Atoms, Molecules, and Ions
Chemistry Timeline #1
B.C.
400 B.C. Demokritos and Leucippos use the term "atomos”

2000 years of Alchemy
1500's
 Georg Bauer: systematic metallurgy
 Paracelsus: medicinal application of minerals
1600's
Robert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of
elements
1700s'
 Georg Stahl: Phlogiston Theory
 Joseph Priestly: Discovery of oxygen
 Antoine Lavoisier: The role of oxygen in combustion, law of conservation of
mass, first modern chemistry textbook
Chemistry Timeline #2
1800's
Joseph Proust: The law of definite proportion (composition)
 John Dalton: The Atomic Theory, The law of multiple proportions
Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules
Amadeo Avogadro: Molar volumes of gases
Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements
 Dmitri Mendeleyev: The periodic table
 J.J. Thomson: discovery of the electron
 Henri Becquerel: Discovery of radioactivity
1900's
 Robert Millikan: Charge and mass of the electron
 Ernest Rutherford: Existence of the nucleus, and its relative size
 Meitner & Fermi: Sustained nuclear fission
 Ernest Lawrence: The cyclotron and trans-uranium elements
Dalton’s Atomic Theory (1808)
 All matter is composed of extremely
small particles called atoms
 Atoms of a given element are
identical in size, mass, and other
properties; atoms of different
John Dalton
elements differ in size, mass, and
other properties
 Atoms cannot be subdivided, created, or destroyed
 Atoms of different elements combine in simple
whole-number ratios to form chemical compounds
 In chemical reactions, atoms are combined,
separated, or rearranged
Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in
size, mass, and other properties; atoms of
different elements differ in size, mass, and
other properties
Modern theory states:
Atoms of an element have a characteristic
average mass which is unique to that
element.
Modern Atomic Theory #2
Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed
in ordinary chemical reactions. However, these
changes CAN occur in nuclear reactions!
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube
to deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Thomson’s Atomic Model
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Mass of the Electron
1909 – Robert Millikan
determines the mass of
the electron.
The oil drop apparatus
Mass of the
electron is
9.109 x 10-31 kg
Conclusions from the Study of
the Electron
 Cathode rays have identical properties regardless
of the element used to produce them. All elements
must contain identically charged electrons.
Atoms are neutral, so there must be positive
particles in the atom to balance the negative
charge of the electrons
 Electrons have so little mass that atoms must
contain other particles that account for most of
the mass
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are
recorded
Try it Yourself!
In the following pictures, there is a target hidden by
a cloud. To figure out the shape of the target, we
shot some beams into the cloud and recorded where
the beams came out. Can you figure out the shape of
the target?
The Answers
Target #1
Target #2
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
“Like howitzer shells bouncing off
of tissue paper!”
Conclusions:
 The nucleus is small
 The nucleus is dense
 The nucleus is positively charged
Atomic Particles
Particle
Charge
Electron
-1
9.109 x 10-31
Electron
cloud
Proton
+1
1.673 x 10-27
Nucleus
0
1.675 x 10-27
Nucleus
Neutron
Mass (kg)
Location
The Atomic
Scale
 Most of the mass of the
atom is in the nucleus
(protons and neutrons)
 Electrons are found
outside of the nucleus (the
electron cloud)
 Most of the volume of
the atom is empty space
“q” is a particle called a “quark”
About Quarks…
Protons and neutrons are
NOT fundamental particles.
Protons are made of
two “up” quarks and
one “down” quark.
Neutrons are made of
one “up” quark and
two “down” quarks.
Quarks are held together
by “gluons”
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon = 12.011
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Atomic Number
Atomic number (Z) of an element is the
number of protons in the nucleus of each atom
of that element.
Element
# of protons
Atomic # (Z)
6
6
Phosphorus
15
15
Gold
79
79
Carbon
Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide
p+
n0
e-
Mass #
Oxygen - 18
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Molecules
Two or more atoms of the same or different
elements, covalently bonded together.
Molecules are discrete structures, and their
formulas represent each atom present in the
molecule.
Benzene, C6H6
Covalent Network Substances
Covalent network substances have covalently
bonded atoms, but do not have discrete
formulas.
Why Not??
Graphite
Diamond
Ions
 Cation: A positive ion
• Mg2+, NH4+
 Anion: A negative ion
 Cl-, SO42 Ionic Bonding: Force of attraction between
oppositely charged ions.
 Ionic compounds form crystals, so their
formulas are written empirically (lowest whole
number ratio of ions).
Periodic Table with Group Names
The Properties of a Group:
the Alkali Metals
Easily lose valence electron
(Reducing agents)
React violently with water
Large hydration energy
React with halogens to form salts
Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions
H+
Li+ Na+
K+
Predicting Ionic Charges
Group 2: Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 13: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
Caution! C22- and C4are both called carbide
Group 14: Loses 4
electrons or gains
4 electrons
Predicting Ionic Charges
N3- Nitride
P3- Phosphide
As3- Arsenide
Group 15: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2- Oxide
S2- Sulfide
Se2- Selenide
Group 16: Gains 2
electrons to form
2- ions
Predicting Ionic Charges
F1- Fluoride
Br1- Bromide
Cl1-Chloride
I1- Iodide
Group 17: Gains 1
electron to form
1- ions
Predicting Ionic Charges
Group 18: Stable
Noble gases do not
form ions!
Predicting Ionic Charges
Groups 3 - 12: Many transition elements
have more than one possible oxidation state.
Iron(II) = Fe2+
Iron(III) = Fe3+
Predicting Ionic Charges
Groups 3 - 12: Some transition elements
have only one possible oxidation state.
Zinc = Zn2+
Silver = Ag+
Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
(
Ba NO3 ) 2
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Not balanced!
Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
( NH4+) SO42-
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
2
Not balanced!
Writing Ionic Compound Formulas
Example: Iron(III) chloride
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
Fe3+ Cl-
3
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum sulfide
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
3+
Al
2
2S
3
Not balanced!
Writing Ionic Compound Formulas
Example: Magnesium carbonate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges
are balanced.
Mg2+ CO32They are balanced!
Writing Ionic Compound Formulas
Example: Zinc hydroxide
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
2+
Zn
3. Balance charges , if necessary,
using subscripts. Use parentheses
if you need more than one of a
polyatomic ion.
( OH- )2
Not balanced!
Writing Ionic Compound Formulas
Example: Aluminum phosphate
1. Write the formulas for the cation
and anion, including CHARGES!
2. Check to see if charges are
balanced.
3+
Al
PO4
3-
They ARE balanced!
Naming Ionic Compounds
• 1. Cation first, then anion
• 2. Monatomic cation = name of the element
• Ca2+ = calcium ion
• 3. Monatomic anion = root + -ide
• Cl- = chloride
• CaCl2 = calcium chloride
Naming Ionic Compounds
(continued)
Metals with multiple oxidation states
 some metal forms more than one cation
 use Roman numeral in name
PbCl2
Pb2+ is the lead(II) cation
PbCl2 = lead(II) chloride
Naming Binary Compounds





Compounds between two nonmetals
First element in the formula is named first.
Second element is named as if it were an anion.
Use prefixes
Only use mono on second element P2O5
CO2
CO
N2O
= diphosphorus pentoxide
= carbon dioxide
= carbon monoxide
= dinitrogen monoxide