Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 2
Atoms, Molecules,
and Ions
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.
Atoms,
Molecules,
and Ions
2. 1
The Atomic
Theory of
Matter
Atoms,
Molecules,
and Ions
Democritus (~400 BC)
• Greek philosopher
• His theory: Matter could not be
divided into smaller and smaller
pieces forever; eventually the
smallest possible piece would
be obtained.
• This piece would be
indivisible
• He was the first to
contemplate the “atom”
• He named the smallest piece of
matter “atomos,” meaning “not
to be cut.”
Atoms,
Molecules,
and Ions
• Democritus’s theory
was forgotten for over
2000 years
• The philosophers of the
time, Aristotle and
Plato, had a more
respected, (and
ultimately wrong)
theory.
• Aristotle and Plato
favored the earth, fire,
air and water approach
to the nature of matter.
Atoms,
Molecules,
and Ions
Atomic Theory of Matter
• The theory that atoms are
the fundamental building
blocks of matter reemerged
in the early 19th century,
championed by John Dalton.
Known as the “Father of
the Atom”
First scientist to support
the existence of the atom
with scientific evidence
Dalton’s Postulates
1) Elements are made up of atoms
2) Atoms of each element are identical (in
mass and properties). Atoms of different
elements are different.
3) Chemical reactions are a rearrangement
of atoms. Atoms are not created or
destroyed.
4) Compounds are formed when
atoms of multiple elements combine.
Each compound has specific
Atoms,
Molecules,
and Ions
numbers and kinds of atoms.
Dalton’s Atomic Theory
• He deduced that all
elements are composed
of atoms. Atoms are
indivisible and
indestructible particles.
• Atoms of the same
element are exactly alike.
• Atoms of different
elements are different.
• Compounds are formed
by the joining of atoms of
two or more elements.
Atoms,
Molecules,
and Ions
Law of Constant Composition
• From Dalton’s 4th postulate
• Also known as the law of definite
proportions.
• The elemental composition of a pure
substance never varies.
 In a compound or element, the numbers
and kinds of atoms are constant.
Atoms,
Molecules,
and Ions
Law of Multiple Proportions
• When chemical elements combine, they do so in a ratio of
small whole numbers.
 For example, carbon and oxygen react to form carbon
monoxide (CO) or carbon dioxide (CO2), but not CO1.3.
• Further, if two elements, A & B, combine to form more than
one compound, the masses of B that combine with a given
mass A are in the ratio of small whole numbers
• Example:
Water H2O vs Hydrogen Peroxide H2O2
Mass O = 16 g
Mass O = 32 g
Mass H = 2 g
Mass H = 2 g
Oxygen  32 g : 16 g = 2
Hydrogen  2 g : 2 g = 1
Atoms,
Molecules,
and Ions
Law of Conservation of Mass
• From Dalton’s 3rd postulate
• The total mass of substances present at
the end of a chemical reaction is the
same as the mass of substances
present before the reaction took place.
Atoms,
Molecules,
and Ions
2. 2
Discovery of
Atomic
Structure
Atoms,
Molecules,
and Ions
J.J. Thomson’s Experiment
• Using a cathode ray tube Thomson found
that passing an electric current makes a
beam of radiation appear to move from
the negative to the positive end
• By adding an electric field to the cathode
ray tube, he found that the beam
deflected towards the positive field so he
deduces that the components of the
beam were negative
Atoms,
Molecules,
and Ions
Cathode Rays & Electrons
• Radiation was passed through a tube from
the cathode (negative end) to the anode
(positive end)
 The radiation would fluoresce (give off light)
 When an electric/magnetic field was applied, the
cathode rays deflected in a manner consistent with
a stream of negative particles
Since the gas was known to be neutral (having no
charge) he reasoned that there must also be
positively charged particles in the atom (but he
Atoms,
could never find them).
Molecules,
(Old TV sets were cathode ray tubes)
and Ions
The Electron
Cathode
Ray
• Streams of negatively charged particles were
found to emanate from cathode tubes.
• J. J. Thomson is credited with their discovery
(1897).
• Thomson measured the charge/mass ratio of Atoms,
Molecules,
8
the electron to be 1.76  10 coulombs/g.
and Ions
Millikan’s Oil Drop Experiment
• Robert Millikan put a charge onto a tiny drop of oil and
dropped it into an electric field
• When the field was manipulated, the oil drop’s free fall
was affected
• By measuring how strong an applied electric field had
to be in order to stop the oil drop in mid air, he was
able to work out the mass of the drop
• He could calculate the force of gravity on one drop
and thus the electric charge that the drop must have.
• By varying the charge on different drops, he noticed
that the charge was always a multiple of
Atoms,
-1.6 x 10-19 C, the charge on a single electron
Molecules,
and Ions
Millikan’s Oil Drop Experiment
Millikan’s
Oil Drop
Experiment
• Once the charge/mass ratio of the electron was
known, determination of either the charge or
Atoms,
the mass of an electron would yield the other. Molecules,
and Ions
Radioactivity:
• The spontaneous emission of radiation
by an atom.
• First observed by Henri Becquerel.
• Also studied by Marie and Pierre Curie.
Atoms,
Molecules,
and Ions
Radioactivity
• Three types of radiation were discovered by
Ernest Rutherford:
  [alpha] particles (positive charge)
  [beta] particles are high speed e– (neg. charge)
  [gamma] rays are unaffected by electric field/charge
Atoms,
Molecules,
and Ions
“Plum Pudding” Model
• In 1897, the English scientist J.J.
Thomson provided the first hint that
an atom is made of even smaller
particles.
• Thomson concluded that the
negative charges came from within
the atom.
• A particle smaller than an atom had
to exist… the atom was divisible!
• Positive sphere of matter with
negative electrons imbedded in it
Atoms,
Molecules,
and Ions
Rutherford’s Gold Foil Experiment
• 1909 - Ernest Rutherford shot  [alpha] particles at
a thin sheet of gold foil and observed the pattern
of scatter of the particles.
Gold Foil
Experiment
Atoms,
Molecules,
and Ions
Results of the Gold Foil Experiment
• Top: Expected results: alpha
particles (+) should pass through the
plum pudding model of the atom
undisturbed.
• Bottom:
Observed results:
Some of the particles
were deflected at
large angles,
indicating a small,
concentrated positive
charge.
Atoms,
Molecules,
and Ions
The Nuclear Atom
• Rutherford predicted that most of the
volume of the Au atoms was open,
empty space. The “plum pudding”
model could not be correct.
• Rutherford concluded that an atom
must have a small, dense, positively
charged center (the nucleus) that
repelled his positive particles
• He also concluded that the e–
were outside of the nucleus
Atoms,
Molecules,
and Ions
Other Subatomic Particles
• Protons
discovered by Rutherford in 1919.
• Neutrons
discovered by James Chadwick in 1932.
Atoms,
Molecules,
and Ions
2. 3
The Modern
View of Atomic
Structure
Atoms,
Molecules,
and Ions
Subatomic Particles
• Protons and electrons
the only particles that have a charge.
• Protons and neutrons
 have essentially the same mass.
• The mass of an electron is so small that we
ignore it (p+ and no are 1800x mass of e–).
Inside the Atom
• Protons and neutrons reside inside the nucleus of
the atom
Though the nucleus is extremely small, it
contains nearly all the mass of the atom
http://ed.ted.com/lessons/just-how-small-is-an-atom
Atomic Mass & Size
• Because atoms have such tiny masses (heaviest
is 4.0 x 10-22 g) we use the atomic mass unit
(amu)
1 amu = 1.66 x10-24 g
• Atoms are very small so another common unit of
length used is the Angstrom (Å)
diameters between 1x10-10 m and 5x10-10 m
= 100-500 pm
=1–5Å
Atomic Numbers, Mass
Numbers, and Isotopes
Atomic
Number (Z)
• the # of protons an atom contains
• is unique to each element
• Thus, # of p+ tells the identity of the atom
• Because atoms are neutral, every atom
has an equal number of protons and
electrons
Mass
Number (A)
• Mass # = protons + neutrons
• Mass of atoms in amu
• Atoms of an element can differ in the # of
neutrons they contain
As a result, atoms of the same element
can have different masses
Practice Counting
+
p,
o
n,
and
–
e
• p+ and e– = 15
• no = 31-15 = 16
• p+ and e– = 26
• no = 56 -26 = 30
• p+ and e– = 10
• no = 20-10 = 10
Isotopes:
• Atoms of the same element (thus same #
protons) with different masses and
different #s of neutrons.
• Standard notation:
11
C
6
12
C
6
13
C
6
14
C
6
• Hyphenated notation:
Carbon -12, Carbon -14, etc
Atoms,
Molecules,
and Ions
More Practice
• How many protons, neutrons, and electrons
are in the following?
1)108Ag
p+ = 47; no = 108-47 = 61; e– = 47
2)Nickel-60
3)209Pb
p+ = 28; no = 60-28 = 32; e– = 28
p+ = 82; no = 209-82 = 127; e– = 82
4)Silicon-30
p+ = 14; no = 30-14 = 16; e– = 14
2. 4
Atomic
Weights
Atoms,
Molecules,
and Ions
Atomic Mass
• Carbon-12 is the standard for
atomic masses of elements:
24
Cr
51.996
• It was assigned an exact mass of 12.00 amu
• All other atomic masses were determined in
comparison to the mass of Carbon-12
• The periodic table gives the masses of an
“average” atom of each element
must be an avg. because there are multiple
isotopes of each element present in nature
known as the average atomic mass
Average Atomic Mass
• Avg. atomic mass is determined by
 mass of each isotope AND
 each isotope’s relative abundance (how much
of that isotope is found in nature)
That means that it is a weighted average of all
the naturally occurring isotopes of that element.
Thus, avg. atomic mass is not a whole #
• Time for some math!
 *Notes Handout – Weighted Average*
Atoms,
Molecules,
and Ions
Weighted Average Example
• On the first day of school, Maggie’s Geometry teacher
told the class that all grades would be calculated by
the weighted average system seen below:
 Participation: 15%
Quizzes: 25%
Homework: 20%
Tests: 40%
• At the end of the MP1, Maggie’s not sure if she is
passing the class. She needs to know what her overall
weigthed class average is. Below are Maggie’s
average grades during MP1. Let’s help her!
 Maggie’s MP1 Averages
Participation: 70
Quizzes: 72
Homework: 74
Tests: 66
Atoms,
Molecules,
and Ions
Avg. Atomic Mass Calculations
- application of a weighted average -
Avg. atomic mass (AAM) =
[(rel, abundance isotope A) x (mass isotope A)]
+ [(rel. abundance isotope B) x (mass isotope B)]
+ [(rel. abundance isotope C) x (mass isotope C)]
…etc
*** Relative Abundance = % abundance ***
100
Atoms,
Molecules,
and Ions
Avg. Atomic Mass Practice
• Sulfur has five naturally occurring isotopes: 95.0%
of sulfur exists as Sulfur-32 (31.97207 amu)
0.76% exists as Sulfur-33 (32.97146 amu), 4.22%
exists as Sulfur-34 (33.96786 amu), 0.006% exists
as Sulfur-35 (34.9690 amu), and 0.014% exists as
Sulfur-36 (35.96709 amu). What is the average
atomic mass of Sulfur?
2. 5
The Periodic
Table
Atoms,
Molecules,
and Ions
The Periodic Table of Elements
Atoms,
Molecules,
and Ions
Periodic Table
• Period/Series =
horizontal rows (#1-7,
lanthanides, actinides)
• Families/Groups =
vertical columns(#1-18,
special names for some)
Elements in the same
family have similar
chemical and physical
properties.
Atoms,
Molecules,
and Ions
Periodicity & Periodic Law
Looking at the elements in
the periodic table, there is
a reoccurring pattern of
chemical and physical
properties at intervals.
Periodic Table
Nonmetals are on
the right side of the
periodic table (with
the exception of H).
Atoms,
Molecules,
and Ions
Periodic Table
Metalloids border
the staircase line.
Atoms,
Molecules,
and Ions
Periodic Table
Metals are on the
left side of the chart.
Atoms,
Molecules,
and Ions
1A – Alkali Metals
(form +1 charge)
Atoms,
Molecules,
and Ions
2A – Alkali Earth Metals
(form +2 charge)
Atoms,
Molecules,
and Ions
1B – Coinage Metals
Atoms,
Molecules,
and Ions
6A – Chalcogens
(form –2 charge)
Atoms,
Molecules,
and Ions
7A – Halogens
(form –1 charge)
Atoms,
Molecules,
and Ions
8A – Noble Gases
(unreactive & very stable)
Atoms,
Molecules,
and Ions
Lanthanide Series
Atoms,
Molecules,
and Ions
Actinide Series
(radioactive; many man-made)
Atoms,
Molecules,
and Ions
2. 6
Molecules &
Molecular
Compounds
Atoms,
Molecules,
and Ions
Chemical Formulas
• The subscript to the right of
an elemental symbol tells the
number of atoms of that
element present in one
molecule of the compound.
 Ex: NH4 = 1 nitrogen atom
& 4 hydrogen atoms
Atoms,
Molecules,
and Ions
Molecular Compounds
• Compounds which are
composed of molecules,
contain more than 1 type of
atom, and are almost
always composed of only
nonmetals
 Ex: Figure 2.20 (left)
• The composition of each
compound is given by its
chemical formula
Atoms,
Molecules,
and Ions
Diatomic Molecules
• Many elements in nature are found most
commonly in their molecular form
 Diatomics are two of the same atom bound together
 ONLY the elements: Br2-I2-N2-Cl2-H2-O2-F2
 When we speak of hydrogen, it is automatically H2
unless specified otherwise (same for other diatomics)
Atoms,
Molecules,
and Ions
Types of Formulas
• Molecular formulas give the exact
number of atoms of each element in a
compound.
Ex: H2O2 or C4H10
• Empirical formulas give the lowest
whole-number ratio of atoms of each
element in a compound.
Ex: HO or C2H5
Atoms,
Molecules,
and Ions
Picturing Atoms
• Chemical formulas only
show what elements are
present in an atom, not how
they are bonded together
 CH4
• Structural formulas show
the order in which atoms are
bonded.
.
Atoms,
Molecules,
and Ions
Molecular or Empirical Formulas?
C6H6
Molecular
C8H18
Molecular
WO2
Empirical
C3H6O2
Empirical
X39Y13
Molecular
Atoms,
Molecules,
and Ions
2. 7
Ions & Ionic
Compounds
Atoms,
Molecules,
and Ions
+
Cations
Ions
–
Anions
• When atoms lose or gain electrons,
they become ions.
 Cations are positively charged
 Anions are negatively charged
Atoms,
Molecules,
and Ions
Ions
• The nucleus of an atom is left unchanged by
a chemical rxn, but atoms will readily loose
or gain e–
 Gaining e– creates negative anions
 Losing e– creates positive cations
Atoms,
Molecules,
and Ions
Atoms to Ions…
• The ion has a –1 charge because it has one
more electron than its number of protons
(18 e– vs.17 p+)
Atoms,
Molecules,
and Ions
Counting Subatomic
Particles in Ions
How many p+ and e– are in Se2– ?
• # Protons ALWAYS = atomic #
 … 34 p+
• # Electrons depends on IONIC CHARGE
 2– charge means two more electrons than
protons (Se atom gained two e–)
 ... 34 p+ + 2 = 36 e–
Atoms,
Molecules,
and Ions
Counting Subatomic
Particles in Ions
How many p+ and e– are in K+ ?
• # Protons ALWAYS = atomic #
 … 19 p+
• # Electrons depends on ionic charge
 +1 charge means one less electron than
protons (K atom lost one e–)
 ... 19 p+ - 1 = 18 e–
Atoms,
Molecules,
and Ions
Practice with Monatomic Ions
Ions
# Protons
# Electrons
N3-
7
10
Br -
35
36
Sr2+
38
36
Li+
3
2
S2-
16
18
Atoms,
Molecules,
and Ions
Polyatomic Ions
• Ions which have two or more different
atoms joined as a molecule
 will have a net negative or positive charge.
Ex: NO3– = nitrate ion
SO42 – = sulfate ion
Atoms,
Molecules,
and Ions
Predicting Ionic Charges
• The goal of any atom is...
 To gain/lose e– so that it will have the
same # of e– as the noble gas nearest to it
 All atoms want to have stable
arrangements like the Nobel gases!
Atoms,
Molecules,
and Ions
Predicting Ionic Charges
- using the periodic table to predict charges -
Atoms,
Molecules,
and Ions
Common Cations
Atoms,
Molecules,
and Ions
Common Anions
Atoms,
Molecules,
and Ions
Ionic Bonds
• Transfer of electrons from one substance to
another
• Ionic compounds (such as NaCl) are generally
formed between metals and nonmetals.
 Ex: Neutral Na atom + Neutral Cl atom …
Na gives e– to Cl  Na+ and Cl–
Because opposite
charges attract,
Na+ and Cl– bind
together to form a
neutral ionic
compound, NaCl.
Atoms,
Molecules,
and Ions
Practice - Ionic vs. Molecular
• Which of the following are ionic compound and
which are molecular compounds?
 N2O
nm + nm = molecular
 Na2O
m + nm = ionic
 CaCl2
ionic
 SF4
molecular
 CrCl2
ionic
Atoms,
Molecules,
and Ions
Ionic Compounds
• Ionic Compound
cation + anion
metal + nonmetal
Neutral compound (no net charge)
• Total + charge and – charge MUST be equal
in order to cancel with each other
Atoms,
Molecules,
and Ions
Criss-Cross Method
• Method for writing empirical ionic formulas
Step 1: Write the cation down first, followed by
the anion right next to it
Step 2: The charge of the cation becomes the
subscript of the anion
Step 3: The charge of the anion becomes the
subscript of the cation
Step 4: If the subscripts are not in the lowest
whole-number ratio, divide by the greatest
common factor (because ionic compounds Atoms,
Molecules,
will only have empirical formulas)
and Ions
Criss-Cross Method Example
• Write the empirical formula for the
compounds formed by the following ions:
 Mg2+ and N3–
Atoms,
Molecules,
and Ions
Formulas w/Polyatomics
• Ionic bonds also form with polyatomic ions
• Treat the polyatomic as a single particle
• After the criss-cross technique, if the subscript
for the entire polyatomic ion is 2 or more,
surround the the polyatomic ion in parentheses
Ex: Ammonium Sulfide
NH4+ and S2–
(NH4)2S
• Do not change the subscripts of the
polyatomic ion itself!!!
Atoms,
Molecules,
and Ions
Criss-Cross Method Practice
• Write the empirical formula for the
compounds formed by the following ions:
1) Na+ and PO43–
Na3PO4
2) Zn2+ and SO42–
ZnSO4
3) Fe3+ and O2–
Fe2O3
4) Al3+ and CO32–
Al2(CO3)3
Atoms,
Molecules,
and Ions
2. 8
Naming
Inorganic
Compounds
Atoms,
Molecules,
and Ions
Chemical Nomenclature
• Organic Compounds
Contain the elements C, O, H, N, S
Have their own naming system
• Inorganic Compounds
Ionic Compounds
Molecular Compounds
Acids
Atoms,
Molecules,
and Ions
Naming Ionic
Compounds
Atoms,
Molecules,
and Ions
Naming Ionic Cations
• Monatomic cations from metal atoms
have the same name as the metal
• Na+
• Zn2+
sodium ion
zinc ion
The following only form only cation:
• group 1A, 2A, Al3+, Ag+, Zn2+
Atoms,
Molecules,
and Ions
Naming Ionic Cations
• Cations from transition metals
Transition metals can form different cations
charge is indicated by a roman numeral in
parentheses after the name of the metal
• Fe2+
• Fe3+
iron (II) ion
iron (III) ion
Older naming system – also distinguishes
differently charged ions:
• Fe2+
• Fe3+
Ferrous (lower charge)
Ferric (higher charge)
Atoms,
Molecules,
and Ions
Naming Ionic Anions
• Monatomic anions formed from nonmetals
named by replacing the ending with the
suffix –ide
•
•
•
•
O2– oxide ion
N3– nitride ion
S2– sulfide ion
some polyatomics end in –ide (OH–, CN–)
Atoms,
Molecules,
and Ions
Naming Ionic Anions
• Polyatomic anions containing oxygen (aka:
oxyanions) either end in:
 –ite (lower # oxygen) or –ate (higher # oxygen)
• SO32–
• SO42–
sulfite ion
sulfate ion
• NO22• NO32-
nitrite ion
nitrate ion
• Prefixes for extended oxyanions:
•
•
•
•
ClO– hypochlorite ion
ClO2– chlorite ion
ClO3– chlorate ion
ClO4– perchlorate ion
(fewest oxygens)
(2nd fewest oxygens)
(2nd most oxygens) Atoms,
Molecules,
(most oxygens)
and Ions
Inorganic Nomenclature - Basics
• Write the name of the cation.
• If the cation can have more than one
possible charge (transition metals),
write the charge as a roman numeral in
parentheses.
• If the anion is an element, change its
ending to -ide; if the anion is a
polyatomic ion, simply write the name of
the polyatomic ion.
Atoms,
Molecules,
and Ions
Practice Naming Ionic Compounds
Examples:
 NH4OH
• NH4+ + OH–
• ammonium
hydroxide
LiCl
• Li+ + Cl–
• lithium chloride
CuI2
• Cu2+ + I–
• Copper (II) Iodide
• Cupric Iodide
MgS
• Mg2+ + S2–
• Magnesium sulfide
Atoms,
Molecules,
and Ions
Practice Writing Ionic Formulas
Examples:
calcium carbonate
• Ca2+ + CO32–
• CaCO3
magnesium hydroxide
• Na2+ + OH–
• Mg(OH)2
ammonium nitrate
• NH4+ + NO3–
• NH4NO3
iron (III) carbonate
ferric carbonate
• Fe3+ + CO32–
• Fe2(CO3)3
Atoms,
Molecules,
and Ions
Naming Acids
Atoms,
Molecules,
and Ions
Some Acid Basics…
• Hydrogen (H+) containing compounds
• For now, the formula of an acid will have
hydrogen listed as the first element
• Consider an acid to be composed of an anion
connected to enough H+ ions to balance the
anion's charge
• Examples of acids:
 HCl  H+ and Cl–
 H2SO4  2H+ and SO4 2–
Atoms,
Molecules,
and Ions
Acid Nomenclature
• If the anion in the acid ends in -ide, change the
ending to -ic acid and add the prefix hydro- :
 HCl:
• H+ and Cl– (chloride ion)
• hydrochloric acid
 HBr:
• hydrobromic acid
 HI:
• hydroiodic acid
Atoms,
Molecules,
and Ions
Acid Nomenclature
• If the anion in the acid ends in -ite, change the
ending to -ous acid:
 HClO
• H+ and hypochlorite
• hypochlorous acid
 HClO2
• H+ and chlorite
• chlorous acid
Atoms,
Molecules,
and Ions
Acid Nomenclature
• If the anion in the acid ends in -ate, change the
ending to -ic acid:
 HClO3
• H+ and chlorate
• chloric acid
 HClO4
• H+ and perchlorate
• perchloric acid
 H2SO4
• H+ and sulphate
• sulfuric acid
Atoms,
Molecules,
and Ions
Atoms,
Molecules,
and Ions
Naming Binary
Molecular
Compounds
[aka: Covalent
Compounds]
Atoms,
Molecules,
and Ions
Binary Molecular Compounds
• Nonmetal + nonmetal
• The element farther to the
left in the periodic table (the
less electronegative atom) is
usually listed first.
• A prefix is used to denote the
number of atoms of each
element in the compound
• The prefix mono- is never
used on the first element
listed, however.)
Atoms,
Molecules,
and Ions
Binary Molecular Compounds
• The ending on the more
electronegative element
(furthest to the right in
periodic table) is changed
to -ide.
 CO2
• carbon dioxide
 CCl4
• carbon tetrachloride
Atoms,
Molecules,
and Ions
Binary Molecular Compounds
• If the prefix ends with a
or o and the name of
the element begins with
a vowel, the two
successive vowels are
often merged into one:
• N2O5
dinitrogen pentoxide
Atoms,
Molecules,
and Ions
Practice Naming Binary
Molecular Compounds
Examples:
 Cl2O
• dichlorine
monoxide
N2O4
• dinitrogen
tetroxide
NF3
• nitrogen
trifluoride
P4S10
• tetraphosphorous
decasulfide
Atoms,
Molecules,
and Ions
Practice Writing
Molecular Formulas
Examples:
phosphourous
triiodide
• PI3
dihygrogen monoxide
• H2O
diborane trioxide
• B2O3
iodine monobromide
• IBr
Atoms,
Molecules,
and Ions
2. 9
Simple
Organic
Compounds
Atoms,
Molecules,
and Ions
Organic Compounds
• Studies compounds which contain carbon
• Typically, organic compounds have carbon
bonded with H, O, S, and P
• Hydrocarbons
 contain only C and H
 In the simplest group, alkanes, each carbon is
bonded to four other atoms
 Although they are binary molecular compounds,
they are NOT named like the inorganics in
section 2.8.
Atoms,
Molecules,
and Ions
Facts about Naming
Organic Compounds
• For alkanes, the ending of the compound will
always be -ane
• Distinct organic prefixes are used to tell you
the # of C atoms in the compound
 Example prefixes:
• meth = 1
• eth = 2
• Prop = 3
etc…
Atoms,
Molecules,
and Ions
Alkanes
• These are the 3 simplest alkanes containing
1, 2 and 3 carbon atoms
Atoms,
Molecules,
and Ions
Derivatives of Alkanes
• Other organic compounds form when some of the H in
an alkane are replaced with functional groups.
 There are several different functional groups
 Ex: an alcohol is an alkane which had hydrogen replaced
with OH (not an ion so NOT the polyatomic hydroxide)
Atoms,
Molecules,
and Ions