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Structure and History of the Atom
History of the Atomic Theory:

DEMOCRITUS, (~400BC) ancient
Greek philosopher, first to suggest that
matter existed in a very small
fundamental particle. He
hypothesized that matter could be
subdivided again and again until it
would finally reach a particle which
could not be divided. This smallest
“invisible” particle he called “atomos.”
History/Atomic Theory continued:

ARISTOTLE hypothesized that matter
was “continuous” in nature. That is,
matter could be subdivided again and
gain an infinite number of times and
you would never reach a basic particle
which could not be divided. This
thought was accepted for the following
200 years. Both Democritus &
Aristotle’s views were NOT based on
experimental fact.
History/Atomic Theory continued:

JOHN DALTON (1808) was an
Englishman who was the first to
develop and publish a theory about
how atoms looked and behaved. He
thought the atom was a solid sphere,
much like a billiard ball.
Dalton’s Atomic Theory:
○ All matter is composed of extremely small
○
○
○
○
particles called atoms.
Atoms of a given element are identical in size,
mass and properties. (This was later
disproved by the discovery of isotopes).
Atoms cannot be subdivided, created or
destroyed. (This was later disproved by the
discovery of protons, neutrons and electrons).
Atoms of different elements combine in simple
whole-number ratios to form compounds.
Atoms are not made or destroyed in a chemical
reaction; they are joined, separated or
rearranged to form a new substance

J. J. THOMPSON (1897) in
his CATHODE RAY TUBE
EXPERIMENT, discovered
electrons.

CATHODE RAY TUBE
EXPERIMENT

Electron Stream w/o applied
electric field
Electron
Stream after applying field, + on
top, - on bottom


In this experiment, particles
were deflected away from
negatively charged plates &
toward positively charged
plates. Since positives
charges attract negative
charges, he concluded that
atoms contain negatively
charged particles called
electrons.
Thompson proposed that the
atom has a mixture of
positive charges and
negative charges. This
model was similar to a
chocolate chip cookie. The
“chips” represent electrons
lying in surrounding “dough”
made of protons.
History/Atomic Theory continued:

ERNEST RUTHERFORD (1911)
worked for Thomson in his lab for a
while. He performed the GOLD FOIL
EXPERIMENT which proved that the
atom as mostly space and that all of
the positive charge was located in a
very small central nucleus.
GOLD FOIL EXPERIMENT


In this experiment, alpha particles were sent through a
thin sheet of gold foil. Most of the particles went
straight through the foil. Some particles were slightly
deflected & a few were deflected back toward the
source. The reason for the deflection was the positive
charge of the nucleus.
Note: If an atom was the size of Pirate Stadium, the
nucleus would be the size of your class ring sitting on the
50 yard line.
History/Atomic Theory continued:

NIELS BOHR (1913)
Danish scientist was the
first to suggest the
“planetary model” of
the atom. Electrons
move around the
nucleus in a set path
called orbits much as
the planets orbit the sun.
Each ring or level
represents a certain
amount of energy.
Electrons farther from
nucleus have greater
energy.
History/Atomic Theory continued:


JAMES CHADWICK (1936) British
chemist who discovered the neutron.
Scientists had looked for another
subatomic particle (other than protons
and electrons) for many years but
couldn’t find them. Chadwick found out
that the neutron is approximately the
same mass as the proton but had NO
electrical charge.
ERWIN SCHRODINGER – an Austrian
physicist, along with Werner Heisenberg
and Louis de Broglie, postulated the
quantum (wave) mechanical model of
the atom which we believe is true today.
History/Atomic Theory continued:


Figure 2-6
In this model, the electrons
do not actually “orbit” the
nucleus but are found only
in definite areas, called
orbitals, based on the
amount of energy they have.
Like the particles in a light
wave, electrons move so
rapidly that their exact
location is unknown – this
creates a general “cloud” of
electrons around the
nucleus.
THE ATOM



PROTON- positively-charged particle which is
found in the nucleus; it is considerably more
massive than the electron. Proton mass is
about 1.6726 x 10-24 g.
ELECTRON- very light negatively-charged
particle which is found somewhere outside the
nucleus. It weighs only 1/1837 that of a proton,
but its negative charge is as powerful as the
positive charge of a proton.
NEUTRON- a particle found in the nucleus
which is approximately the same mass as a
proton but does not have an electrical charge
associated with it – it is neutral.
THE ATOM



ATOMS ARE ALWAYS NEUTRAL PARTICLES—that
is, they contain the same number of protons as
electrons (and it doesn’t matter how many neutrons
they have)
IONS – particles which do not have the same
number of protons and electrons so therefore they do
have an electrical charge associated with them. If
they have more protons than electrons, they will
have a positive charge. If they have more
electrons than protons, they will have a negative
charge.
# of Protons – # of Electrons ( e-) =
charge of ion
Example 2-1:
 If an element has 12 protons & 10 electrons,
what is the charge of this ion?
 If an element has 17 protons & 18 electrons,
what is the charge of this ion?
The Atom

Atomic Number – the number of protons
found in the nucleus of an atom. On the
Periodic Table it is the large whole number in
the upper corner. Notice that the only thing
which makes an element different from another
one is the number of protons it contains. The
number of protons identifies an element.
THE ATOM

Example 2-2: An atom with 6 protons is
___________________; an atom with 5
protons is __________________ and an atom
with 7 protons is ___________________.
THE ATOM


MASS NUMBER – the
number of protons and
neutrons found in the
nucleus of an atom, this
DOES NOT show on
the Periodic Chart.
Used most often to
identify an isotope of an
element.
# of neutrons + # of
protons = mass #
THE ATOM

Example 2-3:
 If Calcium’s atomic number is 20 and mass
number is 46, how many protons and neutrons are
present?
 If Neon’s atomic number is 10 and mass number
is 22, how many protons and neutrons are
present?
THE ATOM


ISOTOPE – atoms with the
same number of protons but
different numbers of neutrons.
This difference in neutrons also
results in different masses. For
instance, “regular” Carbon has
6 protons and 6 neutrons. It is
called Carbon-12. But Carbon14, the radioactive carbon used
in carbon dating, has 6 protons
and 8 neutrons.
Example 2-4:
○ What isotope has 19
protons and 20 neutrons?
○ What isotope has 26
protons and 31 neutrons?
THE ATOM

ATOMIC MASS – the mass of a single atom.
Atoms are too small to measure mass in grams,
so a system of measurement was produced.
The mass of a single atom is found by
comparing it to the standard Carbon-12 atom.
The unit of measurement for a single atom is
Atomic Mass Units (amu).

Note: Atoms are very small – typically 1 x
10-8 cm in diameter. 1.0g of lead contains 2.9
x 1021 atoms of lead. By comparison, the
earth’s entire population is only 5 x 109
people
THE ATOM

AVERAGE ATOMIC MASS – The atomic
mass that shows on the Periodic Table for
each element is really an average of the
masses of all the isotopes of that element,
weighted by their percentage of abundance.
Every element on the Periodic Table has at
least 3 isotopes; some of them have 20 or 30
or more isotopes. Atomic mass is closest to
the mass of the most abundant isotope.
THE ATOM
Equation to find the average atomic mass of an
element:
 Avg. Atomic Mass = (% isotope #1) (mass of #1) +
(% isotope #2) (mass of #2) + (% of isotope 3)
(mass of #3) +……


Example 2-5: Naturally occurring chlorine is
75.53% chlorine-35 and 24.47% chlorine-37. What
is the average atomic mass which should be
placed on the Periodic Table for the element?

Example 2-6: The element neon consists of three
isotopes with masses of 19.99, 20.99 and 21.99
amus. These three isotopes are present in nature
to the extent of 90.92%, 0.25% and 8.83%
respectively. Calculate the atomic mass of neon.
CHEMICAL CONFIGURATION


is a way of writing atoms or ions which
gives you lots of information.
Mass # is on top, element symbol in
middle, and atomic # on the bottom:
CHEMICAL CONFIGURATION
 Example
2-7:
 Write the chemical configuration for a
particle containing 5 protons, 6 neutrons
and 5 electrons.
 Write the chemical configuration for a
particle which contains 8 neutrons, has a
mass number of 14 and contains 6
electrons.
 Write the chemical configuration for a
particle which contains 18 electrons, has a
mass number of 32 and has a -2 charge.
CHEMICAL CONFIGURATION





Sometimes you are not given enough information to write
the entire chemical configuration for a particle but you
can at least identify as to what element it is and whether
it is an atom or a positively-charged or negativelycharged ion.
Example 2-8:
What particle has +2 charge and has 12 protons?
What particle has 44 protons and 44 electrons?
What particle has a -3 charge and contains 36 electrons?
ELECTRON LOCATION
 Electrons
are distributed around the nucleus
in specific ways and each electron has an
"address" that consists of 4 parts – energy
level, sublevel, orbital, & spin.
 Energy Level - indicates how far from the
nucleus the electron is located. Energy
levels closest to nucleus have least energy.
The number of electrons in any energy level
is 2n2 (n=energy level)
ELECTRON LOCATION
Example 2-12:
 How many electrons
could the 2nd
energy level hold?
 How many electrons
could the 3rd energy
level hold?

ELECTRON LOCATION


Sublevel - Each
energy level is
divided into sublevels
which are designated
by the letters s, p, d,
f.
Orbital – regions of
different shape &
energy that have a
high probability for
finding electrons
ELECTRON LOCATION





Each sublevel has a set number of orbitals related
to positions in 3-D space.
s sublevel - 1 orbital; spherical shaped
p sublevel - 3 orbitals; dumb-bell shaped
d sublevel - 5 orbitals;
f sublevel - 7 orbitals;
ELECTRON LOCATION
Each orbital holds only 2 electrons

Spin – Like Earth, an electron in an orbital can
be thought of as spinning on an
 internal axis. The electrons within an orbital must
have opposite spins.

Sublevel
s
p
d
f
# of orbitals
1
3
5
7
Total e2
6
10
14
Rules for Arranging Electrons


Aufbau Principle - electrons
are placed in orbitals of
lowest energy first. The
orbital with the lowest
energy is the 1s orbital. The
2s orbital is the next highest
in energy, then the 2p
orbitals.
Pauli Exclusion Principle an orbital may hold only two
electrons & they must have
opposite spins. Arrows are
used to indicate the direction
of the electron spin.
Rules for Arranging Electrons

Hund’s Rule - when filling orbitals of equal
energy, one electron enters each orbital until
all contain an equal amount. Then orbitals
receive a second electron.
ELECTRON ARRANGEMENT

Knowing the rules
for electron
arrangement allows
us to write a
description of where
in an atom the
electrons are
actually located.
ELECTRON ARRANGEMENT
Example 2-13: Write the electron
configuration for each of the following:

Hydrogen

helium

lithium

beryllium

boron

oxygen

ELECTRON ARRANGEMENT
ELECTRON ARRANGEMENT

Example 2-14: A certain element's electron
configuration is known to be
1s22s22p63s23p5. What element are we
describing?
ELECTRON ARRANGEMENT


An Orbital Diagram provides the same information
with more detail. In this representation, the number
of electrons in each orbital & each sublevel is
indicated. The spin of each electron is also shown.
Example 2-15: Draw the orbital notation for fluorine
(F - 9 electrons)
ELECTRON ARRANGEMENT

Example 2-16: Draw the orbital notation for
calcium (Ca - 20 electrons)
ELECTRON ARRANGEMENT

Order of Electron Fill: Just about the time you think you are
"getting the hang of this" there is an exception. After the 3p fills
up you would think the 3d would fill, but it is not so. Electrons will
fill the areas which require the least amount of energy, and it
takes less energy to fill the 4s than it does to fill the 3d. To help
you know the filling order of electrons, use the diagonal rule or
"tree".
ELECTRON ARRANGEMENT
Example 2-17:
 Write the electron configuration for iron
(Fe) based on the diagonal rule.


Draw the orbital diagram.
ELECTRON ARRANGEMENT

Example 2-18: According to the Diagonal
Rule, which is higher-energy electron:

3d or 4s
 How do you know?

5d or 6p
 How do you know?

6p or 7s
 How do you know?
ELECTRON ARRANGEMENT

Example 2-19: In the element chlorine,
which electron is the highest-energy
electron?

Which electron is physically farthest from
the nucleus?

How many unpaired electrons are in
chlorine?
PERIODIC TABLE and ELECTRON
CONFIGURATION

An element’s placement on the Periodic Table is
related to the arrangement of electrons in an atom. A.
Each horizontal row (called a period) corresponds
to an energy level. The periodic table can be broken
into sections that represent the orbitals. And if you
read the periodic table like a paragraph from left to
right, the order of electron fill is provided.

Note: The "s" and "p" electrons will have an energy
level number same as the period they are on. The "d"
electrons will have an energy level ONE LESS than
the period they are on & "f" electrons will have an
energy level two less than the period they are on
PERIODIC TABLE and ELECTRON
CONFIGURATION
PERIODIC TABLE and ELECTRON
CONFIGURATION







Example 2-20: Use the Periodic Chart to
read the electron configuration for Na (#11)
Example 2-21: Use the Periodic Chart to
write the electron configuration for …
Al (#13)
Ar (#18)
Ni (#28)
Zn (#30)
Br (#35)
PERIODIC TABLE and ELECTRON
CONFIGURATION

Electron configurations are sometimes written using noblegas notation. In this form of notation, instead of writing out
the entire series of energy levels and sublevels, you can
use an abbreviation. First identify the noble-gas that
comes before the element. Write the symbol of the noblegas in brackets & continue writing the energy level,
sublevel, and number of electrons until you reach your
element.

For example: Electron Configuration of
○ Si 1s22s22p6 3s23p2

Noble-gas Notation of
○ Si
[Ne] 3s23p2
PERIODIC TABLE and ELECTRON
CONFIGURATION
Example 2-22: Write the electron
configuration using noble-gas notation
for….
 O (#8)
 P (#15)
 Cu (#29)
 Br (#35)
