Atoms, Molecules and Ions

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Atoms, Molecules and Ions

• Chapter 2

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Chapter 2 The Components of Matter
2.1 Elements, Compounds, and Mixtures An Atomic
Overview 2.2 The Observations That Led to an
Atomic View of Matter 2.3 Daltons Atomic
Theory 2.4 The Observations That Led to the
Nuclear AtomModel 2.5 The Atomic Theory Today 2.6
Elements A First Look at the Periodic Table 2.7
Compounds Formulas, Names, and Masses
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Daltons Atomic Theory (1808)

• Elements are composed of extremely small
  particles called atoms. All atoms of a given
  element are identical. The atoms of one element
  are different from the atoms of all other
  elements.

• Compounds are composed of atoms of more than one
  element. The relative number of atoms of each
  element in a given compound is always the same.

• Chemical reactions only involve the
  rearrangement of atoms. Atoms are not created or
  destroyed in chemical reactions.

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Law of Definite (or Constant) Composition No
matter what its source, a particular chemical
compound is composed of the same elements in the
same parts (fractions) by mass.
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Definite Composition
Consider a 20. 0 g of calcium carbonate
Analysis by mass mass fraction by mass
(g/20.0g) (parts/1.00 part) (parts/100 parts)
40 12 48
8.0 g 2.4 g 9.6 g
0.40 0.12 0.48
Calcium Carbon Oxygen
20.0 g
1.00 part by mass
100 by mass
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Law of Multiple Proportions
If elements A and B react to form two
compounds, the different masses of B that combine
with a fixed mass of A can be expressed as a
ratio of small whole numbers
Example Nitrogen Oxides I II
Nitrogen Oxide I 46.68 Nitrogen and 53.32
Oxygen Nitrogen Oxide II 30.45 Nitrogen and
69.55 Oxygen
Assume that you have 100g of each compound. In
100 g of each compound g O 53.32g for oxide I
69.55g for oxide II
g N 46.68g for oxide I
30.45g for oxide II
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Law of Conservation of Mass The total mass of
substances does not change during a chemical
reaction.
reactant 1 reactant 2
product
56.08g 44.00g
100.08g
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Daltons Atomic Theory
1. All matter consists of atoms.
2. Atoms of one element cannot be converted into
atoms of another element.
3. Atoms of an element are identical in mass and
otherproperties and are different from atoms of
any other element.
4. Compounds result from the chemical combination
of a specific ratio of atoms of different
elements.
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The Atomic Basis of the Law of Multiple
Proportions
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8 X2Y
Law of Conservation of Mass
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The Modern Reassessment of the Atomic Theory
1. All matter is composed of atoms. The atom is
the smallest body that retains the unique
identity of the element. 2. Atoms of one element
cannot be converted into atoms of another element
in a chemical reaction. Elements can only be
converted into other elements in nuclear
reactions. 3. All atoms of an element have the
same number of protons and electrons, which
determines the chemical behavior of the element.
Isotopes of an element differ in the number of
neutrons, and thus in mass number. A sample of
the element is treated as though its atoms have
an average mass. 4. Compounds are formed by the
chemical combination of two or more elements in
specific ratios.
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J.J. Thomson, measured mass/charge of e- (1906
Nobel Prize in Physics)
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Measured mass of e- (1923 Nobel Prize in Physics)
e-
charge -1.60 x 10-19 C Thomsons charge/mass of
e- -1.76 x 108 C/g
e- mass 9.10 x 10-28 g
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(Uranium compound)
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(1908 Nobel Prize in Chemistry)

• particle velocity 1.4 x 107 m/s
• (5 speed of light)

• atoms positive charge is concentrated in the
  nucleus
• proton (p) has opposite () charge of electron
• mass of p is 1840 x mass of e- (1.67 x 10-24 g)

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Rutherfords 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
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Chadwicks Experiment (1932)
H atoms - 1 p He atoms - 2 p mass He/mass H
should 2 measured mass He/mass H 4
neutron (n) is neutral (charge 0) n mass p
mass 1.67 x 10-24 g
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Subatomic Particles (Table 2.1)
mass p mass n 1840 x mass e-
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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 the nucleus
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2.3
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Do You Understand Isotopes?
6 protons, 8 (14 - 6) neutrons, 6 electrons
6 protons, 5 (11 - 6) neutrons, 6 electrons
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2.4
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A molecule is an aggregate of two or more atoms
in a definite arrangement held together by
chemical bonds
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
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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.
anion ion with a negative charge If a neutral
atom gains one or more electrons it becomes an
anion.
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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-
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Do You Understand Ions?
13 protons, 10 (13 3) electrons
34 protons, 36 (34 2) electrons
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Common Monoatomic Ions
Table 2.3
1
-1
2
-2
Common ions are in blue.
3
-3
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Metals With Several Oxidation States
Table 2.4 (partial)
Element
Ion Formula
Systematic Name
Common Name
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2.6
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A molecular formula shows the exact number of
atoms of each element in the smallest unit of a
substance
An empirical formula shows the simplest
whole-number ratio of the atoms in a substance
H2O
CH2O
C6H12O6
O3
O
N2H4
NH2
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• ionic compounds consist of a cation and an anion
• the formula is always the same as the empirical
  formula
• the sum of the charges on the cation and anion
  in each formula unit must equal zero

The ionic compound NaCl
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Formula of Ionic Compounds
Al2O3
Al3
O2-
CaBr2
Ca2
Br-
Na2CO3
Na
CO32-
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Some Polyatomic Ions (Table 2.3)
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Some Common Polyatomic Ions
Formula
Formula
Name
Name
Cations
H3O
hydronium
ammonium
NH4
Common Anions
acetate
CH3COO-
CN-
cyanide
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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
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• Transition metal ionic compounds
• indicate charge on metal with Roman numerals

iron(II) chloride
FeCl2
2 Cl- -2 so Fe is 2
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
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Sample Problem 2.12
Recognizing Incorrect Names and Fromulas of Ionic
Compounds
(a) Ba(C2H3O2)2 is called barium diacetate.
(b) Sodium sulfide has the formula (Na)2SO3.
(c) Iron(II) sulfate has the formula Fe2(SO4)3.
(d) Cesium carbonate has the formula Cs2(CO3).
SOLUTION
(a) Barium is always a 2 ion and acetate is -1.
The di- is unnecessary.
(b) An ion of a single element does not need
parentheses. Sulfide is S2-, not SO32-. The
correct formula is Na2S.
(c) Since sulfate has a 2- charge, only 1 Fe2
is needed. The formula should be FeSO4.
(d) The parentheses are unnecessary. The
correct formula is Cs2CO3.
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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

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Molecular Compounds
HI
hydrogen iodide
NF3
nitrogen trifluoride
SO2
sulfur dioxide
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
N2O
dinitrogen monoxide
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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.
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Naming oxoanions
Figure 2.20
Prefixes
Root
Suffixes
Examples
root
per
ate
ClO4-
perchlorate
ate
root
ClO3-
chlorate
No. of O atoms
ite
root
ClO2-
chlorite
ite
hypo
root
ClO-
hypochlorite
Numerical Prefixes for Hydrates and Binary
Covalent Compounds
Table 2.6
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A base can be defined as a substance that yields
hydroxide ions (OH-) when dissolved in water.
2.7