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Chapter 2
Atoms, Molecules, and Ions
▶ A SECTION THROUGH A GEODE.
A geode is a mass of mineral matter
(often containing quartz) that
accumulates slowly within the shell of
a roughly spherical, hollow rock.
Eventually, perfectly formed crystals
may develop at a geode’s center. The
colors of a geode depend upon its
composition. Here, agate crystallized
out as the geode formed.
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What’s Ahead
Atoms, Molecules, and Ions
2.1 THE ATOMIC THEORY OF MATTER
We begin with a brief history of the notion of atoms—the smallest pieces of matter.
2.2 THE DISCOVERY OF ATOMIC STRUCTURE
We then look at some key experiments that led to the discovery of electrons and to the
nuclear model of the atom.
2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
We explore the modern theory of atomic structure, including the ideas of atomic
numbers, mass numbers, and isotopes.
2.4 ATOMIC WEIGHTS
We introduce the concept of atomic weights and how they relate to the masses of
individual atoms.
2.5 THE PERIODIC TABLE
We examine the organization of the periodic table, in which elements are put in order of
increasing atomic number and grouped by chemical similarity.
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What’s Ahead
Atoms, Molecules, and Ions
2.6 MOLECULES AND MOLECULAR COMPOUNDS
We discuss the assemblies of atoms called molecules and how their compositions are
represented by empirical and molecular formulas.
2.7 IONS AND IONIC COMPOUNDS
We learn that atoms can gain or lose electrons to form ions. We also look at how to use
the periodic table to predict the charges on ions and the empirical formulas of ionic
compounds.
2.8 NAMING INORGANIC COMPOUNDS
We consider the systematic way in which substances are named, called nomenclature,
and how this nomenclature is applied to inorganic compounds.
2.9 SOME SIMPLE ORGANIC COMPOUNDS
We introduce organic chemistry, the chemistry of the element carbon.
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Atoms, Molecules, and Ions
How Do We Explain?
• What makes diamonds transparent and hard, while table salt
is brittle and dissolves in water?
• Why does paper burn, and why does water quench fires?
• Where does the beautiful colors of flowers come from?
• The structure and behavior of atoms are key to understanding
the properties of matter.
• In this chapter we examine the basic structure of atoms and
discuss the formation of molecules and ions, thereby
providing a foundation for exploring chemistry more deeply in
later chapters.
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2.1 THE ATOMIC THEORY OF MATTER
History
• Democritus and Greek philosophers (BC 400).
– The material world must be made up of tiny indivisible particles.
– Atomos: indivisible or uncuttable.
• Plato and Aristotle.
– There can be no ultimately indivisible particles.
– The “atomic” view of matter faded for many centuries.
• The notion of atoms reemerged in Europe during the 17th
century.
– Chemists learned to measure the amounts of elements that
reacted with one another to form new substances.
– The ground was laid for an atomic theory that linked the idea of
elements with the idea of atoms.
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2.1 THE ATOMIC THEORY OF MATTER
Dalton’s Atomic Theory
• The theory came from
the work of John Dalton
during the period from
1803 to 1807.
• Atoms are the
fundamental building
blocks of matter.
• The theory was based
on the four postulates
given in the figure.
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2.1 THE ATOMIC THEORY OF MATTER
Dalton’s Postulates
1) Each element is
composed of
extremely small
particles called
atoms.
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2.1 THE ATOMIC THEORY OF MATTER
Dalton’s Postulates
2) All atoms of a given
element are identical
to one another in
mass and other
properties, but the
atoms of one element
are different from the
atoms of all other
elements.
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2.1 THE ATOMIC THEORY OF MATTER
Dalton’s Postulates
3) Atoms of an element
are not changed into
atoms of a different
element by chemical
reactions; atoms are
neither created nor
destroyed in chemical
reactions.
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2.1 THE ATOMIC THEORY OF MATTER
Dalton’s Postulates
4) Atoms of more than
one element combine
to form compounds;
a given compound
always has the same
relative number and
kind of atoms.
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2.1 THE ATOMIC THEORY OF MATTER
Dalton’s Postulates
• Dalton’s theory explains several simple laws of chemical
combination.
• The law of constant composition.
– In a given compound, the relative numbers and kinds of atoms are
constant.
• The law of conservation of mass (matter).
– The total mass of materials present after a chemical reaction is the
same as the total mass present before the reaction.
• The law of multiple proportions.
– If two elements A and B combine to form more than one compound,
the masses of B that can combine with a given mass of A are in the
ratio of small whole numbers (H2O and H2O2).
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Atomic Structure
• Dalton had no direct evidence for the existence of atoms.
• Scientists have developed methods for more detailed
probing of the nature of matter.
• Many discoveries led to the fact that the atom itself was
made up of smaller (subatomic) particles.
• We can measure the properties of individual atoms and
even provide images of them.
Figure 2.2 An image of the surface of silicon.
The image was obtained by a technique called
scanning tunneling microscopy. The color was
added to the image by computer to help
distinguish its features. Each red sphere is a
silicon atom.
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
The Electron (Cathode Rays)
• Streams of negatively charged particles were found to
emanate from cathode tubes, causing fluorescence.
• J. J. Thomson is credited with their discovery (1897).
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
The Electron
Thomson measured the charge/mass ratio of the
electron to be 1.76 × 108 coulombs/gram (C/g).
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Millikan Oil-Drop Experiment
• In 1909, Robert Millikan (1868–1953) succeeded in
measuring the charge of an electron (1.602 × 10–19 C) by
performing the oil-drop experiment.
Figure 2.5 Millikan’s oil-drop
experiment used to measure the
charge of the electron. Small
drops of oil were allowed to fall
between electrically charged plates.
The drops picked up extra electrons
as a result of irradiation by X-rays
and so became negatively charged.
Millikan measured how varying the
voltage between the plates affected
the rate of fall. From these data he
calculated the negative charge on
the drops. Because the charge on
any drop was always some integral
multiple of 1.602 × 10–19 C, Millikan
deduced this value to be the charge
of a single electron.
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Millikan Oil-Drop Experiment
• He then calculated the mass of the electron by using his
experimental value for the charge (1.602 × 10–19 C) and
Thomson’s charge-to-mass ratio (1.76 × 108 C/g):
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Radioactivity
• Radioactivity is the spontaneous emission of highenergy radiation by an atom.
• It was first observed by Henri Becquerel.
• Marie and Pierre Curie also studied it.
• Its discovery showed that the atom had more
subatomic particles and energy associated with it.
Figure 2.6 Marie Sklodowska Curie (1867–1934). When
Marie Curie presented her doctoral thesis, it was described
as the greatest single contribution of any doctoral thesis in
the history of science. In 1903 Henri Becquerel, Maire Curie,
and her husband, Pierre, were jointly awarded the Nobel
Prize in Physics for their pioneering work on radioactivity (a
term she introduced). In 1911 Marie Curie won a second
Nobel Prize, this time in chemistry for her discovery of the
elements polonium and radium.
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Radioactivity
• Three types of radiation were discovered by Ernest Rutherford:
– α particles: 2+ charge.
• A mass about 7400 times that of an electron.
– β particles: 1– charge.
• A mass of an electron.
– γ rays: No charge, high energy radiation.
• Not consist of particles.
Figure 2.7 Behavior of alpha (α), beta (β), and gamma (γ) rays in an electric field.
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
The Atom, circa 1900
• The prevailing theory was
that of the “plum pudding”
model, put forward by
Thomson.
• It featured a positive sphere
of matter with negative
electrons imbedded in it.
Figure 2.8 J. J. Thomson’s plum-pudding
model of the atom. Ernest Rutherford proved
this model wrong.
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
Discovery of the Nucleus
• Ernest Rutherford shot α particles at a thin sheet of gold
foil and observed the pattern of scatter of the particles.
Figure 2.9 Rutherford’s α-scattering
experiment.
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2.2 THE DISCOVERY OF ATOMIC STRUCTURE
The Nuclear Atom
• Since some particles were deflected
at large angles, Thompson’s model
could not be correct.
Figure 2.10 The structure of the atom. A cloud
of rapidly moving electrons occupies most of the
volume of the atom. The nucleus occupies a tiny
region at the center of the atom and is composed
of the protons and neutrons. The nucleus
contains virtually all the mass of the atom.
• Rutherford postulated a very
small, dense nucleus with the
electrons around the outside
of the atom.
• Most of the volume of the
atom is empty space.
• Protons were discovered by
Rutherford in 1919.
• Neutrons were discovered by
James Chadwick in 1932.
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2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
Subatomic Particles
• Protons (+1) and electrons (–1) are the only particles
that have a charge; neutrons are neutral.
– Every atom has an equal number of protons and electrons, so
atoms have no net charge.
• Protons and neutrons have essentially the same mass.
• The mass of an electron is so small when compared to
that of a proton or a neutron. We ignore it.
• Protons and neutrons are found in the nucleus;
electrons travel around the nucleus.
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2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
Atomic Mass
• Atoms are very tiny, so a mass scale on the atomic level
is used to express the mass of an atom or molecule.
• An atomic mass unit (amu) is the base unit.
• 1 amu = 1.66054 × 10–24 g.
• The SI symbol used for the atomic mass unit is u.
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2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
Symbols of Elements
• What makes an atom of one element different from an atom
of another element: A characteristic number of protons of
the atoms.
• Atomic number: The number of protons in the nucleus. It
is written as a subscript BEFORE the symbol.
• Mass number: The total number of protons and neutrons in
the atom. It is written as a superscript BEFORE the symbol.
• Elements are represented by a one or two letter symbol.
This is the symbol for carbon.
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2.3 THE MODERN VIEW OF ATOMIC STRUCTURE
Isotopes
• Isotopes: Atoms of the same element with different masses.
• Isotopes have different numbers of neutrons.
• The symbol
(read “carbon twelve,” carbon-12) can be
represented simply as
.
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Atomic Mass
2.4 ATOMIC WEIGHTS
• Atomic and molecular masses can be measured with great
accuracy using a mass spectrometer.
Figure 2.11 A mass spectrometer. Cl atoms
are introduced at A and are ionized to form Cl+
ions, which are then directed through a magnetic
field. The paths of the ions of the two Cl isotopes
diverge as they pass through the field.
Figure 2.12 Mass spectrum of
atomic chlorine. The fractional
abundances of the isotopes 35Cl and
37Cl are indicated by the relative signal
intensities of the beams reaching the
detector of the mass spectrometer.
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Atomic Mass
2.4 ATOMIC WEIGHTS
• Atomic mass unit (amu).
– Defined by assigning a mass of exactly 12 amu to an atom of 12C.
•
•
1H:
1.6735 × 10–24 g = 1.0078 amu.
16O: 2.6560 × 10–23 g = 15.9949 amu.
– 1 amu = 1.66054 × 10–24 g and 1 g = 6.02214 × 1023 amu.
• The average atomic mass of an element = The element’s
atomic weight.
– Most elements occur in nature as mixtures of isotopes.
– Calculated from the isotopes of an element weighted by their
relative abundances:
– C: 98.93% 12C and 1.07% 13C.
Atomic weight = (0.9893)(12 amu) + (0.0107)(13.00335 amu)
= 12.01 amu.
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Periodic Table
2.5 THE PERIODIC TABLE
• The periodic table is the most significant tool that chemists
use for organizing and remembering chemical facts.
Figure 2.14 Periodic table of the elements.
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Periodic Table
2.5 THE PERIODIC TABLE
• The periodic table
is a systematic
catalog of the
elements.
• Elements are
arranged in order
of atomic number.
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Periodic Table
2.5 THE PERIODIC TABLE
• The rows on the periodic chart are periods.
• Columns are groups.
• Elements in the same group have similar chemical properties.
– “Coinage metals” (Cu, Ag, and Au) belong to group 11 (or 1B): Less
reactive than most metals.
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Periodicity
2.5 THE PERIODIC TABLE
• Many elements show very strong similarities to one another.
• When one looks at the chemical properties of elements, one
notices a repeating pattern of reactivities.
Figure 2.13 Arranging elements by atomic number reveals a periodic
pattern of properties. This pattern is the basis of the periodic table.
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Groups
2.5 THE PERIODIC TABLE
• These five groups are known by their names.
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2.5 THE PERIODIC TABLE
Metals, Nonmetals, and Metalloids
• Nonmetals are on
the right side of the
periodic table (with
the exception of H).
• Metalloids border
the stair-step line
(with the exception
of Al, Po, and At).
• Metals are on the
left side of the chart.
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2.5 THE PERIODIC TABLE
Metals, Nonmetals, and Metalloids
• Nonmetals generally differ from the metals in appearance
and in other physical properties.
• A metalloid is a chemical element with properties that are
in-between or a mixture of those of metals and nonmetals.
Figure 2.15 Examples of metals and nonmetals.
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2.6 MOLECULES AND MOLECULAR COMPOUNDS
Molecules
• Only the noble-gas elements
are normally found in nature as
isolated atoms.
• A molecule is an assembly of
two or more atoms tightly bound
together.
• Molecules behave in many
ways as a single, distinct object.
Figure 2.18 Molecular models. Notice how
the chemical formulas of these simple
molecules correspond to their compositions.
• The subscript to the right of the
symbol of an element tells the
number of atoms of that
element in one molecule.
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2.6 MOLECULES AND MOLECULAR COMPOUNDS
Molecules and Chemical Formulas
• Many elements are found in nature in molecular form.
– Two different molecular forms of oxygen.
• O2: A diatomic molecule, essential for life, odorless.
• O3: A triatomic molecule, toxic, pungent smell.
• Molecular compounds are composed of molecules
and almost always contain only nonmetals.
• These seven elements occur naturally as molecules
containing two atoms (diatomic molecules): Hydrogen,
nitrogen, oxygen, fluorine, chlorine, bromine, and iodine.
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2.6 MOLECULES AND MOLECULAR COMPOUNDS
Molecular and Empirical Formulas
• Molecular formulas
give the exact number of
atoms of each element in
a compound.
–
–
–
–
H 2O
H 2 O2
C 2H 4
C6H12O 6
• Empirical formulas give
the lowest whole-number
ratio of atoms of each
element in a compound.
–
–
–
–
H 2O
HO
CH2
CH2O
• Why do we need empirical formulas?
– Certain common methods of analyzing substances
(elemental analysis) lead to the empirical formula only.
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2.6 MOLECULES AND MOLECULAR COMPOUNDS
Picturing Molecules
• A structural formulas show the
order in which atoms are attached.
They do NOT depict the threedimensional shape of molecules.
• A perspective drawings show the
three-dimensional order of the atoms
in a compound. These are also
demonstrated using models.
• A Ball-and-stick models shows the
accurate angles between bonds.
• A space-filling model shows the
relative sizes of the atoms.
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2.7 IONS AND IONIC COMPOUNDS
Cations and Anions
• When atoms lose or gain electrons, they become ions.
– Cations are positive and are formed by elements on the left side
of the periodic chart (metal atoms).
– Anions are negative and are formed by elements on the right
side of the periodic chart (nonmetal atoms).
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2.7 IONS AND IONIC COMPOUNDS
Predicting Ionic Charges
• Many atoms gain or lose electrons to end up with the same
number of electrons as the noble gas closest to them in the
periodic table.
– Na  Na+ (the same number of electrons as in Ne).
– Cl  Cl– (the same number of electrons as in Ar).
• The periodic table is very useful for remembering ionic charges.
–
–
–
–
–
The group 1A elements (alkali metals) form 1+ ions.
The group 2A elements (alkaline earths) form 2+ ions.
The group 7A elements (halogens) form 1– ions.
The group 6A elements form 2– ions.
Many of the other groups do not lend themselves to such simple rules.
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2.7 IONS AND IONIC COMPOUNDS
Ionic Compounds
• A compound made up of cations and anions.
• Ionic compounds (such as NaCl) are generally formed
between metals and nonmetals.
– Cf) Molecular compounds (such as H2O) are generally composed
of nonmetals only.
• Electrons are transferred from the metal to the nonmetal.
The oppositely charged ions attract each other. Only
empirical formulas are written.
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2.7 IONS AND IONIC COMPOUNDS
Writing Empirical Formulas
• Because compounds are electrically neutral, one can
determine the formula of a compound this way:
– The charge on the cation becomes the subscript on the anion.
– The charge on the anion becomes the subscript on the cation.
– If these subscripts are not in the lowest whole-number ratio,
divide them by the greatest common factor.
• The ionic compound formed from Mg2+ and N3–:
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2.7 IONS AND IONIC COMPOUNDS
ELEMENTS REQUIRED BY LIVING ORGANISMS
Figure 2.20 Elements essential to life.
• More than 97% of the mass of most organisms:
O, C, H, N, P, and S.
• At least 70% of the mass of most cells: H2O.
• C is the most prevalent element by mass.
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
• Cations.
– Cations formed from metal atoms have the same name as the metal:
– If a metal can form cations with different charges, the positive charge
is indicated by a Roman numeral in parentheses following the name
of the metal:
• Ions of the same element that have different charges exhibit different
properties, such as different colors.
Figure 2.21 Different ions of the same element have
different properties. Both substances shown are
compounds of iron. The substance on the left is Fe3O4,
which contains Fe2+ and Fe3+ ions. The substance on
the right is Fe2O3, which contains Fe3+ ions.
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
• Metals that form only one cation: Metals of group 1A and 2A, Al3+, Ag+,
Zn2+.
• An older method uses the endings -ous and -ic added to the root of the
element’s Latin name:
– Cations formed from nonmetal atoms have names that end in –ium:
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
• Anions.
– The names of monatomic anions are formed by replacing the ending
of the name of the element with -ide:
– A few polyatomic anions also have names ending in -ide:
– Polyatomic anions containing oxygen have names ending in either ate or -ite and are called oxyanions.
• The -ate is used for the most common or representative oxyanion of an
element.
• The -ite is used for an oxyanion that has the same charge but one O
atom fewer:
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
• Prefixes are used when the series of oxyanions of an element extends to
four members, as with the halogens. The prefix per- indicates one more
O atom than the oxyanion ending in -ate; hypo- indicates one O atom
fewer than the oxyanion ending in -ite:
Figure 2.22 Procedure for naming anions. The first part of the element’s
name, such as “chlor” for chlorine or “sulf” for sulfur, goes in the blank.
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
Figure 2.23 Common oxyanions. The composition and charges of
common oxyanions are related to their location in the periodic table.
– Anions derived by adding H+ to an oxyanion are named by adding as
a prefix the word hydrogen or dihydrogen, as appropriate:
• An older method uses the prefix bi-: The HCO3– ion is commonly called
the bicarbonate ion, and HSO4– is sometimes called the bisulfate ion.
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Ionic Compounds
• Ionic compounds.
– Names of ionic compounds consist of the cation name followed by
the anion name:
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2.8 NAMING INORGANIC COMPOUNDS
Names and Formulas of Acids
• If the anion in the acid ends in -ide, change the ending to ic acid and add the prefix hydro-.
• If the anion in the acid ends in -ate, change the ending to ic acid. If the anion in the acid ends in -ite, change the
ending to -ous acid.
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2.8 NAMING INORGANIC COMPOUNDS
Nomenclature of Binary Molecular Compounds
• The procedures used for naming binary (two-element)
molecular compounds.
1. The name of the element farther to the left in the periodic table (less
electronegative atom, closest to the metals) is usually written first.
• An exception occurs when the compound contains oxygen and chlorine, bromine,
or iodine (any halogen except fluorine), in which case oxygen is written last.
2. If both elements are in the same group, the one closer to the bottom
of the tables is named first.
3. The name of the second element is given an -ide ending.
• CO2: carbon dioxide, CCl4: carbon tetrachloride.
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2.8 NAMING INORGANIC COMPOUNDS
Nomenclature of Binary Molecular Compounds
4. Greek prefixes are used to indicate the number of atoms of each
element. The prefix mono- is never used with the first element.
When the prefix ends in a or o and the name of the second element
begins with a vowel, the a or o of the prefix is often dropped.
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2.9 SOME SIMPLE ORGANIC COMPOUNDS
Nomenclature of Organic Compounds
• Organic chemistry.
– The study of carbon.
– Organic chemistry has its own system of nomenclature.
– Organic compounds: Compounds that contain carbon and hydrogen.
• Alkanes (the simplest hydrocarbons).
– Compounds containing only carbon and hydrogen.
– The first part of the names just listed correspond to the number of
carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
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2.9 SOME SIMPLE ORGANIC COMPOUNDS
Some derivatives of Alkanes
• When a hydrogen in an alkane is replaced with something
else (a functional group, like -OH), the name is derived
from the name of the alkane.
• The ending denotes the type of compound.
– An alcohol is obtained by replacing an H atom with a -OH group.
– An alcohol ends in -ol.
• Compounds with the same molecular
formula but different arrangements of
atoms are called isomers.
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Chapter 2. Homework
Exercises 2.3
2.13
2.17
2.22
2.31
2.39
2.41
2.50
2.61
2.70
2.79
2.85
2.111
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