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Unit 2: Atoms, Molecules, and Ions
PART 1
An overview of the evolution of the
Atomic Model from John Dalton to the
Modern Theory.
But first, a little ancient history:
http://www.meta-synthesis.com/webbook/35_pt/pt_database.php
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.
The primary difference between
Dalton’s theory and previous
ones was that Dalton’s was
based on reproducible
laboratory evidence.
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 based his theory on the work of a number of
scientists who developed the following laws:
1. The 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
+
S
→ ZnS
65.4g +
32.1 g = 97.5 g
2. The 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 released 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 E 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)
3. The 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!!
4. The 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!
C
O
C/O
Ratio
CO
12g
16g
16/12
= 1.3
1.3/1.3
=1
CO2
12g
32g
32/12
=2.6
2.6/1.3
=2
CO3
12g
48g
48/12
=4.0
4.0/1.3
=3
5. Guy-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
+
+
• 2H2 +
• 2L +
O2
1L
→
=
SO2
1L
O2
1L
→
=
2H2O
2 Liters
• 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.
“Billiard Ball Model”
Discovery of the electron - 1897
• JJ Thomson studied electrical discharges in
partially, evacuated tubes called cathode –ray
tubes
Diagrams:
Initial:
After:
• J.J. Thomson- In a series of experiments with
a cathode ray tube (CRT) in 1897, discovered
that negatively charged particles of matter
could be removed from atoms.
• This indicated that atoms were not indivisible
but were composed of even smaller particles.
• The discovered particle was the
electron.
• 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)
One of the key findings was that the charge of the
electron was negative!!
Thomson’s Concept of Atoms
• Thomson proposed a “new and improved”
model using Dalton’s theory as a foundation
(remember – The idea that matter was
composed of atoms was not universally accepted
(Mendeleev)
• 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”
A brief aside to discuss 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!!)
• Some materials are radioactive and may produce
different types of radiation.
• These “particles” suggested that there may be
other parts to an atom.
Some types of radiation
Type of Radiation
Symbol
Definition
Alpha particle
α
Like a helium atom
(2 protons fused
together); +2
charge; mass is 7300
times an e-
Beta particle
β
High 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 Kleenex. What would you
expect to happen?
Java – Rutherford scattering
http://www.mhhe.com/physsci/chemistry/essen
tialchemistry/flash/ruther14.swf
Gold foil experiment
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)
Ernest Rutherford - 1911
• Rutherford discovered that the positive
charge 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.
Diagram of Rutherford’s Atom
“Empty Space Model”
Discovery of protons
Henry Moseley – 1913 –
Discovered that the number
of positive charges in an atom
is equal to the element
number. This indicated that
there was a particle in the
nucleus that was the source
of + charge. In 1920,
Rutherford named the
particle “proton”.
Discovery of neutrons
• James Chadwick – 1932 – Explained
the difference between the observed
mass of atomic nuclei and the
number of + charges (also
considering spin) by proposing the
presence of particles with masses
similar to those of protons but with
no charge - the neutron.
The Bohr model of the atom
Niels Bohr – 1913 – Considered
that the Rutherford model of
the atom was unstable and the
spectra of atoms (discrete bands
of light absorbed and emitted by
atoms) to propose a new model
with the electrons confined to
specific energy levels
(sometimes called shells or
orbits).
Neil Bohr’s Concept of the Atom - 1913
• Bohr proposed that electrons are
arranged in definite energy levels(shells)
and follow a definite orbit. Lowest
energy level was nearest the nucleus
• 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
Modern Concept of the Atom – 1920’s to
present
• The modern theory states that the electrons have
wave-like properties as they travel 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. (Much more on this
later!!!!)
Diagram of the Modern Atom
“wave mechanical” model
Evolution of
the atomic
model
We now have a workable model of the atom: It is composed
of three particles, proton, neutron and electron.
Name
Symbol
Mass
Charge
Location
Proton
p+
1.67x10-27 kg
+
nucleus
Neutron
n0
1.67x10-27 kg
None
Nucleus
Electron
e-
9.1x10-31 kg
-
Energy
levels
• The mass of a neutron is actually very slightly more
than that of a proton however, in chemistry we
generally consider them to be the same.
• We use a convenient unit to express this mass, the
amu (atomic mass unit).
• An amu is defined as 1/12 the mass of a carbon-12
atom.
• We generally consider the masses of both p+ and n0
to be 1 amu.
• The mass of an e- is so much less than the other
particles that we considered it to be zero in
calculating the mass of an atom.
So the mass of an atom, in amu’s, is simply the
number of protons plus the number of neutrons.
This is sometimes called “mass number”
atomic mass = #p+ + #n0
The atomic number for an atom is simply its number
of protons.
atomic number = #p+
The total charge on an atom is determined by the
number of p+ and e-. Since these particles have
charges of equal magnitude and opposite sign, their
charges cancel. When an atom has the same
number of p+ and e- , the total charge must be zero
– a neutral atom.
charge = #p+ - #e-
• An ION is an atom with different numbers of
p+ and e-.
• More p+ - positive charge – CATION
• More e- - negative charge – ANION
Note: The only way to form an ion is by gaining
or losing electrons.!!!!!!
Java – build an atom
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
• 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, 18 electrons and a mass
number of 40, and then write the formula
of a different isotope of this element.
IONS
• ION: an atom with a positive or negative
charge.
• CATION: a positive ion
Na+1
•ANION:
Cl-1
Cu+2
Al+3
a negative ion
C2H3O2-1
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. many metals have 1 – 3 electrons in the
outer shell
g. all metals lose electrons during chemical
change. Cations!
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.
Anions
3. Metalloids
• Also called semimetals or semiconductors
• Metalloids have properties of both metals and
nonmetals.
4. Noble or Inert Gases
a. all are gases (no way!)
b. the noble gases generally do not form
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.(1)
2. ALKALI EARTH METALS – reactive (but not as
much as Group IA) metals that form +2
ions. These are the Group IIA (2). These
are often called the Fireworks Metals!!
1. HALOGENS – Group VIIA (17). The word
Halogen means “salt forming”; all halogens
are very reactive and react with metals to
form salts.
Gain 1 e- to form -1 ions.
1. NOBLE OR INERT GASES – Group VIIIA (18);
nonreactive gases/elements
Chemical reactions are represented by
both words and symbols
•WORD:
Zinc and sulfur yields zinc sulfide
• EQUATION
Chemical materials and their reactions are
usually designated by formulas and equations.
This is the language of chemistry
Formulas
…learning the language
• 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
BINARY IONIC COMPOUNDS (FORMULAS)
a. Binary compounds are composed of 2
elements
b.The components of a binary ionic compound
are a monoatomic cation and a monoatomic
anion. (what does monoatomic mean?)
c. Binary compounds end in the suffix –ide
d.Ionic compounds must be 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. Methods for determining subscripts:
visualization
LCM
Criss-cross-reduce
Writing Binary Ionic Formulas
Ex: Potassium chloride
• Calcium bromide
• Iron III oxide
• Calcium sulfide
Rules for writing ionic formulas
1.Write the symbol for the + ion first. Do not write the
charges in the formula.
2.Use subscripts to show the number of each ion
required to give a net charge of zero.
3.Subscripts give the smallest ratio between the ions.
4.Use parentheses only with polyatomic ions to show
more than one of them in the formula.
5.Do not write “1” as a subscript. No subscript means
1.
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. Go by
your ion sheet.
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, use the charges to determine subscripts
e. Third, use parenthesis whenever a polyatomic ion
needs to be taken 2, 3 or 4 times.
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
A simple method for naming binary
covalents!
Number Name Number Name-ide
Never start with mono.
Always end with –ide.
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:
Chem-Nerds of the world unite and
identify these:
• H2
Mg
• SO4 -2
Mg+2
• CaCl2
• Cu+1
• Acid names and formulas…good idea to know
these!!
Names of Acids That Do Not Contain Oxygen
Acid
Name
HF
hydrofluoric acid
HCl
hydrochloric acid
HBr
hydrobromic acid
HI
hydroiodic acid
HCN
hydrocyanic acid
H2S
hydrosulfuric acid
Names of Some Oxygen-Containing Acids
Acid
Name
HNO3
nitric acid
HNO2
nitrous acid
H2SO4
sulfuric acid
H2SO3
sulfurous acid
H3PO4
phosphoric acid
HC2H3O2
acetic acid