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Atoms, Molecules, and Ions
Only in the last fifteen years has it become
possible to “see” individual atoms.
Carbon
atoms in
graphite:
Dalton’s Atomic Theory
The modern view of matter did not emerge until
1806 with John Dalton’s atomic theory:
• Each element is composed of atoms.
• Atoms of a given element are all the same.
Atoms of different elements are different.
• Atoms are not changed into different atoms in
a chemical reaction.
• Compounds are formed when atoms of two or
more elements combine.
Law of Multiple Proportions
In compounds of elements A and B, the ratio of
masses of A and B are small whole numbers.
8 g oxygen
Water: 1 g hydrogen
16
g
oxygen
Hydrogen peroxide:
1 g hydrogen
Hydrogen peroxide must contain twice as much
oxygen as water.
Another example:
The Structure of Atoms
In the mid-1800s, negatively-charged particles
called electrons were discovered using
cathode-ray tubes.
The Mass of an Electron
• The mass/charge ratio was measured by J.J.
Thomson using a cathode ray tube.
• The charge of an electron was measured in a
famous experiment by Robert Millikan.
• The mass of an electron was found to be
about 2000 times less than the lightest atom
(hydrogen).
Radioactivity
Other particles were discovered being emitted
from uranium.
 rays: Helium nuclei
 rays: High speed electrons
 rays: High energy photons
(electromagnetic radiation)
The Nuclear Model
Rutherford’s gold foil experiment proved that
most of an atom’s mass is found in a very small
volume called the nucleus.
150 pm
Nucleus
0.01 pm
Electrons
Protons, Neutrons, and
Electrons
Two more subatomic particles were discovered
in the early 1900s - protons and neutrons:
Particle
proton
neutron
electron
Units:
Charge
+1
0
-1
Mass (amu)
1.0073
1.0087
0.00055
1 charge = 1.60210-19 Coulombs
1 atomic mass unit = 1.66010-24 g
o
1 A = 10-10 m
Atomic Symbols
• The atomic number of an element is the
number of protons in an atom.
• Isotopes have the same number of protons
but different numbers of neutrons.
• The mass number of an isotope is the sum of
protons and neutrons.
Isotopic symbols for two isotopes of carbon:
Carbon-12
12
6
C
Carbon-14
14
6
C
 C
 C
12
14
The Periodic Table
A chart of the elements in order of increasing
atomic number arranged so that elements with
similar chemical properties are in columns, or
groups.
Atomic Elements
• All metals and many semimetals consist of
atoms grouped together.
• The only elements that consist of isolated
atoms are the rare gases He, Ne, Ar, Kr,
Xe, and Rn.
• The chemical formula for an atomic element
is just the atomic symbol.
He
Molecules
• A molecule is an assembly of two or more
tightly bound atoms. Elements or compounds
made up of molecules are called molecular.
• The molecular formula indicates the number
and type of each atom in a molecule.
• Molecular elements: H2, N2, O2, F2, Cl2, Br2,
and I2.
N
N
Molecular Compounds
Molecular compounds are composed of two or
more nonmetals.
Compound
Water
Ammonia
Methane
Methanol
Carbon dioxide
Dinitrogen oxide
Molecular Formula
H 2O
NH3
CH4
CH3OH
CO2
N 2O
Empirical Formulas
The empirical formula gives only the relative
numbers of each element in a compound.
Analysis of a compound usually yields the
empirical formula.
Molecular Formula
C 2H 2
C 3H 6
C4H10
SO2
Empirical Formula
CH
CH2
C 2H 5
SO2
Ions
Atoms often gain or lose electrons to form
charged ions. A postive ion is a cation, a
negative ion is an anion.
Na atom - 1 electron  Na+ cation
11 protons
11 electrons
11 protons
10 electrons
Cl atom + 1 electron  Cl anion
17 protons
17 electrons
17 protons
18 electrons
Isotopic Symbols for Ions
1H +
Protons
1
Neutrons
0
Electrons
0
2H
1
1
1
19F
9
10
10
63Cu+2
29
34
27
140Ce+4
58
82
54
127I
53
74
54
235U
92
143
92
Ionic Compounds
A compound that consists of ions is an ionic
compound. Its chemical formula is an empirical
formula.
Microscopic
view of NaCl:
Ionic Compounds
• Metal atoms tend to lose electrons, nonmetals tend to gain electrons.
• As a general rule, compounds with a metal
and a nonmetal in the formula are ionic
compounds.
• In ionic formulas, the charges sum to zero:
Ca+2 and Cl form CaCl2
Mg+2 and O2 form MgO
Fe+3 and S2 form Fe2S3
Predicting Charges
Group
1 (1A)
2 (2A)
13 (3A)
15 (5A)
16 (6A)
17 (7A)
18 (8A)
Charge
+1
+2
+3
-3
-2
-1
Don’t form ions
Polyatomic Ions
A polyatomic ion is a group of tightly bound
atoms with an overall charge.
Hydroxide:
OH
Nitrate:
Cyanide:
CN
Carbonate:
CO32
Phosphate:
PO43
Ammonium:
NH4+
Sulfate:
NO3
SO42
You must learn the ions in Table 2.5!
What is the Formula?
beryllium & chlorine
BeCl2
potassium & oxygen
K2O
calcium & sulfate
ammonium & sulfur
CaSO4
Mg(NO3)2
(NH4)2S
barium & phosphate
Ba3(PO4)2
magnesium & nitrate
Chemical Nomenclature
• There are different systems for naming
compounds.
• Organic compounds contain carbon. These
compounds have a unique system of
nomenclature.
• We will examine the way inorganic
compounds are named.
Inorganic Nomenclature
Inorganic compounds may have traditional
names such as
H2O - water
NH3 - ammonia
CaSO4 - gypsum
...Or names related to their chemical formulas
such as:
NaCl - sodium chloride
CO2 - carbon dioxide
Na2CO3 - sodium carbonate
Ionic Compounds
Cations
• Cations formed from metal ions have the
same name as the metal:
Mg+2 - magnesium ion
Zn+2 - zinc ion
• If the metal has two or more oxidation states,
the charge is indicated by a roman numeral
Cu+ - copper(I) ion
Cu+2 - copper(II) ion
Ionic Compounds
Polyatomic cations
• There are only a few polyatomic cations usually formed by adding H+ to a molecule:
NH3 - ammonia
NH4+ - ammonium ion
CH3NH2 - methylamine
CH3NH3+ - methylammonium ion
N2H4 - hydrazine
N2H5+ - hydrazonium ion
Ionic Compounds
Anions
• Monatomic anions drop the element name
and add -ide.
Cl - chloride
O2 - oxide
• Polyatomic anions containing oxygen use the
name of the central atom with -ite or -ate.
Lower ox.#
Higher ox.#
NO2 - nitrite
NO3 - nitrate
SO32 - sulfite
SO42 - sulfate
Ionic Compounds
• Some anions add H+ to form a new anion.
They contain the word “hydrogen”:
Na2CO3 - sodium carbonate
NaHCO3 - sodium hydrogen carbonate
Na3PO4 - sodium phosphate
NaH2PO4 - sodium dihydrogen phosphate
Acids
• Acids are molecular compounds that contain
hydrogen and an anion that is dissolved in
water.
• If the anion ends in -ide, the acid has the
prefix hydro- and ends with -ic acid.
HCl - hydrochoric acid
HBr - hydrobromic acid
H2S - hydrosulfuric acid
Acids
• If the anion ends in -ite, the acid ends in
-ous acid.
HNO2 - nitrous acid
H2SO3 - sulfurous acid
• If the anion ends in -ate, the acid ends in -ic
acid.
HNO3 - nitric acid
H2SO4 - sulfuric acid
Binary Molecular
Compounds
• The element farthest to the left in the periodic
table is named first.
• If both elements are in the same group, the
lower one is named first.
• The suffix -ide is added to the second
element.
• Greek prefixes are used to indicate the
number of each element (except for mono-).
P2O5 - Diphosphorus pentoxide
Binary Molecular
Compounds
Greek
prefixes:
Number
1
2
3
4
5
6
7
8
9
10
Prefix
monoditritetrapentahexaheptaoctanonadeca-
Name That Compound
SO3
CuCl2
PbCl2
BaH2
H3PO4
(NH4)2S
N 2O 4
sulfur trioxide
copper(II) chloride
lead chloride
barium hydride
phosphoric acid
ammonium sulfide
dinitrogen tetroxide