CHCA AP Chemistry Summer Assignment
The summer assignment for AP Chemistry is quite simple (but not easy). You need to master the formulas,
charges, and names of the common ions. On the first day of the school year, you will be given a quiz on these
ions. You will be asked to:
• write the names of these ions when given the formula and charge
• write the formula and charge when given the names
I have included several resources in this packet. First, there is a list of the ions that you must know on the first
day. This list also has, on the back, some suggestions for making the process of memorization easier. For
instance, many of you will remember that most of the monatomic ions have charges that are directly related
to their placement on the periodic table. There are naming patterns that greatly simplify the learning of the
polyatomic ions as well.
Also included is a copy of the periodic table used in AP Chemistry. Notice that this is not the table used in first
year chemistry. The AP table is the same that the College Board allows you to use on the AP Chemistry test.
Notice that it has the symbols of the elements but not the written names. You need to take that fact into
consideration when studying for the afore‐mentioned quiz!
I have included a sheet of flashcards for the polyatomic ions that you must learn. I strongly suggest that you
cut them out and begin memorizing them immediately. Use the hints on the common ions sheet to help you
reduce the amount of memorizing that you must do.
Do not let the fact that there are no flashcards for monatomic ions suggest to you that the monatomic ions are
not important. They are every bit as important as the polyatomic ions. If you have trouble identifying the
charge of monatomic ions (or the naming system) then I suggest that you make yourself some flashcards for
those as well.
Doubtless, there will be some students who will procrastinate and try to do all of this studying just before the
start of school. Those students may even cram well enough to do well on the initial quiz. However, they will
quickly forget the ions, and struggle every time that these formulas are used in lecture, homework, quizzes,
tests and labs. All research on human memory shows us that frequent, short periods of study, spread over long
periods of time will produce much greater retention than long periods of study of a short period of time.
I could wait and throw these at you on the first day of school, but I don’t think that would be fair to you. Use
every modality possible as you try to learn these – speak them, write them, visualize them.
Additionally, if you haven’t had chemistry since sophomore year (or if you had CP Chemistry), or you just want to
review to get ahead, please download Chapters 1‐3 Notes off of my website. Work through the problems
(answers are provided for most). We will cover these 3 chapters in 3 weeks!
I look forward to seeing you all at the beginning of the next school year. If you need to contact me during the
summer, you can email me and I will get back to you quickly. Most of you know that I have a website,
mrgansle.site11.com, and there are some review activities in the AP Chemistry section that will give you an
idea whether you have adequately mastered the summer assignment.
Have a blessed summer,
Mr. Gansle
Email: paul.gansle@chca‐oh.org
Common Ions and Their Charges
A mastery of the common ions, their formulas and their charges, is essential to success in
AP Chemistry. You are expected to know all of these ions on the first day of class, when I will give you a
quiz on them. You will always be allowed a periodic table, which makes indentifying the ions on the left
“automatic.” For tips on learning these ions, see the opposite side of this page.
From the table:
Cations
H+
Li+
Na+
K+
Rb+
Cs+
Be2+
Mg2+
Ca2+
Ba2+
Sr2+
Al3+
Name
Hydrogen
Lithium
Sodium
Potassium
Rubidium
Cesium
Beryllium
Magnesium
Calcium
Barium
Strontium
Aluminum
Anions
HFClBrIO2S2Se2N3P3As3Type II Cations
Fe3+
Fe2+
Cu2+
Cu+
Co3+
Co2+
Sn4+
Sn2+
Pb4+
Pb2+
Hg2+
Name
Hydride
Fluoride
Chloride
Bromide
Iodide
Oxide
Sulfide
Selenide
Nitride
Phosphide
Arsenide
Name
Iron(III)
Iron(II)
Copper(II)
Copper(I)
Cobalt(III)
Cobalt(II)
Tin(IV)
Tin(II)
Lead(IV)
Lead(II)
Mercury(II)
Ions to Memorize
Cations
Ag+
Zn2+
Hg22+
NH4+
Name
Silver
Zinc
Mercury(I)
Ammonium
Anions
NO2NO3SO32SO42HSO4OHCNPO43HPO42H2PO4NCSCO32HCO3ClOClO2ClO3ClO4BrOBrO2BrO3BrO4IOIO2IO3IO4C2H3O2MnO4Cr2O72CrO42O22C2O42NH2BO33S2O32-
Name
Nitrite
Nitrate
Sulfite
Sulfate
Hydrogen sulfate (bisulfate)
Hydroxide
Cyanide
Phosphate
Hydrogen phosphate
Dihydrogen phosphate
Thiocyanate
Carbonate
Hydrogen carbonate (bicarbonate)
Hypochlorite
Chlorite
Chlorate
Perchlorate
Hypobromite
Bromite
Bromate
Perbromate
Hypoiodite
iodite
iodate
Periodate
Acetate
Permanganate
Dichromate
Chromate
Peroxide
Oxalate
Amide
Borate
Thiosulfate
Tips for Learning the Ions
“From the Table”
These are ions can be organized into two groups.
1. Their place on the table suggests the charge on the ion, since the neutral atom gains or loses a
predictable number of electrons in order to obtain a noble gas configuration. This was a focus in first
year chemistry, so if you are unsure what this means, get help BEFORE the start of the year.
a. All Group 1 Elements (alkali metals) lose one electron to form an ion with a 1+ charge
b. All Group 2 Elements (alkaline earth metals) lose two electrons to form an ion with a 2+
charge
c. Group 13 metals like aluminum lose three electrons to form an ion with a 3+ charge
d. All Group 17 Elements (halogens) gain one electron to form an ion with a 1- charge
e. All Group 16 nonmetals gain two electrons to form an ion with a 2- charge
f. All Group 15 nonmetals gain three electrons to form an ion with a 3- charge
Notice that cations keep their name (sodium ion, calcium ion) while anions get an “-ide” ending
(chloride ion, oxide ion).
2. Metals that can form more than one ion will have their positive charge denoted by a roman numeral
in parenthesis immediately next to the name of the
Polyatomic Anions
Most of the work on memorization occurs with these ions, but there are a number of patterns that can
greatly reduce the amount of memorizing that one must do.
1. “ate” anions have one more oxygen then the “ite” ion, but the same charge. If you memorize the
“ate” ions, then you should be able to derive the formula for the “ite” ion and vice-versa.
a. sulfate is SO42-, so sulfite has the same charge but one less oxygen (SO32-)
b. nitrate is NO3-, so nitrite has the same charge but one less oxygen (NO2-)
2. If you know that a sufate ion is SO42- then to get the formula for hydrogen sulfate ion, you add a
hydrogen ion to the front of the formula. Since a hydrogen ion has a 1+ charge, the net charge on
the new ion is less negative by one.
a. Example:
PO43Æ
HPO42Æ
H2PO4phosphate
hydrogen phosphate
dihydrogen phosphate
3. Learn the hypochlorite Æ chlorite Æ chlorate Æ perchlorate series, and you also know the series
containing iodite/iodate as well as bromite/bromate.
a. The relationship between the “ite” and “ate” ion is predictable, as always. Learn one and you
know the other.
b. The prefix “hypo” means “under” or “too little” (think “hypodermic”, “hypothermic” or
“hypoglycemia”)
i. Hypochlorite is “under” chlorite, meaning it has one less oxygen
c. The prefix “hyper” means “above” or “too much” (think “hyperkinetic”)
i. the prefix “per” is derived from “hyper” so perchlorate (hyperchlorate) has one more
oxygen than chlorate.
d. Notice how this sequence increases in oxygen while retaining the same charge:
ClOhypochlorite
Æ
ClO2chlorite
Æ
ClO3chlorate
Æ
ClO4perchlorate
Sulfite
Hydrogen
sulfate
Sulfate
Phosphate
Dihydrogen
Phosphate
Hydrogen
Phosphate
Nitrite
Nitrate
Ammonium
Thiocyanate
Carbonate
Hydrogen
carbonate
Borate
Chromate
Dichromate
Permanganate
Oxalate
Amide
Hydroxide
Cyanide
Acetate
Peroxide
Hypochlorite
Chlorite
Chlorate
Perchlorate
Thiosulfate
HSO4- SO42-
SO32-
HPO42-
H2PO4-
PO43 NH4+
NO3-
NO2-
2-
NCSSCN-
HCO3- CO3
Cr2O72-
CrO42-
BO33-
NH2-
C2O42-
MnO4-
C2H3O2CH3COO-
CN-
OH-
ClO2-
ClO-
O22 S2O32 ClO4-
ClO3-
South Pasadena  AP Chemistry
[Keep for Reference]
1  Matter and Measurement
BLUFFER’S
1. Matter
Normally exists in 3 physical states:
LiquidFixed volume, Fluid; Takes on the shape of
lower part of container; well-defined surface
Solid
Rigid Shape; very little volume change as
temperature and pressure change
Gas
Volume expands to fill the container; volume
varies according to temperature and pressure
Kinetic Molecular Theory
The idea that matter consists of molecules or
atoms that are in constant, random motion.
Kinetic Energy = Energy of motion; higher
temperature = more motion
Macroscopic – seen with the eyes.
Microscopic – seen with a microscope
Particulate or Submicroscopic – Structures
at the atomic level (what we think about)
Mixtures
Heterogeneous Mixture – A mixture where
the properties of the mixture vary
throughout. (Like an Oatmeal cookie, the
different components are visible)
Homogeneous Mixture – Also called a
solution, where the components mix at a
molecular level; different properties of the
mixtures are unnoticeable.
Purification – The separation of a mixture
into its components. (techniques: distillation,
filtration, & chromatography)
2. Elements
GUIDE
3. Physical Properties
Properties of a lone sample (ex. mass, volume,
boiling temp, melting temp, conductivity, etc.)
Density is the physical property that relates the
mass of an object to its volume
Mass
Density =
/Volume
Extensive Property – Properties, like mass and
volume, that depend on the amount of substance
Intensive Properties – Properties like color and
density; independent of the amount of substance
Temperature -- how hot a substance is;
physical properties (like density) vary with temp
Celsius 0C for freezing point of water and
100C for melting point of water.
Kelvin – same scale as Celsius; 0C = -273 K;
0 K = no motion; Celsius o + 273 = Kelvin
4. Chemical Properties
How substance interacts with other substances.
Ex. forms gas with acid; burns in air, etc.
5. Physical and Chemical Change
Physical Change – where the identity of all the
substances remain unchanged (melting, boiling,
grinding, pounding into sheets, etc.)
Chemical Change (Reaction) – atoms
rearrange to convert one substance into another
Chemical Equation – A representation of the
chemical reaction taking place
Example: P4 + 6Cl2  4PCl3
6. Measurements/Calculations
Accuracy – how close to a “true value”;
quantified by percent error.
A substance that cannot be decomposed
further by chemical means
Precision – how close measurements are to
each other. Measured by significant figures or
± notation. [I assume you know metric system.]
Names given by symbols: Example: Helium
= He, Gold = Au, Aluminum = Al
Dimensional Analysis – use of a conversion
factor to change units (ex: metric conversions,
mass & volume, time units, etc.)
South Pasadena • AP Chemistry
Name ________________________________
Period ___ Date ___/___/___
2 • Atoms and Elements
STUDY
LIST
The Development of the Atomic Theory:
Molar Mass Calculations:
I can:
 State the four “signpost scientists”, their
experiments, what they added to the atomic
theory, and the name of their model.
 Define the three theories that Dalton explained
in terms of atoms:
o Law of Conservation of Matter
o Law of Definite/Constant Proportions
o Law of Multiple Proportions
 Give examples and solve calculation problems
related to each of the three theories.
 Sketch a cathode ray tube as demonstrated in
class and state how J.J. Thomson’s
experiments led to the idea that atoms have
positive and negative parts, the negative parts
are all the same, and the negative parts (called
electrons) have a certain charge/mass ratio.
 Define cathode rays.
 State the factors that determine how much a
moving charged particle will be deflected by
an electric or magnetic field.
 Explain Millikan’s oil drop experiment & how
it added to the atomic theory.
 Sketch the set-up used by Ernest Rutherford
(the gold-foil experiment), show what he
observed, and explain how these observations
led to the idea that most of the mass of the
atom is concentrated into a tiny, amazingly
massive, positively-charged nucleus.
 Calculate the isotopic mass of an atom given
the resting mass of protons and neutrons.
 Explain that a mole of any element is actually
made up of various isotopes in a constant
percentage abundance.
 Calculate the average atomic mass of an
element using the percent abundance and mass
of each isotope.
 Calculate the percent abundance of isotopes
given the average atomic mass and isotopic
masses of an element.
Parts of the Atom:
 State the three particles that make up an atom,
their symbol, their charge, their mass, and
their location.
 State the number of protons, neutrons, and
electrons in any atom or ion.
 Explain that isotopes are two atoms with the
same atomic number (number of protons) but
different mass numbers (number of
nucleons—protons + neturons).
 Represent the nucleus with isotopic notation,
220
such as: 86 Rn
 Recognize when two nuclei are isotopes of
each other.
The Families of the Periodic Table:
 List the common families of the periodic table
and recognize to which family any element
belongs.
 Recognize metals, non-metals, and metalloids
(semi-metals) on the periodic table.
 State and define the terms conductivity,
malleability, ductility, and sectility.
 State some element facts such as which
elements are too radioactive to exist, which is
the largest non-radioactive element, which
element has the greatest density, and which
element has the highest melting point.
 Explain how Dmitri Mendeleev put together
the periodic table and why we give him credit
for the table even though others were working
along the same lines.
 List the three elements that Mendeleev
predicted and where they are located on the
periodic table.
A Little Nuclear Chemistry:
 State that Henri Becquerel discovered
radioactivity and Marie Curie studied it.
 List the three “Becquerel rays” (alpha, beta,
and gamma) and state why alpha particles
were the perfect tool for Ernest Rutherford to
study the structure of atoms.
 State that the alpha particle is the same as a
helium nucleus, a beta particle is a high-speed
electron, and a gamma ray is a high-energy
form of light.
South Pasadena  AP Chemistry
Name __________________________________
Period ___ Date ___/___/___
3  Chemical Formulas
STUDY
I can:
Formulas
 Look at a formula and state how many
elements and atoms are in that compound.
 Calculate the molecular mass or molar
mass of any compound.
 State that the mass of a molecule is
measured in amu’s and the mass of a mole
is measured in grams.
 Give examples of empirical formulas,
molecular formulas, and structural
formulas.
 Identify a formula as empirical,
molecular, or structural.
Ionic Compounds
I can:
 List the names and formulas of 60 ions.
 State whether a compound is an ionic
compound or a nonmetal compound.
 Write the formula of an ionic compound
given the two ions or its name. Know
when to use parentheses.
 Name an ionic compound given the
formula.
 Determine the charge on an ion from
information in an ionic formula.
Nonmetal Compounds
aka Molecular Compound
 Write the formula of a binary nonmetal
compound (molecular compound) given
its name.
 Name a binary nonmetal compound
(molecular compound) given its formula.
LIST
Percent Composition
 Calculate the percent composition (by
mass) for any compound.
 Calculate the empirical formula from
percent composition data.
 Determine the molecular formula of a
compound given its empirical formula
and molar mass.
Hydrates
 Give examples of hydrates and anhydrous
compounds.
 Calculate the formula of a hydrate from
dehydration data.
The Mole
 State the significance of the mole.
 State the three mole facts for any
substance (molar volume, molar mass,
Avogadro’s number)
1 mole = 22.4 Liters @ STP (gases only)
1 mole = 6.02 x 1023 particles
(particles = molecules or atoms)
1 mole = gram molecular mass of chemical
 Use dimensional analysis to convert
between moles, mass, volume, and number
of particles for a chemical.
 Use density as a conversion factor in mole
problems.
 Use gas density to calculate molar mass.