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Chapter Two:
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
2.1
Chapter 2 | Slide 2
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Three Important Laws
• Law of conservation of mass (Lavoisier):
– Mass is neither created nor destroyed
• Law of definite proportion (Proust):
– A given compound always contains exactly the
same proportion of elements by mass
2.2
Chapter 2 | Slide 3
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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
2.2
Chapter 2 | Slide 4
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Dalton’s Atomic Theory (1808)
• Each element is made up of tiny particles
called atoms.
2.3
Chapter 2 | Slide 5
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Dalton’s Atomic Theory (continued)
• The atoms of a given element are identical;
the atoms of different elements are different
in some fundamental way or ways.
2.3
Chapter 2 | Slide 6
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Dalton’s Atomic Theory (continued)
• Chemical compounds are formed when
atoms of different elements combine with
each other. A given compound always has
the same relative numbers and types of
atoms.
2.3
Chapter 2 | Slide 7
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Dalton’s Atomic Theory (continued)
• Chemical reactions involve reorganization
of the atoms—changes in the way they are
bound together.
• The atoms themselves are not changed in
a chemical reaction.
2.3
Chapter 2 | Slide 8
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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.
Chapter 2 | Slide 9
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Avogadro’s Hypothesis (1811)
• At the same temperature and pressure,
equal volumes of different gases contain
the same number of particles:
– 5 liters of oxygen
– 5 liters of nitrogen
– Same number of particles!
2.3
Chapter 2 | Slide 10
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Gay-Lussac and Avogadro
• 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
2.3
Chapter 2 | Slide 11
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Representing Gay-Lussac’s Results
2.3
Chapter 2 | Slide 12
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Representing Gay-Lussac’s Results
2.3
Chapter 2 | Slide 13
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Representing Gay-Lussac’s Results
2.3
Chapter 2 | Slide 14
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Representing Gay-Lussac’s Results
2.3
Chapter 2 | Slide 15
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Early Experiments to Characterize the Atom
• J. J. Thomson (1898-1903):
– Postulated the existence of 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 the electrons
2.4
Chapter 2 | Slide 16
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Cathode Ray Tube
2.4
Chapter 2 | Slide 17
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Early Experiments to Characterize the Atom
• Robert Millikan (1909):
– Performed experiments involving charged oil
drops
– Determined the magnitude of the electron
charge
– Calculated the mass of the electron
2.4
Chapter 2 | Slide 18
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Millikan Oil Drop Experiment
2.4
Chapter 2 | Slide 19
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Early Experiments to Characterize the Atom
• Ernest Rutherford (1911):
– Explained the nuclear atom
– Atom has a dense center of positive charge
called the nucleus
– Electrons travel around the nucleus at a
relatively large distance
2.4
Chapter 2 | Slide 20
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Rutherford’s Gold Foil Experiment
2.4
Chapter 2 | Slide 21
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The Modern View of Atomic Structure
• The atom contains:
 Electrons
 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
2.5
Chapter 2 | Slide 22
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The Modern View of Atomic Structure
• The nucleus is:
– Small compared with the overall size of the
atom
– Extremely dense; accounts for almost all of the
atom’s mass
2.5
Chapter 2 | Slide 23
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Nuclear Atom
Viewed in
Cross Section
Chapter 2 | Slide 24
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The Modern View of Atomic Structure
• 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
2.5
Chapter 2 | Slide 25
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Two Isotopes of Sodium
2.5
Chapter 2 | Slide 26
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Exercise
A certain isotope X+ contains 54 electrons and 78
neutrons. What is the mass number of this
isotope?
2.5/2.6
Chapter 2 | Slide 27
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Chemical Bonds
• Covalent Bonds:
– Bonds form between atoms by sharing
electrons
– Resulting collection of atoms is called a
molecule
2.6
Chapter 2 | Slide 28
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Covalent Bonding
2.6
Chapter 2 | Slide 29
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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)
2.6
Chapter 2 | Slide 30
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Molecular vs. Ionic Compounds
2.6
Chapter 2 | Slide 31
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The Periodic Table
• Metals vs. Nonmetals
• Groups or Families – elements in the same
vertical columns
– Alkali metals, alkaline earth metals, halogens,
and noble gases
• Periods – horizontal rows of elements
2.7
Chapter 2 | Slide 32
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The Periodic Table
2.7
Chapter 2 | Slide 33
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Naming Compounds
• Binary Compounds:
– Composed of two elements
– Ionic and covalent compounds included
• Binary Ionic Compounds:
– Metal-nonmetal
• Binary Covalent Compounds:
– Nonmetal to nonmetal
2.8
Chapter 2 | Slide 34
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Binary Ionic Compounds (Type I)
• Cation is always named first and the anion
second.
• Monatomic cation has the same name as
its parent element.
• Monatomic anion is named by taking the
root of the element name and adding –ide.
2.8
Chapter 2 | Slide 35
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Binary Ionic Compounds (Type I)
• Examples:
KCl
Potassium chloride
MgBr2
Magnesium bromide
CaO
Calcium oxide
2.8
Chapter 2 | Slide 36
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Binary Ionic Compounds (Type II)
• Metals in these compounds form more than
one type of positive charge.
• 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.
2.8
Chapter 2 | Slide 37
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Binary Ionic Compounds (Type II)
• Examples:
CuBr
Copper(I) bromide
FeS
Iron(II) sulfide
PbO2
Lead(IV) oxide
2.8
Chapter 2 | Slide 38
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Polyatomic Ions
• Must be memorized (see Table 2.5 on pg.
62 in text).
• Examples of compounds containing
polyatomic ions:
NaOH
Mg(NO3)2
(NH4)2SO4
Sodium hydroxide
Magnesium nitrate
Ammonium sulfate
2.8
Chapter 2 | Slide 39
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Formation of Ionic Compounds
2.8
Chapter 2 | Slide 40
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Binary Covalent Compounds (Type III)
• Formed between two nonmetals.
• First element is named first, using the full
element name.
• Second element is named as if it were an
anion.
• Prefixes are used to denote the numbers of
atoms present.
• Prefix mono- is never used for naming the
first element.
2.8
Chapter 2 | Slide 41
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Binary Covalent Compounds (Type III)
• Examples:
CO2
Carbon dioxide
SF6
Sulfur hexafluoride
N 2O 4
Dinitrogen tetroxide
2.8
Chapter 2 | Slide 42
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Overall Strategy for Naming Chemical Compounds
Chapter 2 | Slide 43
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Flowchart for Naming Binary Compounds
Chapter 2 | Slide 44
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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.
2.8
Chapter 2 | Slide 45
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Acids
• If the anion does not contain oxygen, the
acid is named with the prefix hydro- and the
suffix -ic.
• Examples:
HCl
HCN
H2S
Hydrochloric acid
Hydrocyanic acid
Hydrosulfuric acid
2.8
Chapter 2 | Slide 46
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Acids
• If the anion does contain oxygen:
– The suffix -ic is added to the root name if the
anion name ends in -ate.
• Examples:
HNO3
H2SO4
HC2H3O2
Nitric acid
Sulfuric acid
Acetic acid
2.8
Chapter 2 | Slide 47
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Acids
• If the anion does contain oxygen:
– The suffix -ous is added to the root name if the
anion name ends in -ite.
• Examples:
HNO2
H2SO3
HClO2
Nitrous acid
Sulfurous acid
Chlorous acid
2.8
Chapter 2 | Slide 48
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Flowchart for Naming Acids
Chapter 2 | Slide 49
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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) peroxide
tin(IV) hydroxide
phosphorus pentabromide
calcium chromate
2.8
Chapter 2 | Slide 50
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