Download Structure of the Atom

Structure of the Atom
Matter has mass and takes up space. Atoms are basic building blocks of matter, and cannot be
chemically subdivided by ordinary means.
The word atom is derived from the Greek word atom which means indivisible. The Greeks
concluded that matter could be broken down into particles to small to be seen. These particles were
called atoms
The atom is a basic unit of matter that consists of a dense central nucleus surrounded by a cloud of
negatively charged electrons. The atomic nucleus contains a mix of positively charged protons and
electrically neutral neutrons (except in the case of hydrogen-1, which is the only stable nuclide with no
neutrons). The electrons of an atom are bound to the nucleus by the electromagnetic force. Likewise, a
group of atoms can remain bound to each other, forming a molecule. An atom containing an equal
number of protons and electrons is electrically neutral, otherwise it has a positive charge if there are
fewer electrons (electron deficiency) or negative charge if there are more electrons (electron excess). A
positively or negatively charged atom is known as an ion. An atom is classified according to the number
of protons and neutrons in its nucleus: the number of protons determines the chemical element, and
the number of neutrons determines the isotope of the element.
A generic atomic planetary model, or the Rutherford model
Atoms are made from smaller subatomic particles. At the centre of an atom is a nucleus containing
protons and neutrons. Electrons are arranged around the nucleus in energy levels or shells. Make sure
you can label a simple diagram of an atom like this one. Both protons and electrons have an electrical
charge. Both have the same size of electrical charge, but the proton is positive and the electron
negative. The neutron is neutral.
Structure of Atom- Jony Mallik - M.Pharm; MS
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CHARACTERISTICS OF ELECTRON
Charge: It is a negatively charged particle.
Magnitutide of charge: Charge of electron is -1.6022 x 10-19 coulomb.
Mass of electron: Mass of electron is 0.000548597 a.m.u. or 1.1 x 10-31 kg.
Symbol of electron: Electron is represented by "e".
Location in the atom: Electrons revolve around the nucleus of atom in different circular
orbits.
CHARACTERISTICS OF PROTON
Charge: Proton is a positively charged particle.
Magnitude of charge: Charge of proton is 1.6022 x 10-19 coulomb.
Mass of proton: Mass of proton is 1.0072766 a.m.u. or 1.6726 x 10-27 kg.
Comparative mass: Proton is 1837 times heavier than an electron.
Position in atom: Protons are present in the nucleus of atom.
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CHARACTERISTICS OF NEUTRON
charge: It is a neutral particle because it has no charge.
Mass of neutron: . Mass of neutron is 1.0086654 a.m.u. or 1.6749 x 10-27 kg.
Compartive mass: Neutron is 1842 times heavier than an electron.
Location in the atom: Neutrons are present in the nucleus of an atom.
ATOMIC NUMBER
Total number of protons present in the nucleus of an atom is called
"Atomic number" or "Charge number"
Since the total number of protons and the total number of electrons in an atom are equal
therefore atomic number may also be defined as:
"Total number of electron in an atom is called Atomic number"
SYMBOL: It is denoted by "z".
Structure of Atom- Jony Mallik - M.Pharm; MS
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MASS NUMBER
Total number of protons and neutrons present in the nucleus of an atom is called
"Mass number".
SYMBOL: It is denoted by "A".
A=p+n
SYMBOL OF NUCLEUS
A nucleus is represented by the symbol :
zXA
Where X= symbol of element (Cl, Br, H etc.)
A= mass number.
Z= atomic number.
1
23
35
14
27
1H , 11Na , 17Cl , 7N , 13Al
Structure of Atom- Jony Mallik - M.Pharm; MS
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Rutherford Atomic model
The Rutherford model or planetary model is a model of the atom devised by Ernest Rutherford.
Rutherford directed the famous Geiger-Marsden experiment in 1909, which suggested, upon
Rutherford's 1911 analysis, that the so-called "plum pudding model" of J. J. Thomson of the atom was
incorrect. Rutherford's new model for the atom, based on the experimental results, contained the new
features of a relatively high central charge concentrated into a very small volume in comparison to the
rest of the atom and with this central volume also containing the bulk of the atomic mass of the atom.
This region would be named the "nucleus" of the atom in later years.
RUTHERFORD’S ATOMIC MODEL
Rutherford's atomic model shows the existence of nucleus in the atom, nature of charge
on the nucleus and the magnitude of charge on the nucleus.
APPARATUS FOR EXPERIMENT
 Alpha particles.

Gold foil. (0.0004 cm thick)

Zinc sulphide screen.

Electron Gun.
EXPERIMENT
In his experiments, Rutherford bombarded alpha particles on very thin metallic foils such
as gold foil. In order to record experimental observations, he made use of circular screen
coated with zinc sulphide.
OBSERVATIONS
He observed that most of the alpha particles were pass through the foil undeflected.
Very few particles were deflected when passed through the foil.
One particle out of 8000 particles was deflected at 90o.
Structure of Atom- Jony Mallik - M.Pharm; MS
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Few particles were deflected at different angles.
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MAIN POINTS OF RUTHERFORD’S
THEORY
Major portion of the atom is empty.
The whole mass of the atom is concentrated in the center of atom called nucleus.
The positively charged particles are present in the nucleus of atom.
The charge on the nucleus of an atom is equal to (+z.e) where Z= charge number, e =
charge of
proton.
The electrons revolve around the nucleus in different circular orbits.
Size of nucleus is very small as compare to the size of atom.
EXPLANATION OF POSTULATES
1. Since most of the alpha particles were passed through the foil undeflected, therefore, it
was concluded
that most of the atom is empty.
2. Small angles of deflection indicate that positively charged alpha particles were attracted
by electrons.
3. Large angles of deflection indicate that there is a massive positively charged body
present in the atom
and due to repulsion alpha particles were deflected at large angles.
DEFECT OF
RUTHERFORD’S THEORY
There were two fundamental defects in Rutherford's atomic model:
According to classical electromagnetic theory, being a charge particle electron when
accelerated must emit energy. We know that the motion of electron around the nucleus is
an accelerated motion, therefore, it must radiate energy. But in actual practice this does
not happen. Suppose if it happens then due to continuous loss of energy orbit of electron
must decrease continuously. Consequently electron will fall into the nucleus. But this is
against the actual situation and this shows that atom is unstable.
If the electrons emit energy continuously, they should form continuous spectrum .But
actually line
spectrum is obtained
ELECTRONEGATIVITY
"Relative tendency or relative power of an atom to attract shared
pair of electrons towards itself is called ELECTRONEGATIVITY."
Structure of Atom- Jony Mallik - M.Pharm; MS
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E.N depends upon the size of atom .
Small atoms have large values of E.N.
Big atoms have small values of E.N.
E.N decreases in a group.
E.N increases in a period.
Most Electronegatively element is "Flourine". E.N = 4
Bohr’s atomic model
Various postulates of Bohr’s atomic model are:
1. In an atom, the electrons revolve around the nucleus in certain definite circular paths called
orbits, or shells.
2. Each shell or orbit corresponds to a definite energy. Therefore, these circular orbits are also
known as energy levels or energy shells.
3. The orbits or energy levels are characterized by an integer not, where, n can have values 1, 2,
3, 4……. The integer not (= 1, 2, 3…) is called the quantum number of respective orbit. The
orbits are numbered as 1, 2, 3, 4………… etc., starting from the nucleus side. Thus, the orbit for
which n=1 is the lowest energy level.
The orbits corresponding to n = 1,2,3,4…..etc., are also designated as K,L,M,N……….etc.,
shells. When the electron is in the lowest energy level, it is said to be in the ground state.
Since, electronics can be present only in these orbits, hence, these electrons can only have
energies corresponding to these energy levels, i.e., electrons in an atom can have only certain
permissible energies .
4. The electrons present in an atom can move from a lower energy level (Elower) to a level of
higher energy (Ehigher) by absorbing the appropriate energy. Similarly, an electron can jump from
a higher energy level (Ehigher) to a lower energy level (Elower) by losing the appropriate energy.
The energy absorbed or lost is equal to the difference between the energies of the two energy
levels, i.e.,
ΔE= Ehigher - Elower
Structure of Atom- Jony Mallik - M.Pharm; MS
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AMOUNT OF ENERGY
Energy released or absorbed by an electron is equal to the difference of energy of two
energy levels.
Let an electron jumps from a higher energy level E2 to a lower energy level E1.The energy
is emitted in the form of light . Amount of energy released is given by:
E = E 2 - E1
E2 - E1= h
Where
h = Planck's consrant ( 6.6256 x 10-34 j.s)
= Frequency of radiant light
ANGULAR MOMENTUM OF
ELECTRON
Angular momentum of an electron in an energy level is given by:
m v r = nh /2
Where n =1, 2, 3, ………..
m = mass of electron
V = velocity of electron
r = radius of orbit
OR
Only those energy levels or orbits are possible for which angular momentum of electron is
an integral
multiple of h /2
.
There are particular types of Nuclide. They are:



Isotopes
Isobars
Isotones
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Here, A= Mass Number.
Z=Atomic Number
n=Number of charge (+ or-ve)
Isotopes: The atoms having same atomic number but different atomic mass number are called
Isotope.
Isobars: Nuclides having the same mass number but having the different Proton/Atomic number
are called Isobar.
Isotones: Atoms of different elements having different mass number and different atomic
number but same neutron number are called Isotones.
Hydrogen
Hydrogen atom (Z = 1) has no neutrons.
Number of protons = 1Number of electrons =
1Number of neutrons = 0It has been reported that the hydrogen element has atoms with mass
number 2 and 3 also i.e.,
Atoms of elements having the same atomic number
with different mass numbers are called isotopes.
Structure of Atom- Jony Mallik - M.Pharm; MS
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Nuclear composition of isotopes of chlorine:
Nuclear composition of isotopes of carbon:
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Characteristics of Isotopes

All isotopes of an element have the same number of valence electrons thus have identical
chemical properties.

The physical properties of the isotopes are different due to the difference in the number of
neutrons in their nuclei. The densities, melting points and boiling points etc., are slightly different.
Reason for Fractional Atomic Masses of Elements
Atomic masses of many elements are in fractions not in whole numbers.
Example:
Cl - 35.5Cu - 63.5The fractional atomic masses of elements are due to the existence of isotopes having
different masses.
Example:1
Natural chlorine consists of two isotopes:
atomic mass of chlorine.
Structure of Atom- Jony Mallik - M.Pharm; MS
Calculate the average
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Example: 2
A naturally occurring sample of Lithium contains 7.42% of 6Li and 92.58% of 7Li. The relative
mass of6Li is 6.015 and that of 7Li is 7.016. Calculate the atomic mass of a naturally occurring
sample of lithium.
Solution:
Example: 3
Which of the following two nuclei are isotopes of each other?
Solution:
The two isotopes are:
Orbit
As postulated by Bohr, an orbit is a
definite circular path at a definite
distance from the nucleus in which the
electron revolve round the nucleus.
Orbits are designated by the capital
letters K, L, M, N…. etc.
Orbits are circular in shape.
It represents the planer motion of the
electron.
An orbit indicates an exact position of
an electron in an atom.
The maximum number of electro in an
orbit is equal to 2n^2, where n is the
number of the orbit.
Structure of Atom- Jony Mallik - M.Pharm; MS
Orbital
As postulate by wave nature of an
electron, an orbital is a threedimensional region around the nucleus
within which the probability of finding
an electron is maximum., Orbital’s are
designated by s, p, d… etc.
Orbital’s have different shapes.
It represents the three-dimensional
motion of the electro round the nucleus.
An orbital does not specify the exact
position of an electron in an atom.
An orbital cannot accommodate more
than two electrons.
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Structure of Atom- Jony Mallik - M.Pharm; MS
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QUANTUM NUMBER
Different numbers of electron can be present at the different energy shell or orbit that surrounds the
central dense nucleus. From the concept of the spectral analysis of an atom, the shell can be spherical or
semispherical. A very important concept regards to the electron is that, the electron not only surrounds
the nucleus it also surrounds its own during the time of surrounding the nucleus.
Which electron present at which energy shell? The energy shell is spherical or not? What is the axis of
rotation of electron, clockwise or anti-clockwise? All the things can be described by some special
numbers called Quantum numbers.
Electrons can be labelled using the subshell and orbital or by using the four quantum numbers:




n : principal quantum number
l : azimuthal quantum number
ml : magnetic quantum number
ms : spin quantum number
Principal Quantum Number, n
The principal quantum number, n, is always a positive integer and tells us the energy level or shell that
the electron is found in.
The maximum number of subshells permitted for a particular shell is equal to n2.
The maximum number of electrons permitted in a particular shell is equal to 2 x n2.
n Energy Level
1 1st energy level
2 2nd energy level
3 3rd energy level
4 4th energy level
Shell
K
L
M
N
No. Subshells = n2
1
4
9
16
No. electrons = 2n2
2
8
18
32
Azimuthal Quantum Number, l
The azimuthal quantum number tells us which subshell the electron is found in, and therefore it tells us
the shape of the orbital.
l can have values ranging from 0 to n-1.
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The number of orbitals permitted for a particular subshell is equal to 2l + 1.
value of n
l=n-1
subshell
(orbital shape)
No. orbitals = 2l + 1
1
0
s subshell
1 (1 x s orbitals)
2
1
p subshell
3 (3 x p orbitals)
3
2
d subshell
5 (5 x d orbitals)
4
3
f subshell
7 (7 x f orbitals)
Magnetic Quantum Number, ml
The magnetic quantum number, ml, tells us the orientation of an orbital in space.
ml can have values ranging from -l to +l.
It is not always possible to associate a value of ml with a particular orbital.
value of l
subshell
values of ml
possible orbitals
0
s
0
s
1
p
-1, 0, 1
px, py, pz
2
d
-2, -1, 0, 1, 2
dxy, dxz, dyz, dx2-y2, dz2
3
f
-3, -2, -1, 0, 1, 2, 3
Spin Quantum Number, ms
The spin quantum number, ms, tells us the spin of the electron.
ms can have a value of +½ or -½.
Structure of Atom- Jony Mallik - M.Pharm; MS
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1st - Principle
QN
2nd - Orbital
QN
3rd - Magnetic
QN
n
l
ml
4th - Spin QN
ms
l goes from 0 to n1 within an energy
level
n = 1,2,3...7
l values = 0 (for s),
1(for p), 2 (for d), 3
(for f) sublevels
Values of ml go from
+l to - l , which gives
2l + 1 number of
values
has 2 values:
+1/2 (spin up) and
-1/2(spin down)
1. measures the
average distance of
the e- from the
nucleus
1. indicates the
shape of the orbital
( set of probable
locations of the e- )
1. identifies the
direction the e- orbital
has around the
nucleus
1. identifies the
"spin" or rotation
of the e- about its
own axis
2. different values of n
mean different energy
levels
2. diff. values of l
mean diff
sublevels. In a
sublevel all the ehave nearly the
same energy.
2. specifies the eorbital in which the eis located within a
sublevel.
2. shows that each
orbital can contain
only 2 e-
3. different values of n
mean relatively large
differences in the
energies of the e-s
3. different
sublevels within
the same level may
have moderately
large differences in
energy.
3. different values of
ml mean little
difference in energies
of the e-
3. the direction of
spin is either in
one direction or the
other
4. the smallest
avgerage distance and
the lowest energy
occurs when n = 1;
each increase in n
increases those
quantities.
4. within any level,
the lowest energy
sublevel is s, then
p, then d, then f.
4. the number of
possible values of ml
within a sublevel
idenities how many epairs that the
sublevel can hold
4. when 2 e- (in an
atom) have the
same set of QN
except for ms, then
these e- are called
an e- pair
5. the number of epossible in a level is
2n2
5. the number of
possible values of l
for a level is equal
to the value of n
Structure of Atom- Jony Mallik - M.Pharm; MS
5.these e- within
an e- pair have
essentially the
same energy
Page 15
Electronic Configuration of elements
Occupation of Orbitals
The first thing to keep in mind is that electrons fill orbitals in a way to minimize the energy of
the atom. This would mean that the electrons in an atom would fill the principal energy levels in
order of increasing energy (the electrons are getting farther from the nucleus). The order of levels
filled would look like this:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p
One way to remember this pattern, probably the easiest, is to refer to the periodic table and
remember where each orbital block falls to logically deduce this pattern. Another way is to make
a table like the one below and use vertical lines to determine which subshells correspond to each
other.
Structure of Atom- Jony Mallik - M.Pharm; MS
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Pauli Exclusion Principle
The second major fact to keep in mind is the Pauli Exclusion Principle which states that no two
electrons can have the same four quantum numbers. The first three (n,l, and ml) may be similar
but the fourth quantum number must be different. We are aware that in one orbital a maximum of
two electrons can be found and the two electrons must have opposing spins. That means one
would spin up ( +1/2) and the other would spin down (-1/2). This tells us that each subshell has
double the electrons per orbital. The s subshell has 1 orbital that can hold to 2 electrons, the p
subsheel has 3 orbitals that can hold up to 6 electrons, the d subshell has 5 oribtals that hold up to
10 electrons, and the f subshell has 7 oribtals with 14 electrons.
Example
Structure of Atom- Jony Mallik - M.Pharm; MS
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We have the first three quantum numbers n=1, l=0, ml=0. Only two electrons can correspond to
these, which would be either ms = -1/2 or ms = +1/2. As we already know from our studies of
quantum numbers and electron orbitals, we can conclude that these four quantum numbers refer
to 1s subshell. If only one of the ms values are given then we would have 1s1 (denoting
Hydrogen) if both are given we would have 1s2 (denoting Helium). Visually this would be
represented as:
As you can see, the 1s subshell can hold only two electrons and when filled the electrons have
opposite spins.
Hund's Rule
When assigning electrons in orbitals, each electron will first fill all the orbitals with similar
energy (also referred to as degenerate) before pairing with another electron in a half-filled
orbital. Atoms at ground states tend to have as many unpaired electrons as possible. When
visualizing this processes, think about how electrons are exhibiting the same behavior as the
same poles on a magnet would if they came into contact; as the negatively charged electrons fill
orbitals they first try to get as far as possible from each other before having to pair up.
Example
Structure of Atom- Jony Mallik - M.Pharm; MS
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If we look at the correct electron configuration of Nitrogen (Z = 7), a very important element in
the biology of plants: 1s2 2s2 2p3
We can clearly see that p orbitals are half filled as there are three electrons and three p orbitals.
This is because Hund's Rule states that the three electrons in the 2p subshell will fill all the
empty orbitals first before filling orbitals with electrons in them. If we look at the element after
Nitrogen in the same period, Oxygen (Z = 8) its electron configuration is: 1s2 2s2 2p4
Oxygen has one more electron than Nitrogen and as the orbitals are all half filled the electron
must pair up.
The Aufbau Principle
Aufbau comes from the German word "Aufbauen" which means "to build". When writing
electron configurations, we are building up electron orbitals as we proceed from atom to atom.
As we write the electron configuration for an atom, we will fill the orbitals in order of increasing
atomic number. However, there are some exceptions to this rule.
Structure of Atom- Jony Mallik - M.Pharm; MS
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Example
If we follow the pattern across a period from B (Z=5) to Ne (Z=10) the number of electrons
increase and the subshells are filled. Here we are focusing on the p subshell in which as we move
towards Ne, the p subshell becomes filled.
B (Z=5) configuration: 1s2 2s2 2p1
C (Z=6) configuration:1s2 2s2 2p2
N (Z=7) configuration:1s2 2s2 2p3
O (Z=8) configuration:1s2 2s2 2p4
F (Z=9) configuration:1s2 2s2 2p5
Ne (Z=10) configuration:1s2 2s2 2p6
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DALTON ‘S ATOMIC THEORY
Main postulates of Dalton atomic theory are as follows:
1. Matter is composed of very tiny or microscopic particles called "Atom".
2. Atom is an indivisible particle.
3. Atom can neither be created nor it is destroyed.
4. Atoms of an element are identical in size, shape, mass and in other properties.
5. Atoms of different elements are different in their properties.
6. Atoms combine with each other in small whole numbers.
7. All chemical reactions are due to combination or separation of atoms.
DEFECTS IN DALTON’S THEORY:
Postulate number 2, 3, 4 and 6 are not correct as described below:
DEFECT NO: 1
Atom can be divided into a number of sub-atomic particles such as electron, proton
and neutron etc.
DEFECT NO: 2
Atoms of an element may be different in their masses.
For example:
1H
1
, 1H2, 1H3
17Cl
35
,
17Cl
37
DEFECT NO: 3
All compounds do not have small number of atoms.
For example:


Decane C10H22.
Sugar C12H22O11.
DEFECT NO: 4
Atom can be destroyed by fission process in
Atom bomb.
Nuclear reactor.
On the basis of above defects, Dalton's atomic theory has failed now.
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Structure of Atom- Jony Mallik - M.Pharm; MS
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