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Atomos: Not to Be Cut
The History of Atomic Theory
Atomic Models

This model of the
atom may look
familiar to you. This is
the Bohr model. In
this model, the
nucleus is orbited by
electrons, which are
in different energy
levels.

A model uses familiar ideas to
explain unfamiliar facts
observed in nature.

A model can be changed as
new information is collected.
 The
atomic
model has
changed
throughout the
centuries,
starting in 400
BC, when it
looked like a
billiard ball →
Who are these men?
In this lesson, we’ll learn
about the men whose quests
for knowledge about the
fundamental nature of the
universe helped define our
views.
Democritus

This is the Greek
philosopher Democritus
who began the search for
a description of matter
more than 2400 years
ago.
 He asked: Could
matter be divided into
smaller and smaller
pieces forever, or was
there a limit to the
number of times a
piece of matter could
be divided?
400 BC
Atomos



His theory: Matter could
not be divided into
smaller and smaller
pieces forever, eventually
the smallest possible
piece would be obtained.
This piece would be
indivisible.
He named the smallest
piece of matter “atomos,”
meaning “not to be cut.”
Atomos


To Democritus, atoms
were small, hard
particles that were all
made of the same
material but were
different shapes and
sizes.
Atoms were infinite in
number, always
moving and capable
of joining together.
This theory was ignored and
forgotten for more than 2000
years!
Why?

The eminent
philosophers
of the time,
Aristotle and
Plato, had a
more
respected,
(and
ultimately
wrong)
theory.
Aristotle and Plato favored the earth, fire, air
and water approach to the nature of matter.
Their ideas held sway because of their
eminence as philosophers. The atomos idea
was buried for approximately 2000 years.
Dalton’s Model

In the early 1800s,
the English
Chemist John
Dalton performed a
number of
experiments that
eventually led to
the acceptance of
the idea of atoms.
Dalton’s Theory




He deduced that all
elements are composed of
atoms. Atoms are
indivisible and
indestructible particles.
Atoms of the same element
are exactly alike.
Atoms of different elements
are different.
Compounds are formed by
the joining of atoms of two
or more elements.
.
 This
theory
became one
of the
foundations
of modern
chemistry.
Thomson’s Plum Pudding
Model
 In
1897, the
English scientist
J.J. Thomson
provided the first
hint that an atom
is made of even
smaller particles.
Thomson Model
He proposed a
model of the atom
that is sometimes
called the “Plum
Pudding” model.
 Atoms were made
from a positively
charged substance
with negatively
charged electrons
scattered about,
like raisins in a
pudding.

Thomson Model
 Thomson
studied
the passage of
an electric
current through a
gas.
 As the current
passed through
the gas, it gave
off rays of
negatively
charged
particles.
Thomson Model
 This
surprised
Thomson,
because the
atoms of the gas
were uncharged.
Where had the
negative charges
come from?
Where did
they come
from?
Thomson concluded that the
negative charges came from within
the atom.
A particle smaller than an atom had
to exist.
The atom was divisible!
Thomson called the negatively
charged “corpuscles,” today known
as electrons.
Since the gas was known to be
neutral, having no charge, he
reasoned that there must be
positively charged particles in the
atom.
But he could never find them.
Rutherford’s Gold Foil
Experiment

In 1908, the
English physicist
Ernest Rutherford
was hard at work
on an experiment
that seemed to
have little to do
with unraveling the
mysteries of the
atomic structure.
 Rutherford’s
experiment Involved
firing a stream of tiny positively
charged particles at a thin sheet of
gold foil (2000 atoms thick)
Rutherford experiment


Most of the positively
charged “bullets” passed
right through the gold
atoms in the sheet of
gold foil without changing
course at all.
Some of the positively
charged “bullets,”
however, did bounce
away from the gold sheet
as if they had hit
something solid. He
knew that positive
charges repel positive
charges.

http://chemmovies.unl.edu/ChemAnime/R
UTHERFD/RUTHERFD.html




This could only mean that the gold atoms in the
sheet were mostly open space. Atoms were not
a pudding filled with a positively charged
material.
Rutherford concluded that an atom had a small,
dense, positively charged center that repelled
his positively charged “bullets.”
He called the center of the atom the “nucleus”
The nucleus is tiny compared to the atom as a
whole.
Rutherford

Rutherford reasoned
that all of an atom’s
positively charged
particles were
contained in the
nucleus. The
negatively charged
particles were
scattered outside the
nucleus around the
atom’s edge.
Bohr Model
 In
1913, the
Danish scientist
Niels Bohr
proposed an
improvement. In
his model, he
placed each
electron in a
specific energy
level.
Bohr Model

According to
Bohr’s atomic
model, electrons
move in definite
orbits around the
nucleus, much like
planets circle the
sun. These orbits,
or energy levels,
are located at
certain distances
from the nucleus.
Wave Model
The Wave Model
Today’s atomic
model is based on
the principles of
wave mechanics.
 According to the
theory of wave
mechanics,
electrons do not
move about an
atom in a definite
path, like the
planets around the
sun.

The Wave Model


In fact, it is impossible to determine the exact
location of an electron. The probable location of
an electron is based on how much energy the
electron has.
According to the modern atomic model, at atom
has a small positively charged nucleus
surrounded by a large region in which there are
enough electrons to make an atom neutral.
Electron Cloud:




A space in which
electrons are likely to be
found.
Electrons whirl about the
nucleus billions of times
in one second
They are not moving
around in random
patterns.
Location of electrons
depends upon how much
energy the electron has.
Electron Cloud:



Depending on their energy they are locked into a
certain area in the cloud.
Electrons with the lowest energy are found in
the energy level closest to the nucleus
Electrons with the highest energy are found
in the outermost energy levels, farther from
the nucleus.
Indivisible Electron
Greek
X
Dalton
X
Nucleus
Thomson
X
Rutherford
X
X
Bohr
X
X
Wave
X
X
Orbit
Electron
Cloud
X
X
The Periodic Table
I.
II.
III.
IV.
V.
History
Arrangement of Elements
Electron Configuration
Trends
Periodic Trends
Reactivity
A. Johann Dobreiner’s Law of Triads in
1817
B. John Newlands – Law of Octaves
Lothar Meyer (1835-1895 - German)
Properties of elements show a
repetitive pattern when they are
arranged by atomic mass
D.Dimitri Mendeleev (1834-1907- Russian)
(father of modern periodic table)
Published system used today (1869)
2. Elements arranged by increasing mass
3. Left spaces for elements not yet
discovered - predicted properties
(scandium, gallium, germanium)
C.
Dimitri Mendelev
Mendeleev’s Table
His table re-organized
Mendeleev’s Periodic Table
E. Henry Mosley (1887-1915) English
1.Arrange elements
by increasing atomic
number – this led to the
periodic law
2. Periodic Law - properties
of elements are periodic
functions of their atomic #
periodic repetition of
physical and chemical
properties
II. Arrangement of Elements
A.
Periodic Table – arrangement of
elements in order of increasing
atomic number so that elements with
similar properties are in the same
column
period – horizontal row (7)
group(family)- vertical columns (1-18)
periodicity – reoccurrence of similar
properties of elements in groups
C. Special Groups on the Periodic Table
Periodic Table
D. Periodic Table Showing s,p,d,f Blocks
E. Metals – Metalloids - Nonmetals
1. Metals are on the left side – all are solids
except mercury (Hg)
a. elements near the left of a period are more
metallic than those near the right
b. elements near the top of a group are more
metallic than those near the bottom
2. Metalloids – group of elements between
metals and nonmetals(B,Si,Ge,As,Sb,Te)
3. Nonmetals are on the right side – all are
solids or gases except bromine(Br) liquid
Metals – Metalloids - Nonmetals
PROPERTY
Luster
Deformability
METAL
NONMETAL
high
malleable
and ductile
Conductivity
good
Electron gain/lose lose
Ion formed
cation (+)
Ionization energy low
Electronegativity low
low
brittle
poor
gain
anion(-)
high
high
IV.Periodic Trends(Main Group Elements)
A.
Atomic Radii
1. atomic radius is ½ the distance
between nuclei of identical atoms
joined in a molecule
2. decreases across periods (left-right)
a. caused by increasing attraction
between protons and electrons
3. increases from top to bottom
a. caused by adding electrons to
new shells
What is the atomic radius?
Atomic radii include
the region in which
electrons are found
90% of the time
Atomic Size
}
Radius
Atomic
Radius = half the distance between two
nuclei of a diatomic molecule.
Periodic Trends in Atomic Radii
Periodic Trends in Atomic Radii
Trends in Main Groups
Atomic Radii
Period Trends
A. Periodic Trends in Atomic Radii
B. Ionization Energy
1. Energy required to remove an electron
from an atom of an element (KJ/mol)
2. Increases across periods (left to right)
a. result of increased nuclear attraction
3. Decreases down groups (families)
a. electrons added to higher energy levels
b. shielding effect of inner shell electrons
c. repulsion of inner shell electrons
4. Energy to remove second and third
electron is greater
B. Trend in Ionization Energy
B. Periodic Trends in Ionization Energy
Symbol First
H
He
Li
Be
B
C
N
O
F
Ne
1312
2731
520
900
800
1086
1402
1314
1681
2080
Second
5247
7297
1757
2430
2352
2857
3391
3375
3963
Third
11
810
14840
3569
4619
4577
5301
6045
6276
C. Electronegativity
1. Measures how strongly one atom
attracts the electrons of another atom
when they form a compound
2. Increases across periods (left to right)
a. Fluorine has greatest value of 4
3. Decreases down groups
a. electrons far from the nucleus in
larger atoms have less attraction
b. Cesium and Francium with large radii
have the smallest electronegativity
Periodic Trends in Electronegativity
C. Periodic Trends in Electronegativity
C. Periodic Trends in Electronegativity
D. Ionic Radii
1. Ion – atom that acquires a charge by
gaining or losing electrons
a. cation (+) ion
anion (-) ion
2. Period trends
a. cation radii decrease across periods
b. anion radii increase across periods
3. Group trends
a. increase in cation and anion radii
down groups
Formation of an Anion (- ion)
D. Comparison of Atomic and Ionic Radii
D. Periodic Trends in Ionic Radii
E. Electron Affinity
1. Energy change that occurs when an electron is
added to a neutral atom
2. If it is easy to add an electron to an
atom the energy value is negative
a. halogens have large negative values
3. If it is difficult to add an electron to an
atom the energy value is positive
a. atoms in groups 2 and 18 have high
positive values (due to filled subshells)
b. usually higher values in larger atoms
Electron Affinity for Chlorine
Periodic Trends in Electron Affinity
Periodic Trends in Electron Affinity
PERIODIC TRENDS
Periodic Trends in Melting Point
Periodic Trends in Density
V. Reactivity
A.
Reactivity – measure of the tendency
of an element to engage in chemical
reactions by losing, gaining or
sharing electrons
1. atoms of reactive elements are very
likely to gain, lose or share electrons
2. atoms of reactive elements are likely
to form chemical bonds with other
elements
B. Reactivity and the Periodic Table
1. alkali metals (group 1) most reactive
metals
2. alkaline earth metals (group 2)
second most reactive group of metals
3. halogens (group 17) most reactive
nonmetals
4. noble gases (group 18) least reactive
C.Ionization Energy and Electronegativity
1. elements with very high and very
low values are very reactive
Electron Arrangement and Reactivity
Electron Configuration

S block [groups 1 and 2]

P block [groups 13,14,15,16,17,18]

D block [groups 3,4,5,6,7,8,9,10,11,12]

F block (lanthanide and actinide series)
H
Li
1
1s1
group 1
1s22s1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105p6
6s24f145d106p67s1
S- block
s1
s2
Alkali metals all end in s1
 Alkaline earth metals all end in s2
 Should include He but helium has
the properties of the noble gases.
- its outer shell is filled with the
maximum number of electrons
allowed for the first shell (2)

He
2
1s
2
Ne
2
2
6
1s 2s 2p 10
Ar
2
2
6
2
6
1s 2s 2p 3s 3p
18
Kr
2
2
6
2
6
2
10
6
1s 2s 2p 3s 3p 4s 3d 4p
36
1s22s22p63s23p64s23d104p65s24d105p6
Xe
54
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6
Rn
86
The P-block
p1 p2
p3
p4
p5
p6
Transition Metals -d block
d1 d2 d3
s1
d5
s1
d5 d6 d7 d8 d10 d10
F - block

inner transition elements- hold a maximum
of 14 therefore there are 14 elements in
both the actinides and lanthanides
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
Group
Ion Formed
Group 1
 Group 2
 Group 13
 Group 14
 Group 15
 Group 16
 Group 17
 Group 18
X+
X2+
X3+
Xvaries
X3X2XX0

Electron Changes
Loses 1 electron
 Loses 2 electrons
 Loses 3 electrons
 Varies
 Gains 3 electrons
 Gains 2 electrons
 Gains 1 electron
 Does not gain or
lose electrons
