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Chapter 2
Atoms, Molecules,
and Ions
Early History of Chemistry
•
Greeks were the first to attempt to explain why
chemical changes occur.
• Alchemy dominated for 2000 years.
– Several elements discovered.
– Mineral acids prepared.
• Robert Boyle was the first “chemist”.
– Performed quantitative experiments.
– Developed first experimental definition of an
element.
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Three Important Laws
•
Law of conservation of mass (Lavoisier):
– Mass is neither created nor destroyed in a
chemical reaction.
•
Law of definite proportion (Proust):
– A given compound always contains exactly
the same proportion of elements by mass.
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Three Important Laws
(continued)
•
Law of multiple proportions (Dalton):
– When two elements form a series of
compounds, the ratios of the masses of the
second element that combine with 1 gram of
the first element can always be reduced to
small whole numbers.
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Composition of the Atom
• Understanding the structure of the atom
will help to understand the properties of
the elements.
• Keep in mind that these, as all theories,
are subject to constant refinement. The
picture of the atom isn’t final.
Development of the Atomic Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom.
– John Dalton
– early 1800’s
• Much of Dalton’s Theory is still regarded as
correct today. *See starred items.*
Postulates of Dalton’s Atomic Theory
1. All matter consists of tiny particles
called atoms.*
2. An atom cannot be created, divided,
destroyed, or converted to any other
type of atom.
3. Atoms of a particular element have
identical properties.
4. Atoms of different elements have different
properties.*
5. Atoms of different elements combine in
simple whole-number ratios to produce
compounds (stable aggregates of
atoms.)*
6. Chemical change involves joining,
separating, or rearranging atoms.*
* These postulates are still regarded as
true.
CONCEPT CHECK!
Which of the following statements regarding Dalton’s
atomic theory are still believed to be true?
I. Elements are made of tiny particles called atoms.
II. All atoms of a given element are identical.
III. A given compound always has the same relative
numbers and types of atoms.
IV. Atoms are indestructible.
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Dalton’s Atomic Theory
2
16 X
+
8Y
8 X2Y
Gay-Lussac and Avogadro
(1809—1811)
•
Gay—Lussac
– Measured (under same conditions of T and P)
the volumes of gases that reacted with each
other.
• Avogadro’s Hypothesis
– At the same T and P, equal volumes of
different gases contain the same number of
particles.
• Volume of a gas is determined by the number,
not the size, of molecules.
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Representing Gay—Lussac’s
Results
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Representing Gay—Lussac’s
Results
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Representing Gay—Lussac’s
Results
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Representing Gay—Lussac’s
Results
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We will learn that atoms consist of three
primary particles.
electrons
protons
neutrons
• Nucleus - small, dense, positively charged region
in the center of the atom. Contains:
- protons - positively charged particles
- neutrons - uncharged particles
• Surrounding the nucleus is a diffuse region of
negative charge populated by:
- electrons - negatively charged particles
J. J. Thomson (1898—1903)
•
•
•
Postulated the existence of negatively charged
particles, that we now call electrons, using
cathode-ray tubes.
Determined the charge-to-mass ratio of an
electron.
The atom must also contain positive particles
that balance exactly the negative charge carried
by electrons.
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18
Subatomic Particles: Thompson Experiment
Electrons, Protons and Neutrons
• Electrons were the first subatomic
particles to be discovered using the
cathode ray tube.
Indicated that the
particles were
negatively charged.
Cathode Ray Tube
Joseph John Thomson (1856-1940). Photo courtesy of The Cavendish Laboratory.
Robert Millikan (1909)
•
Performed experiments involving charged oil
drops.
• Determined the magnitude of the charge on a
single electron.
• Calculated the mass of the electron
– (9.11 × 10-31 kg).
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Measured mass of e(1923 Nobel Prize in Physics)
e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g
Millikan Oil Drop Experiment
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Henri Becquerel (1896)
• Discovered radioactivity by observing the
spontaneous emission of radiation by
uranium.
• Three types of radioactive emission exist:
– Gamma rays (ϒ) – high energy light
– Beta particles (β) – a high speed electron
– Alpha particles (α) – a particle with a 2+ charge
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(Uranium compound)
Ernest Rutherford (1911)
•
•
•
Explained the nuclear atom.
The atom has a dense center of positive charge
called the nucleus.
Electrons travel around the nucleus at a large
distance relative to the nucleus.
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(1908 Nobel Prize in Chemistry)
 particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1. atoms positive charge is concentrated in the nucleus
2. proton (p) has opposite (+) charge of electron (-)
3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
Rutherford’s Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
“If the atom is the Houston Astrodome, then
the nucleus is a marble on the 50-yard line.”
•
The nucleus is:
– Small compared with the overall size of the
atom.
– Extremely dense; accounts for almost all of
the atom’s mass.
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Nuclear Atom Viewed in Cross
Section
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Discovery of the Proton
• The proton was discovered in 1918 by Ernest
Rutherford. He noticed that when alpha particles were
shot into nitrogen gas, his scintillation detectors
showed the signatures of hydrogen nuclei. Rutherford
determined that the only place this hydrogen could
have come from was the nitrogen, and therefore
nitrogen must contain hydrogen nuclei. He thus
suggested that the hydrogen nucleus, which was
known to have an atomic number of 1, was an
elementary particle. This he named proton, from
protos, the Greek for "first".
Chadwick’s Experiment (1932)
(1935 Noble Prize in Physics)
H atoms - 1 p; He atoms - 2 p
mass He/mass H should = 2
measured mass He/mass H = 4
 + 9Be
1n
+ 12C + energy
neutron (n) is neutral (charge = 0)
n mass ~ p mass = 1.67 x 10-24 g
Atom constitution summary
•
The atom contains:
– Electrons – found outside the nucleus;
negatively charged.
– Protons – found in the nucleus; positive
charge equal in magnitude to the electron’s
negative charge.
– Neutrons – found in the nucleus; no charge;
virtually same mass as a proton.
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Selected Properties of the Subatomic Particles
Name
Charge
Mass(amu) Mass (grams)
Electrons (e) -1
5.4 x 10-4
9.1095 x 10-28
Protons (p)
+1
1.00
1.6725 X 10-24
Neutrons (n) 0
1.00
1.6750 x 10-24
Atomic number, Mass number and Isotopes
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different
numbers of neutrons in their nuclei
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
U
(T)
Charge of
particle
Mass
Number
A
Z
Atomic
Number
X
C
Symbol of
the atom
atomic number (Z) - the number of
protons in the atom
mass number (A) - sum of the number
of protons and neutrons
Calculate the number of protons,
neutrons and electrons in each of the
following:
11
5
B
55
26
Fe
Isotopes
•
•
•
Atoms with the same number of protons but
different numbers of neutrons.
Show almost identical chemical properties;
chemistry of atom is due to its electrons.
In nature most elements contain mixtures of
isotopes.
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Two Isotopes of Sodium
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• Isotopes - atoms of the same element
having different masses.
– contain same number of protons
– contain different numbers of neutrons
Isotopes of Hydrogen
Hydrogen
(Hydrogen
1)
Tritium
Deuterium
(Hydrogen - 2) (Hydrogen - 3)
Do You Understand Isotopes?
How many protons, neutrons, and electrons are
14
in 6 C ?
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are
11
in 6 C ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
EXERCISE!
A certain isotope X contains 23 protons and 28
neutrons.
• What is the mass number of this isotope?
• Identify the element.
Mass Number = 51
Vanadium
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A molecule is an aggregate of two or more atoms in a
definite arrangement held together by chemical forces
H2
H2O
NH3
CH4
A diatomic molecule contains only two atoms
H2, N2, O2, Br2, HCl, CO
A polyatomic molecule contains more than two atoms
O3, H2O, NH3, CH4
Chemical Bonds
•
Covalent Bonds
– Bonds form between atoms by sharing
electrons.
– Resulting collection of atoms is called a
molecule.
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Covalent Bonding
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Chemical Bonds
•
Ionic Bonds
– Bonds form due to force of attraction
between oppositely charged ions.
– Ion – atom or group of atoms that has a net
positive or negative charge.
– Cation – positive ion; lost electron(s).
– Anion – negative ion; gained electron(s).
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Chemical Formulas; Molecular
and Ionic Substances
• Ionic substances
Although many substances are molecular,
others are composed of ions.
An ion is an electrically charged particle
obtained from an atom or chemically bonded
group of atoms by adding or removing
electrons.
Sodium chloride is a substance made up of
ions.
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
How many protons, neutrons and
electrons are in the following ions?
39
19
32
16
24
12
K
S

2-
Mg
2
A monatomic ion contains only one atom
Na+, Cl-, Ca2+, O2-, Al3+, N3-
A polyatomic ion contains more than one atom
OH-, CN-, NH4+, NO3-
Do You Understand Ions?
How many protons and electrons are in
27 3+
13 Al ?
13 protons, 10 (13 – 3) electrons
How many protons and electrons are in
78
2Se
?
34
34 protons, 36 (34 + 2) electrons
Molecular vs Ionic Compounds
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The Periodic Table
• In 1869, Dmitri Mendeleev discovered that if the
known elements were arranged in order of atomic
number, they could be placed in horizontal rows such
that the elements in the vertical columns had similar
properties.
A tabular arrangement of elements in rows
and columns, highlighting the regular
repetition of properties of the elements, is
called a periodic table.
The Periodic Table
• Periods and Groups
A period consists of the elements in one
horizontal role of the periodic table.
A group consists of the elements in any
one column of the periodic table.
The groups are usually numbered.
The eight groups are called main group
(or representative) elements.
The Periodic Table
• Periods and Groups
The “B” groups are called transition
elements.
The two rows of elements at the bottom of
the table are called inner transition
elements.
Elements in any one group have similar
properties.
The Periodic Table
• Periods and Groups
The elements in group IA, often known as
the alkali metals, are soft metals that react
easily with water.
The group VIIA elements, known as the
halogens, are also reactive elements.
The Periodic Table
• Metals, Nonmetals, and Metalloids
A metal is a substance or mixture that has a
characteristic luster and is generally a good
conductor of heat and electricity.
A nonmetal is an element that does not
exhibit the characteristics of the metal.
A metalloid, or semi-metal, is an element
having both metallic and nonmetallic
properties.
Noble Gas
Halogen
Group
Alkali Metal
Alkali Earth Metal
Period
Chemistry In Action
Natural abundance of elements in Earth’s crust
Natural abundance of elements in human body
Chemical Substances;
Formulas and Names
A monatomic ion is an ion formed from a
single atom.
Chemical Substances;
Formulas and Names
• Rules for predicting charges on
monatomic ions
Most of the main group metals form cations
with the charge equal to their group number.
The charge on a monatomic anion for a
nonmetal equals the group number minus 8.
Most transition elements form more than one
ion, each with a different charge.
Chemical Substances; Formulas
and Names
• Rules for naming monatomic ions
Monatomic cations are named after the element – if
there is only one such ion. For example, Al3+ is
called the aluminum ion.
If there is more than one cation of an element, a
Roman numeral in parentheses denoting the
charge on the ion is used. This often occurs with
transition elements. i.e. Fe2+ = Iron(II); Fe3+ =
Iron(III)
The names of the monatomic anions use the stem
name of the element followed by the suffix – ide.
For example, Br- is called the bromide ion.
Naming Compounds
•
Binary Compounds
– Composed of two elements
– Ionic and covalent compounds included
• Binary Ionic Compounds
– Metal—nonmetal
• Binary Covalent Compounds
– Nonmetal—nonmetal
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Binary Ionic Compounds (Type I)
1. The cation is always named first and the anion
second.
2. A monatomic cation takes its name from the
name of the parent element.
3. A monatomic anion is named by taking the root
of the element name and adding –ide.
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Binary Ionic Compounds (Type I)
•
Examples:
KCl
Potassium chloride
MgBr2
Magnesium bromide
CaO
Calcium oxide
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Naming Binary Compounds
• NaF
• LiCl
• MgO
-
Naming Binary Compounds
• NaF
• LiCl
• MgO
-
Sodium Fluoride
Lithium Chloride
Magnesium Oxide
Formula of Ionic Compounds
2 x +3 = +6
3 x -2 = -6
Al2O3
Al3+
1 x +2 = +2
Ca2+
1 x +2 = +2
Na+
O22 x -1 = -2
CaBr2
Br1 x -2 = -2
Na2CO3
CO32-
Formula of Ionic Compounds
1 x +2 = +2
1 x -2 = -2
MgO
Mg2+
O2-
NOT
1 x +2 = +2
Mg2+
1 x -2 = -2
Mg2O2
O2-
Use lowest common denominator
Binary Ionic Compounds (Type II)
•
•
•
•
•
Metals in these compounds form more than one
type of positive ion.
Charge on the metal ion must be specified.
Roman numeral indicates the charge of the
metal cation.
Transition metal cations usually require a Roman
numeral.
Elements that form only one cation do not need
to be identified by a roman numeral.
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Binary Ionic Compounds (Type II)
•
Examples:
CuBr
Copper(I) bromide
FeS
Iron(II) sulfide
PbO2
Lead(IV) oxide
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• Transition metal ionic compounds
– indicate charge on metal with Roman numerals
FeCl2
2 Cl- -2 so Fe is +2
iron(II) chloride
FeCl3
3 Cl- -3 so Fe is +3
iron(III) chloride
Cr2S3
3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide
Polyatomic Ions
•
•
Must be memorized (see Table 2.5 on pg. 65 in
text).
Examples of compounds containing polyatomic
ions:
NaOH
Sodium hydroxide
Mg(NO3)2
Magnesium nitrate
(NH4)2SO4
Ammonium sulfate
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Minimum
Polyatomic Ions You Should Know
•
•
•
•
•
NH4+ - Ammonium
OH- - Hydroxide
CN- - Cyanide
SO42- - Sulfate
ClO4- - Perchlorate
•
•
•
•
O22- - Peroxide
PO43- - Phosphate
CO32- - Carbonate
HCO3- - Hydrogen
carbonate
Chemical Substances;
Formulas and Names
• Polyatomic ions
A polyatomic ion is an ion consisting of two or
more atoms chemically bonded together and
carrying a net electric charge.

NO 3 nitrate

NO 2 nitrite
2
SO 4 sulfate
2
SO 3 sulfite
Prefix thio• The prefix thio- means that an oxygen atom
in the root ion name has been replaced by a
sulfur atom.
Example:
• SO42- --> S2O32sulfate ion
thiosulfate ion
Chemical Nomenclature
• Ionic Compounds
– often a metal + nonmetal
– anion (nonmetal), add “ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
More Practice
• Na2SO4
Sodium Sulfate
• AgCN
Silver Cyanide
• Ca(OCl)2
Calcium Hypochlorite
Na2SO3
Sodium Sulfite
Cd(OH)2
Cadmium Hydroxide
KClO4
Potassium Perchlorate
Formation of Ionic Compounds
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• Molecular compounds
• nonmetals or nonmetals + metalloids
• common names
• H2O, NH3, CH4, C60
• element further left in periodic table
is 1st
• element closest to bottom of group is
1st
• if more than one compound can be
formed from the same elements, use
prefixes to indicate number of each
kind of atom
• last element ends in ide
Binary Covalent Compounds
(Type III)
• Formed between two nonmetals.
1. The first element in the formula is named first,
using the full element name.
2. The second element is named as if it were an
anion.
3. Prefixes are used to denote the numbers of
atoms present.
4. The prefix mono- is never used for naming the
first element.
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Prefixes Used to
Indicate
Number in
Chemical Names
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Binary Covalent Compounds
(Type III)
•
Examples:
CO2
Carbon dioxide
SF6
Sulfur hexafluoride
N2O4
Dinitrogen tetroxide
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Molecular Compounds
HI
hydrogen iodide
NF3
nitrogen trifluoride
SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N2O
dinitrogen monoxide
TOXIC!
Laughing Gas
Acids
•
•
Acids can be recognized by the hydrogen that
appears first in the formula—HCl.
Molecule with one or more H+ ions attached to
an anion.
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An acid can be defined as a substance that yields
hydrogen ions (H+) when dissolved in water.
HCl
•Pure substance, hydrogen chloride
•Dissolved in water (H+ Cl-), hydrochloric acid
An oxoacid is an acid that contains hydrogen,
oxygen, and another element.
HNO3
nitric acid
H2CO3
carbonic acid
H2SO4
sulfuric acid
HNO3
Acids
•
•
If the anion does not contain oxygen, the acid is
named with the prefix hydro– and the suffix –ic.
Examples:
HCl
Hydrochloric acid
HCN
Hydrocyanic acid
H2S
Hydrosulfuric acid
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Oxoacid
•
If the anion does contain oxygen:
– The suffix –ic is added to the root name if the
anion name ends in –ate.
• Examples:
HNO3
Nitric acid
H2SO4
Sulfuric acid
HC2H3O2
Acetic acid
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Oxoacid
•
If the anion does contain oxygen:
– The suffix –ous is added to the root name if
the anion name ends in –ite.
• Examples:
HNO2 Nitrous acid
H2SO3 Sulfurous acid
HClO2 Chlorous acid
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Acid anions (provide
+
H
ions)
• Uses of mono– and di• Rule: to be used if oxyanion is bonded to
one or more hydrogen ions.
Example:
• HPO42- = monohydrogen phosphate ion
• H2PO4- = dihydrogen phosphate ion
Flowchart for Naming Acids
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A base can be defined as a substance that yields
hydroxide ions (OH-) when dissolved in water.
NaOH
sodium hydroxide
KOH
potassium hydroxide
Ba(OH)2
barium hydroxide
EXERCISE!
Which of the following compounds is named
incorrectly?
a) KNO3
b) TiO2
c) Sn(OH)4
d) PBr5
e) CaCrO4
potassium nitrate
titanium(II) oxide
tin(IV) hydroxide
phosphorus pentabromide
calcium chromate
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Worked Example 2.1