situation is to identify whether you have an acid, ionic compound, or
covalent compound. For now follow these rules:
Ionic- contains metal and nonmetal or polyatomic ions
Covalent- contains 2 nonmetals or metalloid and nonmetal
Acid- starts with H (except for water)
*For ionic compounds, you should have most of the common polyatomic ions memorized. You’ll need to know
formula, charge, and name of each ion. I’ve attached a sheet of these for you to study.
*For covalent compounds, you will need to memorize the first 10 prefixes.
1. Go to this website: http://ea008.k12.sd.us/Chemistry%20PP/chemistry_powerpoints.htm
Click on the PowerPoints for Ionic naming and Acid naming, Ionic formulas, and acid formulas, and
covalent naming and formulas. Take notes below.
2. Go to this website to practice both ionic and covalent naming
http://chemistry.csudh.edu/lechelpcs/namingcsn7.html
3. Go to this website: http://misterguch.brinkster.net/pra_namingwkshts.html Click on Naming acids
and bases worksheet. Print it out. Complete the worksheet. Check your work. Attach the worksheet.
4. check out this site for basic information about polyatomic ions:
http://antoine.frostburg.edu/chem/senese/101/compounds/polyatomic.shtml
5. For practice writing formulas: complete the following (keep in mind that covalently bonded substances are
referred to as molecular).
-
Describe naming of ionic compounds
-
Describe writing ionic formulas
-
Describe naming of acids
-
Describe writing acid formulas
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-
Describe naming covalent compounds
-
Describe writing covalent formulas
-
Notes on polyatomic ions
In the following chart, write the proper chemical formula for each of the compounds. Identify each compound
as molecular (M), ionic (I), or acidic (A)
Number Name
Type of compound
Formula
1
Calcium Oxide
2
Phosphorus Pentaiodide
3
Sulfuric Acid
4
Ammonium Perchlorate
5
Barium Hydroxide
6
Molybdenum (II) iodide
7
Aluminum Phosphate
8
Lithium Fluoride
9
Titanium (IV) sulfide
10
Chloric Acid
11
Sodium Peroxide
12
Chromium (III) sulfite
13
Diphosphorus Pentoxide
14
Cobalt (II) chlorite
15
Strontium Permanganate
16
Silver Chromate
17
Carbon Tetrachloride
18
Magnesium nitrite
19
Zinc Oxide
20
Nitrous Acid
21
Potassium Carbonate
22
Lead (IV) nitride
23
Hydrofluoric acid
24
Rubidium nitrate
25
Silicon Dioxide
26
Hydrosulfuric acid
27
Phosphoric acid
28
Sulfurous acid
Balancing equations:
A balanced chemical equation is one that has the same number of atoms on both sides of the equation.
Remember that this must be true due to the law of conservation of mass. Every equation you use this year
MUST be balanced!
1. Go to this website: http://funbasedlearning.com/chemistry/chembalancer/default.htm OR this
website: http://funbasedlearning.com/chemistry/chemBalancer3/default.htm (the first is the easy version. The
second is the harder version. I recommend going to the second. You may turn in a worksheet for either version)
2. Scroll down the page until you see a link called: this worksheet. Click on the worksheet. Fill out the
worksheet as you complete the activity. Attach the worksheet.
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3. If you feel you need more information you may wish to check out one of these sites:
http://www.chemistry.ohio-state.edu/betha/nealChemBal/
http://richardbowles.tripod.com/chemistry/balance.htm
Chemistry Conversions
1. Refresh your memory regarding moles, Avogadro’s number, molar mass, and molar volume by reading
over this website: http://www.cdli.ca/sampleResources/chem2202/unit01_org01_ilo03/b_activity.html
2. Complete the Exercises 1, 2, and 3. Check your work. Attach your work.
3. Perform the practice problems below and check your work using the attached answer key.
Notes- Chemistry Conversions
What is a mole? (please remember that particles can include ions, atoms, molecules, or formula units.
Remember a formula unit is ionic. A molecule is covalent. If a formula unit is broken down it consists of ions.
If a molecule is broken down it consists of atoms.)
What is molar mass?
What is molar volume?
Practice Problems for chemistry Conversions (make sure you show your work- including units and formula of
the chemical. Make sure you answer using significant figures)
1. How many molecules are in 5.6 moles of water?
2. How many atoms are in 5.6 moles of water?
3. How many moles of NaCl are present if you have 5.1 x 1026 formula units of NaCl?
4. How many moles of Cl- are present if you have 7.6 x 1028 formula units of BaCl2?
5. How many grams of (NH4)2SO4 are needed to make a solution using 2.00 moles of (NH4)2SO4?
6. How many liters of hydrogen gas at STP are present if 6.5 moles of hydrogen gas are present?
7. How many atoms of hydrogen are there in 4 L of hydrogen gas at STP?
Stoichiometry:
Make sure you can do Stoichiometry conversions. You will need to be able to perform all types of conversions
without referring to any flow charts or guides. However, if you need help getting started, use the chart below.
**the conversion to liters is not shown on the chart above. Remember that if a gas is at STP, there are
22.4 L in 1 mole of gas.
Page 13
You will need to find the limiting reagent if you are given 2 reactants. Remember that the limiting reagent is
used up first and the excess reagent is left over at the end. BCA tables enable you to easily identify the limiting
reactant, excess reactant, and quantities of product. REMEMBER to use MOLES in your table.
Finding empirical formula (remember that empirical formula is the lowest whole number ratio of elements in a
compound)
1. If you are given a percent, then change the % to a g. If you are given grams, start at #2. If you are
given moles, start at #3.
2. Convert from grams to moles using molar mass
3. Divide all numbers of moles by the smallest number of moles to get the ratio of moles of each
element.
4. If the numbers are whole numbers, use them as subscripts when you write your formula
5. If the number are not whole numbers, multiply all numbers by an appropriate integer to obtain whole
numbers. Then use the new, whole numbers as subscripts when you write the formula
Finding molecular formula. (remember that molecular formula includes every element actually present in the
compound.)
1. Find empirical formula if is was not given already
2. Find the molar mass of the empirical formula
3. Divide the molar mass of the molecular formula by the molar mass of the empirical formula
4. Multiply the subscripts of the empirical formula by the number obtained in number 3.
5. Write your new formula.
Stoichiometry Practice:
1) If a total of 9 mol of NaHCO and 6 mol of C H O react, how many moles of Na C H O will be
produced?
3NaHCO (aq) + C H O (aq) 3CO (g) + 3H O(s) +Na C H O (aq)
2) How many grams of CO are needed to react with an excess of Fe O to produce 8.5 g Fe?
Fe O (s) + 3CO(g) 3CO (g) + 2Fe(s)
3) How many liters of O are needed to react completely with 89 L of H S at STP?
2H S(g) + 3O (g) 2SO (g) + 2H O(g)
4) For the reaction 2Na(s) + Cl (g) 2NaCl(s), how many grams of NaCl could be produced from 5.5 g of Na
and 2.0 L of Cl (at STP)?
5) Solid sodium reacts violently with water, producing heat, hydrogen gas, and sodium hydroxide. How many
molecules of hydrogen gas are formed when 54g of sodium are added to water?
2Na + 2H O 2NaOH + H
6)Consider the following reaction:
2H S(g) + 3O (g) 2SO (g) + 2H O(g)
If you have 4 mol of H S reacting with excess oxygen, and the actual yield is 64.08g of water, what is
the percent yield of the reaction?
7)What is the limiting reagent when 224 g of nitrogen react with 60.6 g of hydrogen?
N (g) + 3H (g) 2NH (g)
8) This equation shows the formation of aluminum oxide: 4Al(s) + 3O2
2O3(s) How many moles of
aluminum are needed to form 7.8 moles of aluminum oxide?
9) How many formula units of KCl are produced by the decomposition of 4.5 x 1025 formula units of KClO3?
2KClO3
2(g)
10) The equation for the combustion of carbon monoxide is:
2CO(g) + O2
2 (g)
How many liters of CO are required to react completely with 54 grams of O2?
Page 14
11)How many formula units of AgNO3 are needed to react completely with 5.74 moles of Cu? The equation is:
2 AgNO3
3)2 + 2 Al
12) If 45.5 g of H are reacted with excess CO, how many grams of CH OH are produced, based on a yield of
98%?
CO(g) + 2H (g) CH OH(l)
13) What is the empirical formula of a compound that is 40% sulfur and 60% oxygen by weight?
14) Calculate the molecular formula of the compound with the following empirical formula and molar mass:
C H , 58 g/mol.
Name:
Hour:
AP CHEMISTRY- Extra Credit PORTION OF THE SUMMER ASSIGNMENT
Scientific Notation:
Write the following numbers in correct scientific notation:
1) 0.000 000 00076
2) 17 000 000 000 000
3) 0.457
4) 1500
5) 3100 x 102
6) 300 x 10-4
Write the following numbers in correct regular (decimal) notation:
7) 5.6 x 105
8) 4.3 x 10-4
9) 1.2 x 103
10) 1.2 x 10-2
11) 15.76 x 10-3
12) 0.098 x 105
Perform the following calculations without the use of a calculator. Write your answers in correct scientific
notation:
13) 5.4 x 105 + 5.4 x 107 =
14) 4.3 x 104 – 1.2 x 105 =
15) 6.7 x 1023 x 2.0 x 1043 =
16) 8.7 x 1054 / 1.2 x 1032 =
Significant figures:
How many significant figures are in each of the following?
1) 1.00 cm
2) 0.000 0045 m
3) 0.000 00450 m
4) 10 m
5) 10. m
6) 15 people
Perform the following calculations. Remember to answer using correct significant figures:
7) 5.432cm + 3.2 cm =
8) 8.857 cm – 1.12 cm =
9) 8.9 cm x 2 cm =
10) 888 cm2 / 2.0 cm =
Page 15
Perform the following calculations without a calculator. Use your knowledge of BOTH scientific notation
AND significant figures.
11) 5.43 x 103 + 5.54 x 105 =
12) 3.2 x 105 – 1.22 x 104 =
13) (4.2 x 1010)(2 x 102) =
14) (8.8 x 1020) / (4 x 1010) =
Naming and writing formulas:
For this section, write from memory the formula and charge for each polyatomic ion indicated:
1) Sulfate
2) Sulfite
3) Nitrate
4) Nitrite
5) Phosphate
6) Phosphate
7) Cyanide
8) Peroxide
9) Hydroxide
10) Chlorate
11) Chlorite
12) Perchlorate (per means 1 more oxygen than the regular –ate)
13) Hypochlorite (hypo- means 1 less oxygen than the regular – ite)
14) Chromate
15) Carbonate
16) Dichromate
17) Bicarbonate (or hydrogen carbonate) (bi means add a hydrogen ion)
18) Oxalate
19) Permanganate
20) Thiosulfate (Thio means: replace one of the oxygen atoms with a sulfur atom)
Indicate the type of compound: Molecular, Ionic, or Acid. Then name the compound.
Formula
Type of Compound
Name
21) CF4
22) KI
23) SF6
24) CaCl2
25) CuCl
26) CuCl2
27) BaO
28) SiO2
29) AlCl3
30) FeSO4
31) Fe2(SO4)3
32) H2CO3
33) HF
34) PbO
35) H2S
36) SnO2
37) H3PO4
38) H3PO3
39) Ba(OH)2
Page 16
Indicate the type of compound: Molecular, Ionic, or Acid. Then write the formula of the compound
Name
Type of compound
Formula
40) Copper (II) carbonate
41) Calcium hydroxide
42) Silicon dioxide
43) hydrochloric acid
44) hydroiodic acid
45) barium phosphate
46) carbon monoxide
47) potassium chloride
48) dinitrogen tetroxide
49) phosphorous acid
50) hydrosulfuric acid
Balancing Equations
Balance the equations below:
1)
Na +
Cl2 
NaCl
2)
Rb +
S8 
Rb2S
3)
H3PO4 +
Ca(OH)2 
Ca3(PO4)2 +
4)
NH3 +
HCl 
NH4Cl
5)
Li +
H2O 
LiOH +
H2
6)
NH3 +
O2 
N2 +
H2O
7)
FeS2 +
O2 
Fe2O3 +
SO2
8)
C+
SO2 
CS2 +
CO
9)
C6H6 +
O2 
CO2 +
H2O
10)
C10H22 +
O2 
CO2 +
H2O
11)
Ca3(PO4)2 +
SiO2 +
C
CaSiO3 +
H2O
CO +
P
Write and balance the following equations:
12) iron + sulfur  iron(II) sulfide
13) zinc + copper(II) sulfate  zinc sulfate + copper
14) silver nitrate + sodium bromide  sodium nitrate + silver bromide
15) potassium chlorate (heated)  potassium chloride + oxygen
16) mercury(II) oxide (heated)  mercury +oxygen
17) potassium iodide + lead(II) nitrate  lead(II) iodide + potassium nitrate
18) aluminum + oxygen  aluminum oxide
19) iron(III) chloride + ammonium hydroxide  iron (III) hydroxide + ammonium chloride
______
Page 17
Precision vs. Accuracy:
1) Balance A measures to the tenths position. Balance B measures to the hundredths position. Balance B is
more
than balance A.
2) Describe the precision and accuracy of the following students: (the true value is 3.5 g)
Student A: 3.5 g, 3.4 g, 3.6 g.
Student B. 5.7 g, 8.7 g, 1.2 g
3) Calculate the percent error of the following experiment. You have determined that the density of Aluminum
is 2.77 g/mL. The true value of the density of Al is 2.70 g/mL.
Classifications of matter:
For each of the following, describe the motion and interaction at a particle level:
1. Solid
2. Liquid
3. Gas
Answer the following questions
4. Two clear solutions combine to form a yellow precipitate and a clear solution. Was this a chemical or
physical change? Explain.
5. Identify each of the following as substances or mixtures. For each substance, identify elements vs
Compounds. For each mixture, identify heterogeneous vs. homogeneous.
a. Cu
b. Zn
c. H2O
d. Sugar water
e. Dirt in water
Chemistry Conversions and Stoichiometry- You must show all work, include units and formulas, and box or
highlight your answers for this section. Be sure to answer using significant figures.
1. How many formula units are in 4.5 moles of NaCl?
2. How many ions are in 4.5 moles of NaCl?
3. How many moles of H2 are present if you have 8.1 x 1026 molecules of H2?
Page 18
4. How many moles of Cl- are present if you have 4.7 x 1027 formula units of AlCl3 ?
5. How many grams of (NH4)3PO4 are needed to make a solution using 2.00 moles of (NH4)3PO4?
6. How many liters of oxygen gas at STP are present if 7.4 moles of oxygen gas are present?
7. How many atoms of chlorine are there in 4.00 L of chlorine gas at STP?
8. How many grams of CO are needed to react with an excess of Fe O to produce 4.25 g Fe?
Fe O (s) + 3CO(g) 3CO (g) + 2Fe(s)
9. How many liters of O are needed to react completely with 45 L of H S at STP?
2H S(g) + 3O (g)
2SO (g) + 2H O(g)
10. For the reaction 2Na(s) + Cl (g)
2.5 L of Cl (at STP)?
2NaCl(s), how many grams of NaCl could be produced from 4.5 g of Na and
11. Solid sodium reacts violently with water, producing heat, hydrogen gas, and sodium hydroxide. How many
molecules of hydrogen gas are formed when 87g of sodium are added to water?
2Na + 2H O
2NaOH + H
12. Consider the following reaction:
2H S(g) + 3O (g)
2SO (g) + 2H O(g)
If you have 8 mol of H S reacting with excess oxygen, and the actual yield is 128.16 g of water, what is the
percent yield of the reaction?
13. The equation for the combustion of carbon monoxide is:
Page 19
2CO(g) + O2(g)  2 CO2 (g)
How many liters of CO are required to react completely with 8.7 grams of O2?
14. How many formula units of AgNO3 are needed to react completely with 7.9 moles of Cu? The equation
is: 2 AgNO3 + Cu  Cu(NO3)2 + 2 Al
15. If 85.7 g of H are reacted with excess CO, how many grams of CH OH are produced, based on a yield of 98%?
CO(g) + 2H (g)
CH OH(l)
16. What is the empirical formula of the compound that is 42.10 % carbon, 5.26 % hydrogen, 24.56 %
nitrogen, and 28.07 % oxygen? If molecular mass of the compound is found to be 171.2 g/mol what is
it's molecular formula?
Atomic Structure:
1. How was the nucleus discovered?
2. How was the electron discovered?
3. How was the proton discovered?
4. How was the neutron discovered?
5. Fill out the following chart:
Element
name
Iodine
Symbol
14
7
Mass number
Atomic
number
# of protons
Page 20
# of neutrons # of electrons
N
Aluminum
28
14
5
7
8
Density - Answers:
1. 3.8 g/mL
2. 1.66 g/cm3 The volume of water displaced by the piece of metal is 50.00 cm3 – 40.66 cm3 = 9.34
cm3. Using mass/ volume = 15.54 g / 9.34 cm3 = 1.66 g/cm3. Not that 40.54 g of water with density
of 0.9971 g/cm3 has a volume of 40.66 cm3.
3. 169 g. From density = m/v, solving for mass and using mass = d x v = 8.96 g/cm3 x 18.88 cm3. You
could also estimate for a multiple choice problem and using 9 x 20 to get an answer of about 180 g.
This would make a multiple choice problem go faster!
4. There are 2 methods. First: Since density is a ratio of the mass of the object compared to the
volume of the object, you might first determine the mass of the sugar cube using a balance. Greater
precision is possible with an analytical balance. Then determine the volume of this regular cube by
measuring its length, with, and depth. Then calculate the volume. V = l x w x d = side 3 Then just
divide mass by the volume. Second method: Use liquid displacement to find volume. However, the
liquid used may not be water because the object is soluble in water. Select a nonpolar liquid to keep
the polar solid from dissolving. Perhaps 1,1,1-trichloroethane. Use a balance to find the mass.
Divide mass by volume.
Types of Error - Answers.
1. Using the same balance is important because it may have some error. If you use the same balance every
time, the error will be systematic error, which is unlikely to matter in some experiments because many
times we will be measuring the difference between 2 masses. If you used 2 different balances, the error
may be magnified if both balances have systematic errors in different directions.
2. Student A was precise but inaccurate. Student B was precise and accurate.
Classification of Matter - Answers:
1. substance- element
2. mixture- homogeneous
3. mixture- heterogeneous
4. substance- compound
5. Substance- compound
6. Substance- element
7. Substance- element
8. Substance- element
9. Substance- compound
Practice finding electrons, protons, and neutrons - Answers:
Element
name
Carbon
Carbon-14
(less common
isotope of C)
Fluorine
Magnesium
Oxygen
Chlorine
Symbol
12
6
14
6
19
9
24
12
16
8
35
17
C
C
F
Mg
O
Cl
# of protons
12
Atomic
number
6
Page 21
# of neutrons # of electrons
6
6
6
14
6
6
8
6
19
9
9
10
9
24
12
12
12
12
16
8
8
8
8
35
17
17
18
17
Mass number
Formula practice – Answers:
Number Name
1
Calcium oxide
2
Phosphorus Pentaiodide
3
Sulfuric Acid
4
Ammonium Perchlorate
5
Barium Hydroxide
6
Molybdenum (II) iodide
7
Aluminum Phosphate
8
Lithium Fluoride
9
Titanium (IV) sulfide
10
Chloric Acid
11
Sodium Peroxide
12
Chromium (III) sulfite
13
Diphosphorus Pentoxide
14
Cobalt (II) chlorite
15
Strontium Permanganate
16
Silver Chromate
17
Carbon Tetrachloride
18
Magnesium nitrite
19
Zinc Oxide
20
Nitrous Acid
21
Potassium Carbonate
22
Lead (IV) nitride
23
Hydrofluoric acid
24
Rubidium nitrate
25
Silicon Dioxide
26
Hydrosulfuric acid
27
Phosphoric acid
28
Sulfurous acid
Type of compound
I
M
A
I
I
I
I
I
I
A
I
I
M
I
I
I
M
I
I
A
I
I
A
I
M
A
A
A
Formula
CaO
PI5
H2SO4
NH4ClO4
Ba(OH)2
MoI2
AlPO4
LiF
TiS2
HClO3
Na2O2
Cr2(SO3)3
P2O5
Co(ClO2)2
Sr(MnO4)2
Ag2CrO4
CCl4
Mg(NO3)2
ZnO
HNO2
K2CO3
Pb3N4
HF
Rb(NO3)2
SiO2
H2 S
H3PO4
H2SO3
Page 22
Chemistry Conversions – Answers to practice problems for chemistry conversions: (remember that definitions
and definite numbers have infinite significant figures, so the only significant figures you need to worry about
are the ones in the beginning number)
1)(5.6 mol H2O)(6.022 x 1023 molecules of H2O) = 3.4 x 1024 molecules of H2O
1 mol H2O
23
2) (5.6 mol H2O)(6.022 x 10 molecules of H2O)(3 atoms of H and O) = 1.0 x 1025 atoms of H and O
1 mol H2O
1 molecule H2O
26
3) (5.1 x 10 formula units NaCl)(1 mol NaCl
) = 8.5 x 102 moles NaCl or 850 moles
23
6.022 x 10 formula units NaCl
28
4) (7.6 x 10 formula unitsBaCl2)(1 mol BaCl2
)(2 mol Cl) = 2.5 x 105 mol Cl23
6.022 x 10 formula units BaCl2
1 mol BaCl2
5) (2.00 mol (NH4)2SO4)(132.17 g (NH4)2SO4) = 264.34 g (NH4)2SO4 = 264 g (NH4)2SO4
1 mol (NH4)2SO4
6) (6.5 moles H2)(22.4 L H2) = 145.6 L H2 = 150 L H2
1 mol H2
7) (4 L H2)(1 mol H2) = 0.2 mol H2
22.4 L H2
Stoichiometry Practice - Answers:
1. (9 mol NaHCO3)(1 mol C6H8O7) = 3 mol of C6H8O7 are needed to use up all the NaHCO3. We have
3 mol NaHCO3
6 mols of C6H8O7, so NaHCO3 is the limiting reagent.
(9 mol NaHCO3)(1 mol Na3C6H5O7) = 3 mol Na3C6H5O7
3 mol NaHCO3
2. Fe O (s) + 3CO(g) 3CO (g) + 2Fe(s)
(8.5 g Fe) (1 mol Fe) (3 mol CO) (28.01 g CO) = 6.4 g CO
55.85 g Fe 2 mol Fe 1 mol CO
3. 2H S(g) + 3O (g) 2SO (g) + 2H O(g)
(89 L H2S) (1 mol H2S) (3 mol O2 ) (22.4 L O2) = 133.5 L O2
22.4 L H2S 2 mol H2S
1 mol O2
4. 2Na(s) + Cl (g) 2NaCl(s)
(5.5 g Na)(1 mol Na) (1 mol Cl2)(22.4 L Cl2) = 2.7 L Cl2 needed to use up all the Na. We have 2.0 L Cl2, so Cl2
22.99 g Na 2 mol Na 1 mol Cl2
is the limiting reagent.
(2.0 L Cl2)(1 mol Cl2)(2 mol NaCl) (58.44 g NaCl) =10. g NaCl
22.4 L Cl2 1 mol Cl2 1 mol NaCl
5. (54 g Na)(1 mol Na )(1 mol H2 )(6.022 x 1023 cules H2) = 7.1 x 1023 cules H2
22.99 g Na 2 mol Na
1 mol H2
6. (4 mol H2S)(2 mol H2O)(18.02 g H2O) = 72.08 g H2O – Theoretical Yield
2 mol H2S 1 mol H2O
64.08 g H2O = 89 % yield of H2O
72.08 g H2O
7. (224 g N2)(1 mol N2 ) (3 mol H2) (2.02 g H2) = 48.4 g H2 are needed to use up all the N2. We have
28.02 g N2 1 mol N2 1 mol H2
60.6 g of H2, so N2 is the limiting reagent.
8. 4Al(s) + 3O2(g)  2Al2O3(s)
(7.8 mol Al2O3)(4 mol Al
) = 15.6 mol Al
2 mol Al2O3
9. 2KClO3(s)  2KCl (s) + 3O2(g)
( 4.5 x 1025 formula units of KClO3)(1 mol KClO3
)(2 mol KCl ) (6.022 x 1023 f.u. KCl) =
6.022 x 1023 f.u. KClO3 2 mol KClO3
1 mol KCl
25
4.5 x 10 f.u. KCl
10. 2CO(g) + O2(g)  2 CO2 (g)
Page 23
(54 g O2) (1 mol O2 )(2 mol CO) (22.4 L CO) = 76 L CO
32.00 g O2 1 mol O2
1 mol CO
11. 2 AgNO3 + Cu  Cu(NO3)2 + 2 Al
(5.74 moles of Cu) (2 mol AgNO3)(6.022 x 1023 f.u. AgNO3) = 6.91 x 1024 f.u. AgNO3
1 mol Cu
1 mol AgNO3
12. (45.5 g of H2) (1 mol H2 )(1 mol CH3OH) (32.05 g CH3OH) = 361 g CH3OH
2.02 g H2
2 mol H2
1 mol CH3OH
361 g CH3OH x .98 = 353.7 g CH3OH
13. 40% sulfur and 60% oxygen
40 g S(1 mol S
) = 1.247 mol S
(60g O)( 1mol O
) = 3.75 mol O
32.07 g S
16.00 g O
Divide by the smallest number of moles:
S: 1.247 mol = 1
O: 3.75 mol O = 3
1.247 mol
1.247 mol
Formula SO3
14) C2H5 58 g/mol Molar mass of empirical: (2)(12.01g/mol) + (5)(1.01 g/mol) = 29.07 g/mol
Divide molar mass of molecular formula by molar mass of empirical:
58 g/mol
29.07 g/mol
C4H10
= 2 Multiply the subscripts of the empirical formula by 2 to get the molecular formula