Atomic structure

1. ATOMIC STRUCTURE
-: made easy:-
by
M.,Anwar Sohail
Bachelor of Science and Education
Master of Science (Organic Chemistry)
CHEMISTRY EDUCATOR
Pelham High
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2. Democritus’ atom
[Hypothetical / Not based on
experiments]
Democritus proposed
that matter is composed
of tiny indivisible
particles called ‘atom’
The word ‘atom’ means
unable to be divided.
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3. Dalton’s atomic theory
(1808)
[Based on experiments]
 Every element is made of tiny,
unique particles called atoms that
cannot be subdivided.
 Atoms of the same element are
exactly alike.
 Atoms of different elements can
join to form molecules.
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4. Discovery of
fundamental or
subatomic particles
The electrons, protons and neutrons
are called fundamental particles or
fundamental subatomic particles.
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5. Canal Rays and Protons
(1886)
 Eugene Goldstein noted streams of positively charged
particles in cathode rays in 1886.
– Particles move in opposite direction of cathode rays.
– Called “Canal Rays” because they passed through holes
(channels or canals) drilled through the negative
electrode.
 Canal rays must be positive.
– Goldstein postulated the existence of a positive
fundamental particle called the “proton”.
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6. Discovery of Electrons (1897)
Electrons are discovered
by J.J. Thompson when high
voltage is applied across a sealed
glass tube called the ‘discharge
tube’ or CRT at very low pressure. The Discharge
He found that what was called as Tube
cathode rays until his time was not
“rays” but “particles” travelling
from cathode to anode.
He called them electrons.
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7. Discovery of
Neutrons (1932)
 James Chadwick in 1932 analyzed
the results of α-particle scattering
on thin Be films.
 Chadwick recognized existence of
massive neutral particles which he
called neutrons.
– Chadwick discovered the neutron.
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8. Characteristics of
subatomic particles at a
glance
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9. Thomson’s Atomic
model(1898)
(Also called Plum-pudding
model)
Thomson puts together the
subatomic particles and comes
forward with his atomic model.
In the atom the mass and the
positive charge is evenly
distributed throughout the atom
(like pudding) and the negatively
charged electrons are embedded
in it like the plum.
He could not experimentally
prove his model.
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10. Rutherford’s Alpha
particles scattering
experiment (1911)
 Rutherford
bombarded Alpha
particles on a very
thin(0.00006cm) gold
foil.
 Most of the particles
passed through, some
deflected at large
angles and 1 in 20000
deflected back to its
own path.
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11. Inference from Rutherford’s
experiment
 Almost all the Alpha
particles passed
through the gold foil
means Most of the
atom is empty space.
 Some of the + charged
alpha particles are
deflected at large
angles because there is
a very tiny dense core
of mass and + charge
located in the atom.
(Called nucleus)
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12. Rutherford’s Alpha Rays Scattering Experiment
Most of the Alpha Particles passed Through.
Alpha
One
Particles in every 20,000 deflected back on its own path.
Source
Gold Foil
Some Deflected at large angles.
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13. Rutherford’s Atomic model
(1920)
Based on his experiment he postulated a
model. The important postulates are:
 The atom is mostly hollow.
 The mass and the positive charge
(protons & neutrons) are located
at the center at a very small
portion called nucleus.
 The electrons revolve around the Also called Planetary
nucleus like the planets revolve model
around the sun.
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14. Atomic (z)# and Mass #
 Atomic # is the # of protons present
inside the nucleus of an atom.
 It’s unique to each element therefore,
its the identity of an element.
 No two elements can have the same
atomic number.
 Elements are listed in the periodic
table in the increasing order of their
atomic numbers.
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15. Mass # (A)
 Mass # is the sum of the # of protons
and neutrons present inside the
nucleus of an atom. Therefore, must
be a whole # not decimal.
 A = p + n OR A = Z + n
 The periodic table lists the average
atomic mass not the mass #.
 Atomic mass rounded to nearest
whole number is the Mass #.
 Mass # has no units.
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16. Isotopes
 Atoms of the same element with
different mass #.
 They have
1. Same atomic #
2. Same symbol
3. Same # of protons and electrons
4. Different # of neutrons & mass #
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17. Isotopic Symbol
Net Charge
Mass # (A)
Atomic # (Z)
Symbol
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18. The 3 Nuclie of H isotopes
P P P
N N N
Z Z Z
M M M
e e e
Complete Isotopic Symbol Worksheet
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19. Average Atomic Mass
 Weighted average of the atomic
masses of all the naturally occurring
isotopes of an element is called
average atomic mass.
 It is measured in amu (atomic mass
unit)
 1 amu is 1/12th of the mass of C-12
atom.
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20. Calculating Average
Atomic Mass
(Mass of A*%) + (Mass of B*%)
100
Complete Average Atomic Mass – 1 Worksheet
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21. Isotones
 Isotones are atoms of different
elements with same # of neutrons.
 Examples:
 S – 32 and P – 31
 Ca – 40 and K – 39
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22. Problems with Rutherford’s
model
 As per the classical laws of Physics: if a particle
(electron) is revolving around oppositely charged
particle (positive nucleus), the revolving particle
loses its energy continuously and finally falls in
to the central particle. Therefore the atom should
collapse.But this is not happening in nature.
 If the negatively charged electron is revolving
around positively charged nucleus, the atomic
spectra should be a band spectrum but in nature
the atomic spectrum is line spectrum.
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23. Band spectrum
 When white light is passed through a
prism, it splits in to 7 different colors and
they appear as bands of 7 colors on a film
or screen. (Example in nature: Rainbow)
 This is called a band spectrum. It is not a
characteristic of an atomic spectrum
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24. Line Spectrum or Atomic
emission spectrum
 When electricity is passed through a tube filled
with a gas (Ex.CRT), light will be emitted. If the
emitted light is passed through the prism and its
image is recorded on a film it appears as ‘sharp
lines on black background’. This is called line
spectrum or “atomic emission spectrum” .Every
element has a characteristic emission spectrum
of its own.
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25. Absorption
spectrum
 An absorption spectrum is formed by
shining a beam of white light through a sample
of gas.
– Absorption spectra indicate the wavelengths of
light that have been absorbed by the gas.
– It appears as dark lines on bright background.
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26. Characteristics of Light
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27. Characteristics of Light
 Velocity (c): Distance traveled by light in 1 second.
It’s a constant 3.00 x 10 8 m/s
c=ν λ
 Wave length (λ): Distance between any two similar
points on successive waves. Measured in m or nm
(nano meters) 1nm = 10 – 9 m = 10 – 7 cm
λ =c/ν
 Frequency (ν ): # of waves that cross a given point
in 1 second. Measured in Hertz (Hz) or cycles per
Hz
second (cps) ν = c / λ
cps
 Amplitude: Height of a crest or depth of a trough.
Refers to the intensity of light.
Energy (E): Energy contained in a wave. Measured
in Joules (J) E = h ν
Where h is Planck’s constant (6.626 x 10-34 J)
08/19/12 Complete Characteristics of light Worksheet
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28. Frequency, Wavelength & Energy
relationships
When frequency increases:
 Energy increases
 Wavelength decreases
When Wavelength increases:
 Energy decreases
 Frequency decreases
When amplitude decreases:
 intensity (brightness of light) decreases
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29. Bohr’s atomic model
Neils Bohr presented his atomic
model retaining the basic idea
of Rutherford’s model. The
important postulates are:
1. Electrons revolve around the
nucleus in definite, closed,
circular paths called orbits.
2. Each orbit is associated with a
definite amount of energy
therefore also called as
energy level.
3. These orbits or energy levels
are numbered 1,2,3,4….. or
K,L,M,N…. from inside
onwards. Bigger the orbit, They are also called principal
greater is the energy quantum levels, represented by
associated with it. ‘n’.
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30. Bohr’s model
Continued:-
4. More than one energy
levels are possible for an
electron. However, as long
as an electron is in an
energy level its energy
remains constant.
5. When an electron gains
energy it jumps from lower
energy level to higher.
6. When it jumps back from
higher energy level to
lower, it loses energy in 7. The energy released ( ∆E )can be
the form of light. calculated by:
Where ‘h’ is Planck’s constant, ‘v’ is the frequency of light emitted.
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31. Modern Model
or
Wave Mechanical Model
or
Quantum Mechanical Model
of the atom
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32. Particle nature of light(1901)
Max Planck's Quantum Theory
 Max Planck studied the radiation emitted by various
objects at high temperatures and came to a conclusion that:
 Light is absorbed or emitted by matter in the form of
discrete packets of energy. Each energy packet is called
Photon and the energy it holds is termed Quantum.
 The energy contained in each PHOTON of light is directly
proportional to its frequency and can be calculated by the
equation: E=hv
 Where h is Planck’s constant (6.626 x 10-34 J)
 Planck’s quantum theory helped understanding the
phenomenon of Photoelectric effect (ejection of electrons
from the surface of metal when light of a certain frequency
08/19/12 falls on it.)
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33. DeBroglie’s
Dual nature of electron
 Based on Planck’s quantum
theory and Bohr’s quantized
orbits, De Broglie suggested
that:
 every moving particle exhibits
a wave nature so also the
electrons.
 electrons behave more like
waves on a vibrating string
than like particles.
 The wave length of any
particle wave can be
calculated by the equation: λ =h/mv (De Broglie’s equation)
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34. Heisenberg’s uncertainty
principle
 Its impossible to find out both the position and the
speed of an electron accurately at the same time.
 It is because to locate an electron, light, having
wave length shorter than the size of an electron
should fall on it and be reflected. Light with such
a short wave length will have very high energy,
which will energize the electron. Therefore, its
velocity around the nucleus will increase.
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35. Splitting of Bohr’s Spectral lines
Introduction to
Orbitals/Quantum #s
With a regular spectroscope
Principal Quantum # =(n)
sharp principal diffused fundamental
Under High resolution
spectroscope
Azimathul Quantum # =(l)
Under Magnetic Field
Magnetic Quantum # =(m)
Under Electric Field
Spin Quantum # =(s)
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36. Principle Quantum #
 Bohr’s spectral lines.
 = Energy Level or Bohr’s atomic orbits.
 Values: any # 1,2,3,….so on from inside
outwards.
 Value can’t be zero
 The total # of electrons that can be
accommodated in an energy level is given
by 2n2 where n is the Principle Quantum #
or energy level.
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37. Azimathul Quantum # l
 Splitting of Bohr’s spectral lines under high resolution
spectroscope.
 = sub energy Levels or orbitals.
 Values: 0 to n – 1
 Total # of “l ” values will be equal to n
 Ex: For n=4 l values will be: 0,1,2,3
 Figure out the l values for 1st, 2nd & 3rd energy levels
 l = 0 : s orbital
 l = 1 : p orbital
 l = 2 : d orbital
 l = 3 : p orbital
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38. Magnetic Quantum # (m)
 Splitting of high resolution lines in
magnetic field.
 Also called angular momentum Q #.
 # of m values for each l value = 2l +1
(How many?)
 Value ranges from – l to 0 to + l
(What are they?)
Practice:
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39. Figure out the m values
1. How many m values are there for s orbital?
2. What are they?
3. How many m values are there for p orbital?
4. What are they?
5. How many m values are there for d orbital?
6. What are they?
7. How many m values are there for f orbital?
8. What are they?
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40. Spin Quantum # (s)
 Indicates electron spin in the orbital or
electron cloud (either clockwise or
counterclockwise)
 Values: for each m value there are 2 s
values; they are +1/2 and -1/2
 This indicates that in each m there are
2 electrons one spinning clockwise
and the other counterclockwise.
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41. Atomic Orbitals
 As it is impossible to locate an electron’s exact
position at a given time, therefore:
 The area around the nucleus where the probability
of finding an electron is maximum is called an
orbital.
 There are 4 atomic orbitals discovered so far. They
are s,p,d,f
 The s orbital is spherical shaped electron cloud, the
p orbital is a dumbbell shaped electron cloud and
the d orbital is a double dumbbell and the f orbital
is an 8 lobbed dumbbell.
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42. Orbitals and Electrons
Energy level Types of # of Orbitals Electrons Total # of electrons in the
Principal Quantum Orbitals Magnetic Quantum Spin Quantum # (s) energy level
# (n) Azimuthal Quantum # (m)
# (l)
1 s 1 2 2
2 s 1 2 8
p 3 6
3 s 1 2 18
p 3 6
d 5 10
4 s 1 2 32
p 3 6
d 5 10
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43. s Orbitals
There is one s Orbitals in each energy level
Each one can hold 2 electrons.
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44. p Orbitals
There are 3 p Orbitals in each energy level
(from 2nd energy level on wards)
Each one can hold 2 electrons therefore 6 electrons in each p sublevel
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45. d Orbitals
There are 5 p Orbitals in each energy level
(from 3rd energy level on wards)
Each one can hold 2 electrons therefore 10 electrons in each d sublevel
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46. What does the modern
atom look like?
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47. Electron Configuration
 Arrangement of electrons in the various orbitals
of the atom of an element is called electron
configuration.
 It is governed by 3 laws:
 Aufbau Principle: Electrons occupy the lowest
energy orbital available.
 Pauli’s exclusion principle: No more than two
electrons in each orbital.
 Hunds rule: When degenerate orbitals are
available, Pairing of electrons takes place after
half filling.
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48. Atomic Orbital Energy
Diagram
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49. End of
Pr esentation
Remember the Atomic
Structure.
It is the key to learn Chemistry
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