Chemistry
Third Edition
Julia Burdge
Lecture PowerPoints
Chapter 2
Atoms, Molecules, and
Ions
Copyright © 2012, The McGraw-Hill Compaies, Inc. Permission required for reproduction or display.
CHAPTER
2.1
2.2
2.3
2.4
2.5
2.6
2.7
2
Atoms, Molecules, and Ions
The Atomic Theory
The Structure of the Atom
Atomic Number, Mass Number, and Isotopes
The Periodic Table
The Atomic Mass Scale and Average Atomic Mass
Molecules and Molecular Compounds
Ions and Ionic Compounds
2
2.1
The Atomic Theory
Topics
The Atomic Theory
3
2.1
The Atomic Theory
The Atomic Theory
Dalton’s Atomic Theory
1. Elements are composed of extremely small particles called
atoms. All atoms of a given element are identical, having
the same size, mass, and chemical properties. The atoms
of one element are different from the atoms of all other
elements.
2. Compounds are composed of atoms of more than one
element. In any given compound, the same types of atoms
are always present in the same relative numbers.
4
2.1
The Atomic Theory
The Atomic Theory
Dalton’s Atomic Theory
3. A chemical reaction rearranges atoms in chemical
compounds; it does not create or destroy them.
5
2.1
The Atomic Theory
The Atomic Theory
Proust’s law of definite proportions :
6
2.1
The Atomic Theory
The Atomic Theory
The law of multiple proportions :
7
2.1
The Atomic Theory
The Atomic Theory
The law of multiple proportions :
8
SAMPLE PROBLEM
2.1
(a) Both water (H2O) and hydrogen peroxide (H2O2) are
composed of hydrogen and oxygen. When water is
decomposed into its constituent elements, it produces
0.125 g hydrogen for every gram of oxygen. When
hydrogen peroxide is decomposed, it produces 0.063 g
hydrogen for every gram of oxygen.
Determine the whole number ratio of g H : 1.00 g O in
water to g H : 1.00 g O in hydrogen peroxide to show how
these data illustrate the law of multiple proportions.
9
SAMPLE PROBLEM
2.1
Solution
(a)
10
SAMPLE PROBLEM
(b)
2.1
Sulfur and oxygen can combine to form several
compounds including sulfur dioxide (SO2) and sulfur
trioxide (SO3).
Sulfur dioxide contains 0.9978 g oxygen for every gram
of sulfur. Sulfur trioxide contains 1.497 g oxygen for
every gram of sulfur.
Determine the whole number ratio of g O : 1.00 g S in
sulfur dioxide to g O : 1.00 g S in sulfur trioxide.
11
SAMPLE PROBLEM
2.1
Solution
(b)
12
2.2
The Structure of the Atom
Topics
Discovery of the Electron
Radioactivity
The Proton and the Nucleus
Nuclear Model of the Atom
The Neutron
13
2.2
The Structure of the Atom
Discovery of the Electron
14
2.2
The Structure of the Atom
Discovery of the Electron
© The McGraw-Hill Companies,
Inc./Charles D. Winters, photographer
15
2.2
The Structure of the Atom
Discovery of the Electron
16
2.2
The Structure of the Atom
Discovery of the Electron
J. J. Thomson: Charge-to-mass ration of the electron
Millikan: Charge of the electron
17
2.2
The Structure of the Atom
Radioactivity
18
2.2
The Structure of the Atom
The Proton and the Nucleus
Rutherford’s Experiment
19
2.2
The Structure of the Atom
Nuclear Model of the Atom
The atom’s positive charges, Rutherford proposed, were all
concentrated in the nucleus, which is an extremely dense
central core within the atom.
The positively charged particles in the nucleus are called
protons.
20
2.2
The Structure of the Atom
The Neutron
When Chadwick bombarded a thin sheet of beryllium with 
particles, a very high energy radiation was emitted by the
metal that was not deflected by either electric or magnetic
fields.
Although similar to  rays, later experiments showed that the
rays actually consisted of a third type of subatomic particle,
which Chadwick named neutrons because they were
electrically neutral particles having a mass slightly greater
than that of protons.
21
2.2
The Structure of the Atom
The Neutron
22
2.2
The Structure of the Atom
The Neutron
23
2.3
Atomic Number, Mass Number, and Isotopes
Topics
Atomic Number, Mass Number, and Isotopes
24
2.3
Atomic Number, Mass Number, and Isotopes
Atomic Number, Mass Number, and Isotopes
The atomic number (Z) is the number of protons in the
nucleus of each atom of an element.
The mass number (A) is the total number of neutrons and
protons present in the nucleus of an atom of an element.
25
2.3
Atomic Number, Mass Number, and Isotopes
Atomic Number, Mass Number, and Isotopes
Most elements have two or more isotopes, atoms that have
the same atomic number (Z) but different mass numbers (A).
26
2.3
Atomic Number, Mass Number, and Isotopes
Atomic Number, Mass Number, and Isotopes
27
SAMPLE PROBLEM
2.2
Determine the numbers of protons, neutrons, and electrons
in each of the following species:
Setup
Number of protons = Z, number of neutrons = A – Z, and
number of electrons = number of protons.
Recall that the 14 in carbon-14 is the mass number.
28
SAMPLE PROBLEM
2.2
Solution
(a) The atomic number is 17, so there are 17 protons. The
mass number is 35, so the number of neutrons is 35 – 17
= 18.
The number of electrons equals the number of protons,
so there are 17 electrons.
29
SAMPLE PROBLEM
2.2
Solution
(b) Again, the atomic number is 17, so there are 17
protons.
The mass number is 37, so the number of neutrons is
37 – 17 = 20.
The number of electrons equals the number of protons,
so there are 17 electrons, too.
30
SAMPLE PROBLEM
2.2
Solution
(c)
The atomic number of K (potassium) is 19, so there are
19 protons. The mass number is 41, so there are 41 –
19 = 22 neutrons.
There are 19 electrons.
(d)
Carbon-14 can also be represented as 14C. The atomic
number of carbon is 6, so there are 6 protons and 6
electrons. There are 14 – 6 = 8 neutrons.
31
2.4
The Periodic Table
Topics
The Periodic Table
32
2.4
The Periodic Table
The Periodic Table
33
2.4
The Periodic Table
The Periodic Table
Horizontal rows called periods and in vertical columns called
groups or families.
Elements in the same group tend to have similar physical and
chemical properties.
The elements can be categorized as metals, nonmetals, or
metalloids.
34
2.4
The Periodic Table
The Periodic Table
A metal is a good conductor of heat and electricity, whereas a
nonmetal is usually a poor conductor of heat and electricity.
A metalloid has properties that are intermediate between
those of metals and nonmetals.
35
2.4
The Periodic Table
The Periodic Table
36
2.5
The Atomic Mass Scale and Average Atomic
Mass
Topics
The Atomic Mass Scale and Average Atomic Mass
37
2.5
The Atomic Mass Scale and Average Atomic
Mass
The Periodic Table
According to international agreement, atomic mass is the
mass of an atom in atomic mass units.
One atomic mass unit (amu) is defined as a mass exactly
equal to one-twelfth the mass of one carbon-12 atom.
38
2.5
The Atomic Mass Scale and Average Atomic
Mass
The Periodic Table
Experiments have shown that a hydrogen atom (1H) is only
8.3985 percent as massive as the carbon-12 atom.
Thus, if the mass of one carbon-12 atom is exactly 12 amu,
the atomic mass of hydrogen must be 0.083985 × 12 amu, or
1.0078 amu.
39
2.5
The Atomic Mass Scale and Average Atomic
Mass
The Periodic Table
When you look up the atomic mass of carbon in a table such
as the one on the inside front cover of the text, you will find
that its value is 12.01 amu, not 12.00 amu.
The difference arises because most naturally occurring
elements (including carbon) have more than one isotope.
The atomic masses in the periodic table are average atomic
masses. The term atomic weight is sometimes used to mean
average atomic mass.
40
2.5
The Atomic Mass Scale and Average Atomic
Mass
The Periodic Table
41
SAMPLE PROBLEM
2.3
Oxygen is the most abundant element in both Earth’s crust
and the human body. The atomic masses of its three stable
isotopes, 168O (99.757 percent), 178O (0.038 percent), and 188O
(0.205 percent), are 15.9949, 16.9991, and 17.9992 amu,
respectively.
Calculate the average atomic mass of oxygen using the
relative abundances given in parentheses. Report the result to
four significant figures.
42
SAMPLE PROBLEM
2.3
Solution
43
2.6
Molecules and Molecular Compounds
Topics
Molecules
Molecular Formulas
Naming Molecular Compounds
Empirical Formulas
44
2.6
Molecules and Molecular Compounds
Molecules
A molecule is a combination of at least two atoms in a specific
arrangement held together by electrostatic forces known as
covalent chemical bonds.
The hydrogen molecule, symbolized as H2, is called a diatomic
molecule because it contains two atoms.
Other elements that normally exist as diatomic molecules are
nitrogen (N2), oxygen (O2), and the Group 7A elements—
fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2).
45
2.6
Molecules and Molecular Compounds
Molecules
In homonuclear diatomic molecules, both atoms in each
molecule are of the same element.
A diatomic molecule can also contain atoms of different
elements (heteronuclear diatomic molecules).
46
2.6
Molecules and Molecular Compounds
Molecular Formulas
A chemical formula denotes the composition of the
substance.
A molecular formula shows the exact number of atoms of
each element in a molecule.
An allotrope is one of two or more distinct forms of an
element. Two of the allotropic forms of the element carbon—
diamond and graphite—have dramatically different
properties.
47
2.6
Molecules and Molecular Compounds
Molecular Formulas
The structural formula shows not only the
elemental composition, but also the general
arrangement of atoms within the molecule.
48
SAMPLE PROBLEM
2.4
Write the molecular formula of ethanol
based on its ball-and-stick model, shown
here.
49
SAMPLE PROBLEM
2.4
Write the molecular formula of ethanol
based on its ball-and-stick model, shown
here.
Solution
C2H6O
50
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
Binary molecular compounds consist of just two different
elements.
To name such a compound, we first name the element that
appears first in the formula.
We then name the second element, changing the ending of
its name to –ide.
HCl
SiC
hydrogen chloride
silicon carbide
51
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
Use Greek prefixes to denote the number of atoms of each
element present.
52
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
53
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
The prefix mono– is generally omitted for the first element.
SO2, for example, is named sulfur dioxide, not monosulfur
dioxide.
For ease of pronunciation, we usually eliminate the last letter
of a prefix that ends in “o” or “a” when naming an oxide.
Thus, N2O5 is dinitrogen pentoxide, rather than dinitrogen
pentaoxide.
54
SAMPLE PROBLEM
2.5
Name the following binary molecular compounds:
(a) NF3 and (b) N2O4.
55
SAMPLE PROBLEM
2.5
Name the following binary molecular compounds:
(a) NF3 and (b) N2O4.
Solution
(a) nitrogen trifluoride and (b) dinitrogen tetroxide
56
SAMPLE PROBLEM
2.6
Write the chemical formulas for the following binary
molecular compounds:
(a) sulfur tetrafluoride and (b) tetraphosphorus decasulfide.
57
SAMPLE PROBLEM
2.6
Write the chemical formulas for the following binary
molecular compounds:
(a) sulfur tetrafluoride and (b) tetraphosphorus decasulfide.
Solution
(a) SF4 and (b) P4S10
58
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
The names of molecular compounds containing hydrogen do
not usually conform to the systematic nomenclature
guidelines.
59
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
Acids make up another important class of molecular
compounds.
One definition of an acid is a substance that produces
hydrogen ions (H+) when dissolved in water.
Several binary molecular compounds produce hydrogen ions
when dissolved in water and are, therefore, acids.
In these cases, two different names can be assigned to the
same chemical formula.
60
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
For example, HCl, hydrogen chloride, is a gaseous compound.
When it is dissolved in water, however, we call it hydrochloric
acid.
The rules for naming simple acids of this type are as follows:
remove the –gen ending from hydrogen (leaving hydro–),
change the –ide ending on the second element to –ic,
combine the two words, and add the word acid.
61
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
62
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
Organic compounds contain carbon and hydrogen,
sometimes in combination with other elements such as
oxygen, nitrogen, sulfur, and the halogens.
The simplest organic compounds are those that contain only
carbon and hydrogen and are known as hydrocarbons.
Among hydrocarbons, the simplest examples are compounds
known as alkanes.
63
2.6
Molecules and Molecular Compounds
64
2.6
Molecules and Molecular Compounds
65
2.6
Molecules and Molecular Compounds
66
2.6
Molecules and Molecular Compounds
Naming Molecular Compounds
Many organic compounds are derivatives of alkanes in which
one of the H atoms has been replaced by a group of atoms
known as a functional group.
The functional group determines many of
the chemical properties of a compound
because it typically is where a chemical
reaction occurs.
67
2.6
Molecules and Molecular Compounds
68
2.6
Molecules and Molecular Compounds
Empirical Formulas
Molecular substances can also be represented using empirical
formulas.
The word empirical means “from experience” or, in the
context of chemical formulas, “from experiment.”
The empirical formula tells what elements are present in a
molecule and in what whole-number ratio they are combined.
69
2.6
Molecules and Molecular Compounds
70
SAMPLE PROBLEM
2.7
Write the empirical formulas for the following molecules:
(a) glucose (C6H12O6)
(b) adenine (C5H5N5)
(c) nitrous oxide (N2O)
71
SAMPLE PROBLEM
2.7
Write the empirical formulas for the following molecules:
(a) glucose (C6H12O6)
(b) adenine (C5H5N5)
(c) nitrous oxide (N2O)
Solution
(a) CH2O
(b) CHN
(c) N2O
72
2.7
Ions and Ionic Compounds
Topics
Atomic Ions
Polyatomic Ions
Formulas of Ionic Compounds
Naming Ionic Compounds
Naming Oxoanions and Oxoacids
Hydrates
Familiar Inorganic Compounds
73
2.7
Ions and Ionic Compounds
Atomic Ions
An atomic ion or monatomic ion is one that consists of just
one atom with a positive or negative charge.
The loss of one or more electrons from an atom yields a
cation, an ion with a net positive charge.
74
2.7
Ions and Ionic Compounds
Atomic Ions
An anion is an ion whose net charge is negative due to an
increase in the number of electrons.
Sodium chloride (NaCl), ordinary table salt, is called an ionic
compound because it consists of cations (Na+) and anions
(Cl–).
75
2.7
Ions and Ionic Compounds
Atomic Ions
An atom can lose or gain more than one electron.
76
2.7
Ions and Ionic Compounds
Atomic Ions
Designate different cations with Roman numerals, using the
Stock system.
77
2.7
Ions and Ionic Compounds
Atomic Ions
An older nomenclature system – for cations with two possible
charges - assigns the ending –ous to the cation with the
smaller positive charge and the ending –ic to the cation with
the greater positive charge:
78
2.7
Ions and Ionic Compounds
Atomic Ions
A monatomic anion is named by changing the ending of the
element’s name to –ide, and adding the word ion.
Thus, the anion of chlorine (Cl–), is called chloride ion.
79
2.7
Ions and Ionic Compounds
Polyatomic Ions
Ions that consist of a combination of two or more atoms are
called polyatomic ions.
80
2.7
Ions and Ionic Compounds
Polyatomic Ions
81
2.7
Ions and Ionic Compounds
Polyatomic Ions
82
2.7
Ions and Ionic Compounds
Formulas of Ionic Compounds
83
2.7
Ions and Ionic Compounds
Formulas of Ionic Compounds
Aluminum Oxide
Calcium Phosphate
84
2.7
Ions and Ionic Compounds
Naming Ionic Compounds
85
SAMPLE PROBLEM
2.8
Name the following ionic compounds:
(a) MgO
(b) Al(OH)3
(c) Fe2(SO4)3
Setup
(a) Mg2+
O2–
(b) Al3+
OH–
(c) Since the charge on SO42– is –2, the charge on Fe must be
+3: Fe3+
86
SAMPLE PROBLEM
2.8
Name the following ionic compounds:
(a) MgO
(b) Al(OH)3
(c) Fe2(SO4)3
Solution
(a) magnesium oxide
(b) aluminum hydroxide
(c) iron (III) sulfate
87
SAMPLE PROBLEM
2.9
Deduce the formulas of the following ionic compounds:
(a) mercury(I) chloride
(b) lead(II) chromate
(c) potassium hydrogen phosphate
Setup
(a) Hg22+
(b) Pb2+
(c) K+
Cl–
CrO42–
HPO42–
88
SAMPLE PROBLEM
2.9
Deduce the formulas of the following ionic compounds:
(a) mercury(I) chloride
(b) lead(II) chromate
(c) potassium hydrogen phosphate
Solution
(a) Hg2Cl2
(b) PbCrO4
(c) K2HPO4
89
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
Oxoanions are polyatomic anions that contain one or more
oxygen atoms and one atom (the “central atom”) of another
element.
Examples include the chlorate (ClO3– ), nitrate (NO3– ), and
sulfate (SO42–) ions.
90
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
Starting with the oxoanions whose names end in –ate, we can
name these ions as follows:
1. The ion with one more O atom than the –ate ion is called
the per . . . ate ion. Thus, ClO3– is the chlorate ion, so ClO4–
is the perchlorate ion.
2. The ion with one less O atom than the –ate anion is called
the –ite ion. Thus, ClO2– is the chlorite ion.
91
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
Starting with the oxoanions whose names end in –ate, we can
name these ions as follows:
3. The ion with two fewer O atoms than the –ate ion is called
the hypo . . . ite ion. Thus, ClO– is the hypochlorite ion.
92
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
An important class of acids known as oxoacids, which ionize
to produce hydrogen ions and the corresponding oxoanions.
The formula of an oxoacid can be determined by adding
enough H+ ions to the corresponding oxoanion to yield a
formula with no net charge.
For example, the formulas of oxoacids based on the nitrate
(NO3– ) and sulfate (SO42–) ions are HNO3 and H2SO4,
respectively.
93
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
The names of oxoacids are derived from the names of the
corresponding oxoanions using the following guidelines:
1. An acid based on an –ate ion is called . . . ic acid. Thus,
HClO3 is called chloric acid.
2. An acid based on an –ite ion is called . . . ous acid. Thus,
HClO2 is called chlorous acid.
94
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
The names of oxoacids are derived from the names of the
corresponding oxoanions using the following guidelines:
3. Prefixes in oxoanion names are retained in the names of
the corresponding oxoacids. Thus, HClO4 and HClO are
called perchloric acid and hypochlorous acid, respectively.
95
2.7
Ions and Ionic Compounds
Naming Oxoanions and Oxoacids
Many oxoacids, such as H2SO4 and H3PO4, are polyprotic—
meaning that they have more than one ionizable hydrogen
atom.
In these cases, the names of anions in which one or more (but
not all) of the hydrogen ions have been removed must
indicate the number of H ions that remain.
96
SAMPLE PROBLEM
2.10
Name the following species:
(a) BrO4–
(b) HCO3–
(c) H2CO3
Setup
(a) BrO3– is the bromate ion.
(b) CO3– is the carbonate ion.
97
SAMPLE PROBLEM
2.10
Name the following species:
(a) BrO4–
(b) HCO3–
(c) H2CO3
Solution
(a) perbromate ion
(b) hydrogen carbonate ion (or bicarbonate ion)
(c) carbonic acid
98
SAMPLE PROBLEM
2.11
Determine the formula of sulfurous acid.
Setup
The sulfite ion is SO32–.
99
SAMPLE PROBLEM
2.11
Determine the formula of sulfurous acid.
Solution
The formula of sulfurous acid is H2SO3.
100
2.7
Ions and Ionic Compounds
Hydrates
Hydrates are compounds that
have a specific number of water
molecules within their solid
structure.
© The McGraw-Hill Companies,
Inc./Charles D. Winters, photographer
101
2.7
Ions and Ionic Compounds
Familiar Inorganic Compounds
102