Download Atomic Theories

Atomic Theories
The idea of an atom -- the smallest particle of matter -- has intrigued mankind
since the beginning of civilization. Throughout the centuries the "view" of the atom has
changed. New ideas, and new technologies have influenced the model of the atom. This
view of the atom is still a Theory and therefore it is still subject to change. The modern
model of the atom is called the Quantum Model and you will study this model in future.
The chart below summarizes the various atomic models that have been developed
during the course of history.
Scientist &
approximate
Date
Democritus
c.300 BC
Name of Model,
Sketch
and main idea of
theory
Atom the indivisible
particle
Atomos (in ancient
Greek) means "that
which cannot be
further broken down
into smaller pieces".
The solid sphere
model Atoms are
seen
as
solid,
indestructible spheres
(like billiard balls)
Dalton
c.1800
Importance and
Improvement on
previous model
Talks about the atom as
the smallest particle of
matter.
Defines the atom as an
indivisible particle
Explains certain natural
occurrences such as the
existence of elements
Explains a lot of
chemical
properties
such as how atoms
combine
to
form
molecules.
Explains
chemical
change better than the
Particle Theory
Confirms the basic
Laws of Chemistry:
Conservation of Mass
& definite Proportions
Shortcomings Problems
or why was it
changed
Does not give a
scientific view of
the atom only a
conceptual
definition
Does not talk
about subatomic
particles
(Electrons,
Protons,
Neutrons)
Does not include
the existence of
the nucleus
Does not explain
the existence of
ions or isotopes
Does not talk
about subatomic
particles
(Electrons,
Protons,
Neutrons)
1
The raisin bun Model
or the chocolate chip
cookie model:
Atoms are solid
spheres made-up of a
solid positive mass
(or core) with tiny
negative
particles
embedded in the
positive core.
Infers on the existence
of
electrons
and
protons.
Introduces the concept
of the nucleus.
Ifers on the relative
nuclear density and
atom mass of different
atoms
Does not explain
the existence of
electrons outside
the nucleus does
not explain the
role of electrons
in bonding.
Does not talk
about
neutrons
therefore
can't
explain
radioactivity and
the existence of
isotopes
First real modern view
of the atom.
Explains
why
the
electron spins around
the
nucleus
(Bohr's Contribution)
Proposes that the atom
Famous Gold Leaf is really mostly empty
Experiment proves space
that the nucleus is
positive and the
electrons are outside
the nucleus.
Does not place
electrons
in
definite energy
levels around the
nucleus.
Doesn't include
neutrons in the
nucleus.
Electrons in Definite
energy Levels around
the nucleus
Used atomic spectra
to
prove
that
electrons are placed
in definite orbitals
(called shells) around
the nucleus.
It
does
not
explain
the
shapes
of
molecules
or
other
abnormalities that
result
form
unevenly shared
pairs of electrons
(such as
the
abnormal
behaviour
of
water,
the
difference
in
Carbon-Carbon
Bonds between
diamond
and
graphite etc..)
J.J. Thomson
c.1850
The Planetary Model
Rutherford
c. 1905
(Neils Bohr)
BohrRutherford
c. 1920
Explains the role of
valence electrons in
bonding.
Relegates the number
of valence electrons to
the Periods of a
periodic table.
Fully explains ionic and
covalent bonding.
Places electrons in
definite energy levels
2 e- in the first
8 e- in the second
8 e- in the third
(see example below)
Does Not relate
the
valence
electrons atomic
charge
2
Modern Theory
Many Scientists
Contributed.
Some of the more
famous are:
Schroedinger
Einstein
Luis De Broglie
Max Planck
Frank Hertz
Maxwell
Fermi
Quantum Mechanical
Model
or Electron Cloud
Model
The analogy here is
that of a "beehive"
where the bees are
the electrons moving
around the nucleus in
a "cloud" of energy
levels.
Advanced
Theories
will
explain bonding
and other facts
about
the
behaviour
of
atoms and their
chemical
and
physical
properties
in
forming
new
compounds.
Other important facts about the particles of an atom:
Subatomic
Symbol
Charge
Relative Mass* Location
Particle
Proton
p+
positive
2000
nucleus
orbits
around
Electron
enegative
1
nucleus
Neutron
n0
neutral (zero)
2000
nucleus
The mass number and the atomic number from the Periodic Table are very important
numbers because they tell us how many subatomic particles are contained in a given
atom.
The atomic number tells us the number of protons., i.e. Atomic Number = Number of
Protons. The atomic mass tells the total number of particles in the nucleus, i.e. Atomic
Mass = # of protons + number of neutrons.
For example: The square where the element Boron is located on the Periodic Table
looks like this
From this we can obtain the following information:
3
Element
Symbol
Boron
B
Atomic
Number
5
Atomic #
of #
of
# of neutrons
Mass
protons electrons
11
5
5
11 - 5 = 6
Recall that the Bohr-Rutherford Model places the electrons around the nucleus in
definite orbitals or energy shells as summarized by the table below.
Energy Shell
1
2
3
4
Name of Shell Maximum
Number
or Symbol
Electrons it can contain
K
2
L
8
M
8
N
18
of
Now we use this and the
information from the
Periodic Table to draw a
Bohr-Rutherford diagram
for the Boron Atom as
illustrated below:
To draw Bohr-Rutherford
diagram for Boron we
place the first 2 electrons
in first shell.
The first shell can only
hold a maximum of 2
electrons so we start
filling the second shell
The Boron atom has 5
electrons therefore we
have another 3 electrons
to place.
We place these electrons in the Second shell and we space them apart from one another
If there are more than four electrons in the second shell (as in the case of the Fluorine
atom), we pair the electron up. This pairing of electrons is explained by more advanced
theories which propose that to counteract the repulsive forces between the electrons'
negative charges, one elctron spins in the opposite direction of the second electron.
Example:
Draw a Bohr-Rutherford diagram for the element Sodium.
Solution:
Using the periodic table we obtain the following information about the sodium atom:
Element Symbol
Sodium
Na
Atomic Atomic of
Number Mass protons
11
23
11
of
electrons
11
of
Atomic Diagram
neutrons
12
4
Electronic Configuration
Electronic Configuration is the method of listing the locations of all the electrons on a
specific atom. The system locates each electron by energy level and orbital type. The
number of electrons in each orbital type is indicated with a superscript.
Orbitals
Orbitals are probability diagrams. Specifically, an orbital describes a region in space
where there is a 90% change of finding an electron. The electron is never restricted to
an orbital as in travels around a nucleus, but it seems to keep returning to this particular
region even though its behavior is random. The concept of the orbital differs from
Bohr's concept of the orbit. Bohr considered an orbit to be a path that the electron
always followed much like a train stays on a track. The concept of the orbital was
developed in Schrodinger's work to avoid violating the Heisenberg Uncertainty
Principle. In the Modern Theory of Atomic Structures a picture of an orbital is also
called a Probability Diagram. By agreement among chemists, the orbital is a 90%
Probability Diagram. This idea allows the electron to be found anywhere and still
indicates where the electron spends most of its time.
Their names, and shapes, are based on the value of the Azimuthal Quantum Number.
l = 0, the orbital is called "s".
l = 1, the orbital is called "p".
l = 2, the orbital is called "d".
l = 3, the orbital is called "f".
Filling the electron energy levels
Main Energy Number of electrons in level in each sub-level
This is a useful pattern of remembering how the levels are filled.
http://www.webchem.net/notes/atomic_structure/electron_configuration.htm
http://www.lenntech.com/periodic-chart.htm
http://acswebcontent.acs.org/periodic/tools/PT.html
http://www.webelements.com/
5
PERIODIC PROPERTIES
Description: The properties of the elements exhibit trends and these trends can be
predicted with the help of the periodic table. They can also be explained and understood
by analyzing the electron configurations of the elements. This is because, elements tend
to gain or lose valence electrons to achieve the stable octet formation.
In addition to this activity, there are two other important trends. First, electrons are
added, one at a time, moving from left to right across a period. And, as this happens, the
electrons of the outermost shell experience increasingly strong nuclear attraction. As a
result, the electrons become closer to the nucleus and more tightly bound. The second
trend is the moving down a column in the periodic table, where the outermost electrons
become less tightly bound to the nucleus. And these trends explain the periodicity
observed in the elemental properties of atomic radius, ionization energy, electron
affinity, and electronegativity. But, before going into that we need to know a bit more
about the above mentioned terms:
Atomic Radius
The atomic radius of an element is half of the distance between the centers of two atoms
of an element that are in contact with each other. Generally, the atomic radius decreases
across a period, from left to right and increases down a given group. Therefore, the
atoms with the largest atomic radii are located in Group I and at the lower half of
groups.
.
.
Ionization Energy
Ionization energy or ionization potential is the energy required to completely remove an
electron from a gaseous atom or ion. And, the closer and more tightly an electron is
bound to the nucleus, the more difficult it is to remove and the higher its ionization
energy. Ionization energy is also required to remove a second valence electron from the
univalent ion to form the divalent ion, and so on.
A. Increase across the period (L  R) due to:
i)
ii)
More protons
Decreasing radius
stronger
attraction
B. Decrease IE down the group due to: More electrons shielding less attraction thus
lower IE
Electron Affinity
Electron affinity is the energy change that occurs when an electron is added to a gaseous
atom. It reflects the ability of an atom to accept an electron. And the atoms with
stronger effective nuclear charge have a greater electron affinity. Therefore, some
generalizations can be made about the electron affinities of certain groups in the
periodic table. The alkaline earths have low electron affinity values. This is because
they have filled sub shells. But, the halogens have high electron affinities because of the
6
addition of an electron to an atom results in a completely filled shell. Noble gases have
zero electron affinities, since each atom possesses a stable octet and will not accept an
electron readily.
Electronegativity
An atom with higher electro negativity has a great capacity for attracting bonding
electrons. Therefore, electronegativity is a measure of the attraction of an atom for the
electrons in a chemical bond.
Trends for Electronegativity:Electronegativity increases across the period. It decreases
down the group.
Reasons for trends:Increase across periods due to more protons, stronger attraction,
small radius thus large electronegativity.
Decrease down the group due to increase in the number of electrons an atom already
has, which means the atom is less able to attract another atom’s electrons, thus small
electronegativity.
periodic trends
at
form
d
n
u
po
te r
com
arac
t
h
n
c
e
l
c
l li
Cova
meta
ency
electron affinity
end
ion t
Ionization energy
electron affinity
electronegativity
atomic and ionic radii
Ionization energy
electronegativity
atomic and ionic radii
.
CHEMICAL BOND
The process of two or more atoms joining together to form a molecule is called
bonding. In general, bonding is a chemical change that occurs during chemical
reactions. During bonding electrons are involved in "fusing or gluing" two or more
atoms together.
There are three general ways in which electrons play a role in bonding.
They can either be electrostatically moved from the atom of one element to another
atom of a second element. This is known as ionic bonding
They can be shared between two different atoms or they can be shared among the same
element of the same atom. This sharing type of bonding is called molecular or covalent.
Some elements are surrounded by a cloud of free electrons which are shared among all
atoms of the same element. This mainly takes place in metallic elements and explains
the many properties of metals. This type of bonding is called metallic bonding.
7
The physical behavior of materials is thought to be determined by the structure of the
material. The phrase "Structure Determines Function/Behavior" is often used as a
reminder that the observable behavior of a material depends on its structure.
Metals
The metallic elements are found on the left-hand side of the Periodic Table. Metals are
lustrous, malleable, and ductile. They conduct heat and electricity quite well. Metals
tend to have high melting points with several notable exceptions of which mercury(Hg)
is the most well known. Metals are not soluble in water.
The proposed structure for samples of metals is that of the "band theory" of metallic
bonding. Metal samples are considered to be composed of metal atoms that are held
together by delocalized bonds formed by all of the atoms in the sample. It is proposed
that electrons can move easily within the sample. An earlier theory was sometimes
called the "electron sea model" which pictured the sharing of valence electrons by all
atoms in the sample.
Properties of Metals:
 Solid and liquid samples of metals can conduct an electric current due to the
mobility of the valence electrons of the atoms in the sample.
 Most metals will have high melting points.
 Metals are lustrous(shiny), malleable(able to flattened into sheets) and
ductile(able to be drawn into wires).

Ionic Compounds
Ionic compounds are formed when metallic elements from the left-hand side of the
Periodic Table react with nonmetallic elements from the right-hand side of the Table.
Ionic compounds have high melting and boiling points. Many ionic compounds are
soluble in water. Ionic compounds are brittle and do not conduct electricity when in the
solid form. The molten (liquid) form of ionic compounds will conduct an electric
current. In a molten sample of an ionic compound the ions are free to move. The cations
can move to the negative electrode and acquire electrons. Similarly, the anions can
move to the positive electrode and release electrons, thereby causing a flow of electricity
in the outer circuit.
The proposed structure for ionic compounds is that positive cations ( formed from the
elements on the left-hand side of the Periodic Table) attract negative anions(formed
from the elements on the right-hand side of the Periodic Table) to form a network or
lattice of oppositely charged ions arranged in three-dimensional patterns that depend on
the size and charge of the ions.
Properties of Ionic Compounds
 Solid ionic compounds do not conduct an electric current.
 Molten samples of ionic compounds can conduct an electric current due to the
mobility of the ions which are free to move to the electrodes and react.
 Ionic compounds have high melting and boiling points.
 Many ionic compounds can dissolve in water.
 Dissolved ionic compounds separate into cations and anions in solution.
 The mobile ions move to the electrodes and react to accept and release electrons
creating a flow of electricity in the outer circuit.
 Ionic compounds that are water soluble are strong electrolytes.
Covalent Compounds
Molecular compounds form when two or more nonmetal atoms form units that are
called molecules. Molecular compounds generally have low melting and boiling points.
8
Molecular compounds do not conduct electricity in the solid form or in the liquid form.
Some molecular compounds dissolve in water and some do not.
The proposed model for molecular compounds is that atoms of nonmetallic elements can
bond by sharing electrons to form units called molecules. A very common example is
the water molecule that forms when an oxygen atom shares electrons with two hydrogen
atoms to form one unit that is called a water molecule. A sample of water is thought to
be a set of these molecular units.
Molecules are attracted to other molecules by attractions that are called Intermolecular
Attractions (attractions between molecules). In a physical change such as melting or
evaporation these intermolecular attractions must be broken. Each molecule is still a
molecular unit but it is separated from the other molecular units. Intermolecular
attractions are generally weaker attractions and do not require as much energy to break
as do metallic, ionic or covalent bonds. Therefore, molecular compounds generally have
low melting points and low boiling points.
Molecules may be polar molecules which have a charge separation due to the shape of
the molecule and the polarity of the bonds between the atoms forming the molecule.
These molecules are pictured as having a positive end and a negative end and are said to
be dipoles or polar molecules. Other molecules are non-polar molecules. These are
molecules that have no charge separation again due to the shape of the molecule and the
polarity of the bonds between the atoms forming the molecule. These non-polar
molecules are pictured as having no dipole character.
LIKE DISSOLVES LIKE. Polar molecules tend to dissolve materials with charged
nature such as ionic compounds and other polar molecules. Likewise, non-polar
molecules tend to dissolve other non-polar molecules.
Properties of Molecular Compounds
 Solid molecular compounds do not conduct an electric current.
 Molten samples of molecular compounds do NOT conduct electricity.
 Molecular compounds have low melting and boiling points.
 Polar molecular compounds can dissolve in water. Some very polar molecular
compounds can be "ripped" into ions by the water molecules. These are said to
be "dissociated" into ions by the dipole nature of the water molecules.
 Polar molecular compounds are usually not soluble in non-polar solvents.
 Non-polar molecular compounds are usually not soluble in water. They are
soluble in other non-polar solvents such as toluene and other organic solvents.
 Polar molecular compounds that are dissociate into ions by water electrolytes.
http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom1s2_11.swf
http://www.chemguide.co.uk/index.html#top
http://www.ktf-split.hr/periodni/en/abc/e-config.html#7Period
9