Unit 2 – Atoms: The Building Blocks of Matter
I.
The Atom: From Philosophical Idea to Scientific Theory
A.
B.
C.
The Philosophers
1.
Democritus (~400 BC) – Greek philosopher; proposed the
existence of atoms (“atomos”…indivisible)
2.
Another Greek philosopher of the time, Aristotle, believed that all
matter was continuous…no such thing as atoms. Aristotle’s opinion
was accepted for over 2000 years!
Foundations of Atomic Theory
1.
Law of Conservation of Mass – mass is neither created nor
destroyed during ordinary chemical reactions (Antoine
Lavoisier…”The Father of Chemistry”)
2.
Law of Definite Proportions – a chemical compound contains the
same elements in exactly the same proportions by mass regardless
of the size or source of the sample (Joseph Proust)
3.
Law of Multiple Proportions- if 2 or more different compounds are
composed of the same 2 elements, the ratio of the masses of the
2nd element to 1.0 g of the 1st element is always a small whole
number (John Dalton)
John Dalton’s Atomic Theory (1808)
1.
Basic Points
a.
b.
c.
d.
e.
2.
The Law of Conservation of Mass is explained by point e
3.
The Law of Multiple Proportions is explained by point d
1
D.
Modern Atomic Theory
1.
Some of Dalton’s points were not exactly accurate and have been
modified throughout the years
2.
Two important concepts remain unchanged:
a.
b.
II.
The Structure of the Atom
A.
Discovery of the Electron
1.
Cathode Rays and Electrons (Late 1800s)
a.
b.
When current was passed through the cathode ray tube, the
surface of the tube opposite the cathode glowed
1.
Hypothesized that the glow was caused by a stream
of particles called “cathode rays”
2.
B.
Rays seemed to travel from the cathode to the anode
c.
JJ Thomson (1897) –
d.
Robert Milliken (1909) –
Discovery of the Atomic Nucleus
1.
The Gold Foil Experiment (1911)
a.
Carried out by Ernest Rutherford, Hans Geiger, and Ernest
Marsden (among others)
b.
The Experimental Design
1.
A thin, gold foil was bombarded with alpha particles
which were positively-charged and four times the
mass of the hydrogen atom
2
2.
2.
3.
The foil was surrounded by a screen that would detect
where the alpha particles hit after passing through the
foil
The Results
a.
Most of the alpha particles went straight through the foil or
were only slightly deflected (this was expected)
b.
HOWEVER,
Rutherford’s Explanation (1913)
a.
b.
c.
C.
Particle
Composition of the Atomic Nucleus
1.
Except for the simplest type of hydrogen atom, all atomic nuclei are
made up of two kinds of particles: protons and neutrons
2.
Electrons travel around the nucleus, occupying the majority of the
atom’s volume
3.
Summary of Subatomic Particles
Symbols
Relative
Electric
Charge
Mass
Number
(A)
Relative
Mass
(amu)
Actual
Mass
(g)
Electron
Proton
Neutron
D.
The Sizes of Atoms
1.
The region of the atom occupied by the electron can be thought of
as
3
2.
III.
Radius of an atom -
Counting Atoms
A.
B.
Atomic Number (Z) –
1.
Every element has a different atomic number…this identifies the
element!
2.
Found on the periodic table…in each element block near the
symbol of the element
3.
In a neutral atom, there will be
Mass Number (A) –
1.
MUST BE GIVEN TO YOU!!! NOT on the periodic table!
2.
Atoms of the same element may have different numbers of
neutrons and, therefore, different mass numbers
a.
These atoms are referred to as isotopes…
b.
Although isotopes have different masses, they do not
significantly differ in chemical behavior
c.
Example - Hydrogen
Protium (Most
Deuterium
common form of H…
“Heavy Hydrogen”
99.985%)
(0.015%)
Tritium
“Radioactive
Hydrogen” (trace)
#
protons
#
neutrons
Mass
Number
4
C.
Designating Isotopes
1.
Nuclear Symbol
2.
Hyphen Notation
3.
Using either designation, one can determine the number of protons,
neutrons, and electrons…
4.
Nuclear
Symbol
a.
# protons = atomic number (Z) = # electrons
b.
# neutrons =
Examples
Hyphen
Notation
#
Protons
#
Neutrons
#
Electrons
20
21
20
35Cl
131I
Uranium-235
5.
D.
Nuclide –
Relative Atomic Masses of the Elements
1.
Because atoms are so small, using mass in units of grams is
inconvenient and hard to understand
1 atom hydrogen-1 = 1.673 x 10-24 grams
2.
Scientists set up a relative scale of atomic masses…known as the
atomic mass scale
a.
Carbon-12 was arbitrarily assigned a mass of 12 atomic
mass units (amus)
5
E.
b.
One atomic mass unit (amu), or 1 amu, is
c.
The mass of all other nuclides are determined by comparing
each one with the mass of a carbon-12 atom
Average Atomic Masses of the Elements
1.
The mass found in each element block on the periodic table must
account for all isotopes of that element that are in existence and
how abundant they are in nature
2.
Average atomic mass -
a.
To do this, use the following formula for each isotope:
(% occurrence  100) (mass)
b.
3.
Add all values together to get the average atomic mass
In my examples, an element’s atomic mass will be rounded to one
place past the decimal before it’s used in a calculation
6
F.
Relating Mass to Numbers of Atoms
1.
2.
3.
IV.
The Mole
a.
All scientists around the world use the unit “mole” when
describing the amount of substance present in a sample
b.
The mole is a counting unit…just like “pair” means “2,”
“dozen” means “12,” and “gross” means “144.”
c.
Mole (mol) -
Avogadro’s Number
a.
Named after the 19th century Italian scientist Amedeo
Avogadro…he helped explain the relationship between mass
and number of atoms
b.
The number of particles in exactly one mole of a pure
substance is 6.0221367 x 1023 particles…we will use
c.
One mole of ANYTHING will consist of 6.02 x 1023 particles
MOLE DAY – OCTOBER 23…6:02 AM TO 6:02 PM
Basic Bonding
A.
Atoms are held together (forming compounds) by chemical bonds;
there are two types
B.
1.
Ionic Bonding -
2.
Covalent Bonding -
Determining the Type of Bonding Based on the Chemical Formula
7
1.
If the compound is ionically bonded…
2.
If the compound is covalently bonded…
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