Atoms, Molecules, and Ions
The Evolution of the Atomic Model
(from John Dalton to the Modern
Theory – just the big kahunas!)
1. John Dalton’s Atomic Theory – 1803
Dalton stated a group of assumptions to
explain the nature and behavior of
chemical systems. These became known
as Dalton’s Atomic Theory and he
proposed this the year 1803.
Dalton’s Atomic Theory
Four Assumptions:
1. All substances are composed of small, dense
particles called ATOMS.
2. Atoms of a given substance are identical in
mass, size and shape.
3. An atom is the smallest part of an element
that enters into a chemical reaction.
4. Molecules are produced by a combination of
atoms.
Dalton’s Atomic Theory
• Dalton’s Atomic Theory is often called the
Billiard Ball Model. His theory, however, fails
to explain many different types of behavior in
chemical reactions.
• It was through Dalton’s work that evidence of
the following two subatomic particles were
discovered:
ELECTRON: negatively charged particles.
PROTON: positively charged particles.
“Billiard Ball Model”
2. Thomson’s Concept of Atoms
• JJ Thomson, early 1900’s
• Thomson proposed a “new and improved”
model using Dalton’s theory as a foundation
(remember – people thought Dalton was
crazy when he said in the early 1800’s that
everything was made of atoms!!)
• Thomson proposed that atoms consist of a
solid bulk of positive charge with electrons
dispersed throughout. His model is known
as the Plum Pudding Model.
Diagram of Thomson’s Atom
“Plum Pudding Model”
3. Ernest Rutherford - 1911
• Rutherford discovered that the positive
charge (the proton) and the mass was
concentrated in the center of the atom
(called the NUCLEUS).
• He postulated that the electrons were
moving at fast speeds around the nucleus
but were contained by a certain
boundary.
3. Diagram of Rutherford’s Atom
“Empty Space Model”
4. Neil Bohr’s Concept of the Atom 1913
• Bohr proposed that electrons are
arranged in definite energy
levels(shells) and follow a definite
orbit.
• Bohr’s concept or model is known as
the “satellite” or “solar system”
model of the atom .
Diagram of Bohr’s Atom
• Solar system model
5. Modern Concept of the Atom –
1920’s to present
• The modern theory states that the electrons have
wave-like properties as they orbit around the
central nucleus. The + charged nucleus is
surrounded by electrons with definite energy
levels (called orbitals).
• The paths of the e- are described in terms of the
probability of being found in certain regions. The
e- do not follow a prescribed path.
• The modern concept of the atom is known as the
“wave-mechanical” model of the atom.
Diagram of the Modern Atom
“wave mechanical” model
Evolution of
the atomic
model
Chemical reactions are represented by
both words and symbols
•WORD:
Zinc and sulfur yields zinc sulfide
• EQUATION
Law of Conservation of Mass
• Or the Law of Conservation of Matter
• The law states that in ordinary chemical
reactions, the mass of the system remains
constant.
Zn
+
65.4g +
S
→ ZnS
32.1 g = 97.5 g
Law of Conservation of Energy
• The heat lost by the system (reaction) is equal
to the heat gained by the surroundings (or, in
ordinary chemical reactions, the energy of the
system remains constant).
ENDOTHERMIC: heat energy is absorbed during
a chemical reaction (it gets colder)
EXOTHERMIC: heat energy is lost during a
chemical reaction (it gets hotter!)
95% of all reactions are exothermic!!
• Heat + nitrogen + oxygen → nitric oxide
• q + N2
+ O2
→ 2NO
(q is one of the symbols used for heat or energy)
If q is written on the left side, this indicates the
reaction is endothermic (since energy is being
absorbed).
If q is written on the right side of the equation,
the reaction is exothermic (since heat is being
released to the surroundings)
Law of Definite Composition
• When elements combine and form specific
compounds, they do so in definite proportions by
mass.
Zn
+ S
→ ZnS
65.4g + 32.1 g = 97.5 g
Zinc and sulfur will always combine in a definite
fixed ratio of 65.4 parts to 32.1 parts by mass.
If Zn or S were present in any other ratio, the
one in excess would remain unchanged or
unused. The excess would remain unreacted!!
Law of Multiple Proportions
• When two elements combine and form more
than one compound, the masses of one
element that combine with a fixed mass of the
other are in the ratio of small, whole
numbers.
• Look at the formulas of C to O in your notes:
notice that in none of those formulas do you
see C1.2O2.67 or C.98O3.11
• Formulas are always WHOLE NUMBERS!
Gay-Lussac’s Law of Combining
Volumes
• Gases react chemically with a volume of small,
whole numbers!!
• 1 volume H2 + 1 volume Cl2 → 2 volumes HCl
• S
1L
+
+
O2
1L
→
=
SO2
1L
• 2H2 + O2 → 2H2O
2L + 1 L =
2 Liters
(it’s all about the small, whole numbers baby!)
Discovery of the Electron
• JJ Thomson studied electrical discharges in
partially, evacuated tubes called cathode –ray
tubes (does anyone know the connection to
our TV’s?) NOTE: this is the after picture!!!
More Cathode Ray fun!!!
• Since the ray was attracted to the positive
electrode, Thomson called the stream of
particles electrons.
• Thomson also determined the charge-to-mass
ratio of the electron to be:
e/m = -1.76 x 108 Coulombs/gram
(where e = charge of the e- in Coulombs and m =
mass in grams)
WHY is the e/m ratio important? It showed
that the electron has a NEGATIVE CHARGE!
Radioactivity
• The French scientist Henri Becquerel found
that a piece of Uranium produced its image on
a photographic plate. So, as a third arm was
growing out of his ribs, he figured out that this
U had some weird energy coming from it
(kind of like a
chicken patty sandwich)
• RADIOACTIVITY: the spontaneous
emission of radiation (particles with lots
of energy!!)
Three Types of Radiation
Type of Radiation
Symbol
Definition
Alpha particle
α
Beta particle
β
Like a helium atom
(2 protons fused
together); +2
charge; mass is 7300
times an eHigh speed/High
energy electrons
Gamma ray
γ
“high energy” light;
very damaging to
our DNA; most
damaging of the 3
Discovery of the Nucleus
• ERNEST RUTHERFORD – most famous for his
discovery of the nucleus.
• Rutherford used a radioactive source that
emitted alpha (α)particles (these were his
“bullets”), a piece of VERY THIN gold foil, a
fluorescent screen, and a lead block.
• His experiment would be similar to if you took
your genuine Red Ryder BB gun and shot BB’s
at a brick wall. What would you expect to
happen?
Diagram of Rutherford’s Experiment
Rutherford’s Interpretation
Rutherford’s 3 Assumptions
1. The slightly deflected α particle had a close
encounter with the “positive center” of the
atom.
2. Most of the atom is empty space (because
most of the α particles passed through)
3. The alpha particles that were completely
deflected hit head on with the nucleus
(because likes repel – both were + charged)
Major Subatomic Particles
Particle
Mass
Charge
Electron
9.11 x 10-28 g
-1
Proton
1.67 x 10-24 g
+1
1.68 x 10-24 g
0
(in nucleus)
Neutron
(in nucleus)
A,Z,X Method (atomic number, atomic
weight, etc.)
MASS NUMBER
(or atomic weight) – the
Total number of protons and neutrons
In the atom.
SYMBOL
ATOMIC NUMBER – the number of
protons (and also = to the number of
electrons if the atom is neutral!)
What if it is an ION ?
• An ion is an atom with a net positive or a net
negative charge.
+1
A positive ion means that the element has lost
electrons (in this case 1 electron). Na now has one
more proton than electrons. A negative ion means
that element has gained electrons (it now has more
electrons than protons).
Isotopes
• An isotopes are atoms with the same number
of protons but different numbers of neutrons.
• Examples of the isotopes of Hydrogen:
•
protium
deuterium
tritium
•
•
p+ =
no =
p+ =
no =
p+ =
no =
Write the symbol of the element
which has 20 protons and a mass
number of 40, and then write the
formula of an isotope of this element.
Formulas
…can you say pay attention to this!
• CHEMICAL FORMULA: The symbol
for the elements are used to indicate
the types of atoms present and the
subscripts are used to show the
relative numbers of atoms.
• Example of a chemical formula:
CO2
STRUCTURAL FORMULAS
• A formula showing the individual bonds (using
lines to show the bonds).
•
H2O
CH4
IONS
• ION: an atom with a positive or negative
charge.
• CATION: a positive ion
Na+1
•ANION:
Cu+2
Al+3
a negative ion
C2 H3 O2 -1
BO3-3
SO4-2
Introduction to the fabulous PERIODIC
TABLE!!!
• Identify the following sections on a blank
periodic table:
– Metals
nonmetals
– Hydrogen
alkali metals
– Halogens
metalloids
– Al family
Carbon family
– Oxygen family
– Rare Earth elements
– Lanthanide Series
noble gases
alkali earth metal
transition met.
Nitrogen family
Groups IA-VIIIA
Actinide series
Four Classifications of Elements
1. METALS
a. shiny luster
b. good conductors; poor insulators
c. malleable – can be hammered into
thin sheets
d. ductile – can be drawn into a thin
wire
e. All are solids except Ga and Hg
Properties of Metal (con’d)
f. all metals have 1 – 3 electrons in the outer
shell
g. all metals lose electrons during chemical
change.
2. Nonmetals
a. very brittle but pretty colors
b. poor conductors; good insulators
c. nonmetals are solids, liquids and gases
d. all nonmetals have 5 – 7 e- in the outer
shell
e. nonmetals gain e- in chemical reactions
3. Metalloids
• Also called semimetals
• Metalloids have properties of both metals and
nonmetals.
4. Noble or Inert Gases
a. all are gases (no way!)
b. the noble gases form NO compounds
c. they are unreactive – they have 8 electrons
in the outer shell (except Helium which only
has and only needs 2 electrons)
SPECIFIC TYPES OF METALS
1. ALKALI METALS – very active metals that
form ions with a +1 charge. Group IA.
2. ALKALI EARTH METALS – reactive (but not as
much as Group IA) metals that form +2
ions. These are the Group IIA. These are
often called the Fireworks Metals!!
1. HALOGENS – Group VIIA. The word Halogen
means “very active nonmetal”; all
halogens react with metals to form salts
containing ions with a -1 charge.
2. NOBLE OR INERT GASES – Group VIIIA ;
nonreactive gases/elements
a.
b.
c.
d.
BINARY IONIC COMPOUNDS (FORMULAS)
Binary compounds are composed of 2
elements
The components of a binary ionic compound
are a monoatomic cation and a monoatomic
anion. (what does monoatomic mean?)
Binary compounds end in the suffix –ide
Ionic compounds are electrically neutral.
Writing Formulas
e. In writing a formula, we must exactly balance
the positive charge of the cation with the
negative charge of the anion (the net charge
should be 0).
f. We use the Crisscross
method”
Ie. Potassium chloride
(makes you want to “…jump, jump”)
Writing Binary Ionic Formulas
• Calcium bromide
• Iron III oxide
• Calcium sulfide
a. Binary ionic compounds are named by
writing the name of the cation followed by
the anion (ending in –ide).
b. When the cation has more than one possible
ionic charge, it is important to use the Roman
numeral..
c. DON’T overkill the use of the Roman
numerals. The representative (Group A)
metals only have one charge so a Roman
numeral is NOT needed.
Naming Binary Ionic Compounds
• AlI3
• FeO
• Cu2S
• CaSe
3. Ternary Ionic Compounds
a. Ternary ionic compounds contains atoms of
three or more different elements.
b. Ternary ionic compounds usually contain one
or more polyatomic ions.
c. First, write down the symbol of the ions
d. Second, “crisscross”
the charges.
e. Third, use parenthesis whenever a
polyatomic ion needs to be taken 2, 3 or 4
times.
(without jumping)
Examples of Ternary Ionic Compounds
Calcium nitrate
Potassium sulfate
Magnesium hydroxide
Ammonium sulfide
Examples…ternary ionic
• Aluminum bicarbonate
• Chromium III benzoate
• LiCN
• Sr(H2PO4)2
More examples…ternary
• NH4C2H3 O2
• Fe(ClO3)3
4. Binary Molecular Compounds
a. Composed of two or more nonmetallic
elements
b. Most of the elements that form binary
molecular compounds are not charged atoms
(not ions).
c. When two nonmetallic elements combine,
they often do so in more than one way.
CO CO2 CO3 CO4
Binary Molecular Compounds
d. Binary molecular compounds end in –ide.
e. Greek prefixes are used to show how many
atoms of each element are present in each
molecule.
f. Note that the vowel at the end of the prefix
“mono” is dropped when the name of the
element begins with a vowel; monoxide not
monooxide!
g. The prefix mono is omitted if there is just a
single atom of the first element in the name.
Greek Prefixes
GREEK PREFIX
NUMBER
1
2
3
4
5
6
7
8
9
10
Examples of Binary Molecular…
• N2O
• PCl3
• SF6
• N4O7
Binary Molecular Compounds
• Dinitrogen tetrahydride
• Diphosphorus trioxide
• Carbon tetrachloride
• Nonapotassium monophosphide 
Types of Combinations of Atoms
1. FREE ELEMENT – one type of atom; no charge
example:
2. MONOATOMIC ION – one type of atom; charged
example:
3. COMPOUNDS/MOLECULES: more than one type
of atoms; no charge
example:
4. POLYATOMIC IONS: more than one type of atom
but with a charge.
example:
Nerds of the world unite and identify
these:
• H2
• SO4 -2
• CaCl2
• Cu+1
• Acid names and formulas…good idea to know
these!!