AP Chemistry Ch 2 Atoms, Molecules, and Ions
2.1 Contains historical information. Read it if you would like.
2.2 Fundamental Chemical Laws
 Antoine Lavoisier was the first chemist to insist on quantitative experimentation. He got guillotined,
but not for that reason.
 The law of conservation of mass: matter is neither created nor destroyed.
 The law of definite proportions: a given compound always contains exactly the same
proportions of elements by mass.
Mass of Oxygen that combines
 The law of multiple proportions: when two elements combine to
with 1 gram of C
form a series of compounds, the ratios of the masses of the
Compound I
1.33 g
second element that combine with one gram of the first element
Compound II
2.66 g
can always be reduced to small whole numbers. You can see an
example of this on the right. The likely formulas for these compounds would be CO and CO2.
Exercise 2.1 Illustrating the Law of Multiple Proportions
The following data were collected for several compounds of nitrogen and oxygen:
Mass of Nitrogen That Combines With 1 g of Oxygen
Compound A 1.750 g
Compound B 0.8750 g
Compound C 0.4375 g
Show how these data illustrate the law of multiple proportions.
2.3 Dalton’s Atomic Theory
 Dalton’s Theory (partially correct, partially not)
o All matter is made of atoms. These indivisible and indestructible objects are the ultimate
chemical particles.
o All the atoms of a given element are identical, in both weight and chemical properties.
However, atoms of different elements have different weights and different chemical
properties.
o Compounds are formed by the combination of different atoms in the ratio of small whole
numbers.
o A chemical reaction involves only the combination, separation, or rearrangement of atoms;
atoms are neither created nor destroyed in the course of ordinary chemical reactions
o Two modifications were made when subatomic particles and isotopes were discovered.
 Avogadro’s Hypothesis
o At the same temperature and pressure, equal
volumes of different gases contain the same
number of particles.
2.4 Early Experiments to Characterize the Atom
 The electron
o J.J Thomson found that when high voltage was
applied to an evacuated type, a “ray” he called a cathode ray was produced. The ray was
produced at the electrode (also called the cathode) and was repelled by the negative pole
of an applied electric field. He postulated that the ray was a stream of negative particles
(now called electrons). He then measured the deflection of beams of electrons to
determine the charge-to-mass ratio. Thomson discovered that he could repeat this
deflection and calculations using different metal electrodes, showing that all metals contain
electrons and all atoms contain electrons. He also deduced that since atoms were neutral,
there must be a positive charge within the atom, giving rise to the “plum pudding” model.
o Next up, Robert Millikan sprayed charged oil drops into a chamber. He halted their fall (due
to gravity) by adjusting the voltage across two charged plates. He used the stop-drop
voltage and Thomson’s charge-mass ratio to determine the charge on one drop of oil, which
was a whole number multiple of the electron charge.


o The mass of an electron is 9.11 x 10-31 kg.
Radioactivity
o Henry Becquerel famously (and accidentally) discovered radiation when he left a uranium
ore in a closed drawer with a photographic plate. When he realized that the plate had been
exposed, he realized that a form of radiation other than light had penetrated it. The
uranium, of course, was the culprit.
o Three types of radioactive emission
 Alpha (particles): helium nuclei, relatively massive and slow, poorly penetrating,
somewhat dangerous
 Beta (particles): electrons, relatively light and fast, moderately penetrating, a little
more dangerous
 Gamma (rays): just energy, most penetrating, most dangerous
 These are not the only kinds of radioactive emission. We will discuss more in the
spring.
The nuclear atom
o Rutherford’s famous gold foil experiment proved that a positively-charged and somewhat
bulky nucleus could be found in the center of an atom. He also found that atoms are mostly
empty space.
2.5 The Modern View of Atomic Structure (an introduction)
 Elements
o All matter composed of only one type of atom is an element. 92 elements are naturallyoccurring; the rest are manmade.
 Atoms
o The atom is the smallest particle of an element that
retains the chemical properties of that element. It
consists of a bulky, dense nucleus (protons and
neutrons) and electrons shells/clouds (which of course
contain electrons).
o We can find a few pieces of
information about each element using
isotope notation.
 Mass number = #protons +
#neutrons for specific
isotopes of an element
 Actual mass is not an integral number! mass defect--causes this and is
related to the energy binding the particles of the nucleus together
 Atomic number = #protons = #electrons in a neutral atom = identity of the element
Exercise 2.2 Writing the Symbols for Atoms
Write the symbol for the atom that has an atomic number of 9 and a mass number of 19. How many
electrons and how many neutrons does this atom have?

Isotopes
o Isotopes are atoms that have the same number of protons (and therefore are the same
element) but different numbers of neutrons (and therefore different masses).
 Most elements have at least two stable isotopes. Exceptions include Al, F, P.
 Hydrogen isotopes are important because they have special names.
 0 neutrons = hydrogen
 1 neutron = deuterium
 2 neutrons = tritium
2.6 Molecules and Ions
 Electrons are responsible for bonding and chemical reactivity.
o Chemical bonds—forces that hold atoms together
o
o
o
o
o
o
o
o
o
Covalent bonds—atoms share electrons and make molecules [independent units]; H2, CO2,
H2O, NH3, O2, CH4 to name a few.
Molecule--smallest unit of a compound that retains the chem. characteristics of the
compound; characteristics of the constituent elements are lost.
Molecular formula--uses symbols and subscripts to represent the composition of the
molecule. (Strictest sense--covalently bonded)
Structural formula—bonds are shown by lines [representing shared e- pairs]; do not always
indicate shape
Ions--formed when electrons are lost or gained in ordinary chem. reactions; dramatically
affect size of atom
Cations--(+) ions; often metals since metals lose electrons to become + charged
Anions--(-) ions; often nonmetals since nonmetals gain electrons to become - charged
Polyatomic ions--units of atoms behaving as one entity--MEMORIZE formula and charge!
Ionic solids—Electrostatic forces hold ions together. Strong ions held close together solids.
2.7 An Introduction to the Periodic Table
 Metals—malleable, ductile & have luster; most of the
elements are metals—exist as cations in a “sea of
electrons” which accounts for their excellent conductive
properties; form oxides [tarnish] readily and form
POSITIVE ions [cations]. Why must some have such
goofy symbols?
 Groups or families--vertical columns; have similar
physical and chemical properties (based on similar
electron configurations!!)
o Group A—Representative elements
o Group B--transition elements; all metals; have
numerous oxidation/valence states
 Periods --horizonal rows; progress from metals to metalloids [either side of the black “stair step”
line that separates metals from nonmetals] to nonmetals
 MEMORIZE:
o ALKALI METALS—1A
o HALOGENS—7A
o ALKALINE EARTH METALS—2A
o NOBLE (RARE) GASSES—8A
2.8 Naming Simple Compounds
 Binary Ionic Compounds (Type I and Type II)
o In general—consist of a metal cation and a nonmetal anion.
 The cation is written first. The charges from the cation and anion must cancel; we
use subscripts to make this happen.
 The names of ionic compounds do not contain prefixes such as mono- or di- unless
that is part of the name of a polyatomic ion in the compound.
 Monatomic ions end in –ide. Ex. NaF is sodium fluoride.
o Type I contain non-transition metals, which have only one charge when they are cations.
 Group 1A = +1, Group 2A = +2, Aluminum = +3
 Zinc, silver, and cadmium also fit into this category; silver ions always have a +1
charge, while zinc and cadmium ions always have a +2 charge.
 Writing the name of a Type I Binary Ionic compound is simple. Ex. MgCl2 is
magnesium chloride. The formulas are also simple, but you have to swap-and-drop
to get the correct formula. Ex. sodium oxide is Na2O, and calcium nitride is Ca3N2.
o Type II contain transition metals, as well as a few others such as lead, tin, and mercury.
 These ions have variable charges which are reflected in the formula using roman
numerals. For example, FeCl3 would be iron (III) chloride and SnO2 would be tin
(IV) oxide. Conversely, lead (II) chloride would be PbCl2.
 Some of them are real weirdoes. For example, the mercury (II) ion is Hg2+ which
makes sense, but the mercury (I) ion is Hg22+.
Exercise 2.3 Naming Type I Binary Compounds
Name each binary compound:
a. CsF
b. AlCl3
Exercise 2.4 Naming Type II Binary Compounds
Give the systematic name of each of the following compounds.
a. CuCl
b. HgO
c. Fe2O3
d. MnO2
Exercise 2.5 Naming Binary Compounds
Give the systematic name of each of the following compounds.
a. CoBr2
b. CaCl2
c. Al2O3
d. CrCl3

c. LiH
Ionic Compounds With Polyatomic Ions
o Same as the other ionic names/formulas we’ve seen, but you need to look out for
polyatomic ions. I’ve given you a sheet of them, but you will not be given a list for the AP
Exam.
Exercise 2.6 Naming Compounds Containing Polyatomic Ions
Give the systematic name of each of the following compounds.
a. Na2SO4
b. KH2PO4
c. Fe(NO3)3
d. Mn(OH)2
e. Na2SO3
f. Na2CO3
g. NaHCO3
h. CsClO4
i. NaOCl
j. Na2SeO4
k. KBrO3

Binary Covalent Compounds
o Consist of two nonmetals bonded together
o Use prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
o Don’t forget the –ide ending
Exercise 2.7 Naming Type III Binary Compounds
Name each of the following compounds.
a. PCl5
b. PCl3
c. SF6
d. SO3
e. PbCl2
e. SO2
f. CO2

Acids
o Hydrogen is listed first in the formula; the anion is listed second
o -ide →hydro [negative ion root]ic ACID
o -ate →-ic ACID
o -ite → -ous ACID
Exercise 2.8 Naming Acids
Give the systematic name for each of the following acds.
a. H2SO4
b. HClO3
c. HNO3
d. H3PO4
e. HCl
f. H2CO3
g. H2SeO3
h. HBrO2
Exercise 2.9 Writing Acid Formulas
Give the formula for each of the following acids.
a. Hydrobromic acid
b. Perchloric acid
c. Sulfurous acid
d. Acetic acid
f. Dichromic acid

e. Iodic acid
Annoying Things That People Can’t Let Go Of
o Water (easy)
o Ammonia NH3
o Hydrazine N2H4
Exercise 2.10 Naming Various Types of Compounds
Give the systematic name for each of the following compounds.
a. P4O10
b. Nb2O5
c. Li2O2
o
o
o
Phosphine PH3
Nitric oxide NO
Nitrous oxide (laughing gas) N2O
d. Ti(NO3)4
Exercise 2.11 Writing Compound Formulas from Names
Given the following systematic names, write the formula for each compound.
a. Vanadium(V) fluoride
b. Dioxygen difluoride
c. Rubidium peroxide
d. Gallium oxide