2. Learning Standards
• Atomic Structure
Broad Concept:
Atomic models are
used to explain
atoms and help us
understand the
interaction of
elements and
compounds
observed on a
macroscopic scale.,
3. –Recognize discoveries from Dalton (atomic
theory), Thomson (the electron), Rutherford (the
nucleus), and Bohr (planetary model of atom) and
understand how these discoveries lead to the
modern theory.
–Describe Rutherford’s “gold foil” experiment that
led to the discovery of the nuclear atom. Identify the
major components (protons, neutrons, and electrons)
of the nuclear atom and explain how they interact.
–Write the electron configurations for the first
twenty elements of the periodic table.
4. Atomic theory proposed by
John Dalton
All matter is composed of atoms
Atoms cannot be made or destroyed
All atoms of the same element are identical
Different elements have different types of
atoms
Chemical reactions occur when atoms are
rearranged
Compounds are formed from atoms of the
7. Discovery of Protons
• 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.
8. Canal rays must be positive.
Goldstein postulated the
existence of a positive
fundamental particle called
the “proton”.
9. Thomson’s Experiment And
Discovery of Electrons
- Voltage source +
Passing an electric current makes a beam
appear to move from the negative to the
positive end.
10. Thomson’s Experiment
Voltage source
+
-
By adding an electric field he found
that the moving pieces were negative.
11. The electron was discovered in
1897 by Thomson. He imagined
the atom as a “raisin pudding” with
electrons stuck in a cake of
positive charge.
12. J.J. Thomson’s Model of Atom
• Plum Pudding Model,
1896
• Thought an atom was
like plum pudding
– Dough was cloud
– Raisins were electrons
– Didn’t know about
neutrons at this time
13. Rutherford’s experiment and
discovery of nucleus
• English physicist Ernest Rutherford
(1911)
• Shot alpha particles at fluorescent
screen.
• When an alpha particle hits a fluorescent
screen, it glows.
20. Rutherford’s Findings
Most of the particles passed right
through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing
off of tissue paper!”
Conclusions:
a) The nucleus is small
b) The nucleus is dense
c) The nucleus is positively
charged
21. The model created by Rutherford had
still some serious discordance.
According to the classic science,
electron moving around the nucleus
should emit an electromagnetic wave.
Electron should than move not by the
circle but helical and finally collide with
the nucleus. But atom is stable.
22. Rutherford also realized that the
nucleus must contain both neutral and
positively charged particles. The
neutron was then discovered in 1932
by Chadwick.
23. Atomic number and Mass number :-
• a) Atomic number (Z) :-
• The atomic number of an element is the number of protons present in the
• nucleus of the atom of the element.
• All the atoms of an element have the same atomic number.
• Eg :- Hydrogen – Atomic number = 1 (1 proton)
• Helium - Atomic number = 2 (2 protons)
• Lithium - Atomic number = 3 (3 protons)
• b) Mass number (A) :-
• The mass number of an element is the sum of the number of protons and
• neutrons (nucleons) present in the nucleus of an atom of the element.
• The mass of an atom is mainly the mass of the protons and neutrons in the nucleus of the atom.
• Eg :- Carbon – Mass number = 12 (6 protons + 6 neutrons) Mass = 12u
• Aluminium – Mass number = 27 (13 protons + 14 neutrons) Mass = 27u
• Sulphur – Mass number = 32 (16 protons + 16 neutrons) Mass = 32u
• In the notation of an atom the Mass number Symbol of 14
• atomic number and mass number element E g :- N
7
• are written as :- Atomic number
24. Isotopes
• Isotopes are atoms of the same element having the same
atomic numbers but different mass numbers.
• Eg :- Hydrogen has three isotopes. They
are1 Protium, Deuterium 3(D) and Tritium (T).
2
• 1H 1H
1
H
• Protium Deuterium Tritium
• Carbon has two isotopes. They are :-
12 14
• 6C 6C
• Chlorine has two isotopes They are :-
35 37
• 17 Cl Cl
17
•
25. Isobars
• Isobars are atoms of different elements having
different atomic numbers but same mass numbers.
• These pairs of elements have the same number of
nucleons.
• Eg :- Calcium (Ca) – atomic number - 20 and Argon
(Ar) – atomic number 18 have different atomic
numbers but have the same mass numbers – 40.
40 40
20 Ca 18 Ar
• Iron (Fe) and Nickel (Ni) have different atomic
numbers but have the same atomic mass numbers –
58.
• 58
Fe
58
Ni
27 28
•
26. Bohr’s Model of the Atom
• Similar to
Rutherford’s model
• Thought atom was
mostly empty space
• Neils Bohr, 1913
– Nucleus in center
is dense,
positively charge
– Electrons revolve
around the
nucleus.
27. Following Rutherford’s
planetary model of the atom, it
was realized that the attraction
between the electrons and the
protons should make the atom
unstable
Bohr proposed a model in
which the electrons would
stably occupy fixed orbits, as
long as these orbits had special
quantized locations
28. Parts of an Atom
Each element has a different number of protons
in its nucleus
Protons have positive charge p
Change the number of protons change
elements
This is called nuclear physics
The element also has the same number of
electrons
Electrons have negative charge
e
Change the number of electrons ionize
the element
This is called chemistry
Some elements also have neutrons
Neutrons have no charge n
They are in the nuclei of atoms
29. Subatomic particles
Actual
Name Symbol Charge mass (g)
Electron e- 9.11 x 10-28
Proton p+ 1.67 x 10-24
Neutron n0 1.67 x 10-24
30. Bohr’s model
• Electrons move around the nucleus at
stable orbits without emitting radiation.
• Electron in one of these stable orbit has
a definite energy.
• Energy is radiated only when electrons
make transitions from high energy orbit
to a low energy orbit.
31. In the Bohr model, the electron can
change orbits, accompanied by the
absorption or emission of a photon
of a specific color of light.
32. Wave Nature of Electromagnetic Radiation
Waves have 3 primary characteristics:
1. Wavelength (λ): distance between two consecutive peaks in a wave.
2. Frequency (ν): number of waves (cycles) per second that pass a given
point in space.
3. Speed: speed of light is 2.9979 * 108 m/s. We will use 3.00 x108
m/s.
Wavelength and frequency can be interconvert and they have an inverse relationship
v = c/λ
v = frequency (s1)
λ = wavelength (m)
c = speed of light (m s1)
Wavelength is also given in nm (1 nm = 10-9 m) and Angstroms (Å) (1 Å = 10-10 m).
The frequency value of s1 or 1/s is also called “hertz (Hz)” like KHz on the radio.
34. Blackbody Radiation and the Quantization
of Energy
A B
A. The interior of a cold ceramic-firing kiln approximates a blackbody, an
object that absorbs all radiation falling on it and appears black. A hot
kiln emits light characteristic of blackbody radiation. B Planck’s formula
generates a curve that fits perfectly the changes in energy and intensity
of light emitted by blackbody at different wavelength for a given
temperature
35. Planck’s Formula
• To find a physical explanation of blackbody Planck
made a radical assumption that the hot, glowing
object could emit (or absorb) only certain quantities of
energy:
• E = nhν
• Where E is the energy of the radiation, ν is its
frequency, n is a positive integer (1, 2, 3 and so on)
called a quantum number and h is a proportionality
constant now called Planck’s constant and has value
= 6,626x10-34 J.s
36. The Photoelectric Effect and The Photon
Theory of Light
• Current flow when monochromatic light of
sufficient energy shines on a metal plate
• The photoelectric effect had certain features: the
presence of a threshold frequency and the absence
of a time lag
• Carrying Planck’s idea of packeted energy,
Einstein proposed that light itself is particulate,
occurring as quanta of electromagnetic energy,
called photon
• In terms of Planck’s work we can say that each
atom changes its energy whenever it absorbs or
emits one photon, one “particle” of light, whose
energy is fixed by its frequency
• Ephoton = hν = ∆Eatom
37. Bohr's model for hydrogen atom
• Niels Bohr adopted Planck’s assumption and
explained these phenomena in this way:
1.Electrons in an atom can only occupy certain orbits
(corresponding to certain energies).
• Niels Bohr adopted Planck’s assumption and
explained these phenomena in this way:
2. Electrons in permitted orbits have specific,
“allowed” energies; these energies will not be
• Niels Bohr from the Planck’s assumption and
radiated
adopted
atom.
explained these phenomena in this way:
3.Energy is only absorbed or emitted in sucha way
as to move an electron from one “allowed” energy
state to another; the energy is defined by
E = hν
38. Bohr's model for hydrogen atom
• Lyman series
The atom will remain in the excited state for a
short time before emitting a photon and
returning to a lower stationary state. All
hydrogen atoms exist in n = 1 (invisible).
• Balmer series
When sunlight passes through the atmosphere,
hydrogen atoms in water vapor absorb the
wavelengths (visible). H atoms exist in n=2.
Similarly it will fill in:
• Paschen From n=4,5……… till n=3
• Brackutt from n=5,6……… till n=4
• Pfund from n=6,7……… till n=5
39. Bohr's model for hydrogen atom
The energy absorbed or
emitted from the process of
electron promotion or
demotion can be calculated
by the equation:
1 1
∆E = −RH ( nf2
- n2 )
i
where RH is the Rydberg
constant, 2.18 × 10−18 J, and
ni and nf are the initial and
final energy levels of the
electron.
40. Limitation of Bohr’s Model
• Bohr’s model only works for hydrogen atom and
other one-electron (hydrogen-like) ionic species,
such as He+, Li2+, etc.
• For H-atom, electronic energy:
En = -2.178 x 10-18 J(1/n2)
• For other one-electron particle: En = -2.178 x 10-18 J(Z2/n2)
– (Z = atomic number)
• Bohr’s model cannot explain atomic spectra of atoms having more
than one electron;
• Bohr’s model also cannot explain why each line in the hydrogen
spectrum appears as double-lines if the discharge tube is placed in
magnetic field.
• Perhaps his treatment of electron as having only particulate
properties is insufficient.
41. Aufbau Principle
• As protons are added
one by one to the nucleus to
build up the elements,
electrons are similarly added
to these hydrogen-like
orbital.
• H : 1s1, He : 1s2, Li :
1s2 2s1, Be : 1s2 2s2
• B : 1s2 2s2 2p1, C : 1s2
2s2 2p2.
42. Hund’s Rule
• The lowest energy
configuration for an atom is the
one having the maximum
number of unpaired electrons
allowed by the Pauli principle in
a particular set of degenerate
orbitals.
• N : 1s2 2s2 2p3, O : 1s2 2s2
2p4,
• F : 1s2 2s2 2p5, Ne : 1s2
2s2 2p6,
• Na : 1s2 2s2 2p63s1 OR
[Ne] 3s1
43. Heisenberg Uncertainty Principle
x = position h
mv = momentum ∆ x ⋅ ∆ (m v ) ≥
h = Planck’s constant 4π
The more accurately we know a particle’s position, the
less accurately we can know its momentum. Both the
position and momentum of a particle can not be
determined precisely at a given time. The uncertainty
principle implies that we cannot know the exact motion of
the electron as it moves around the nucleus.
44. Quantum Numbers (QN)
• The principal quantum number (n) is a positive integer (1, 2,
3 and so forth). It indicates the relative size of the orbital and
therefore the relative distance from the nucleus of the peak in the
radial probability distribution plot
Principal QN (n = 1, 2, 3, . . .)
• The angular momentum number (l) is an integer from 0 to n-
1. it is related to the shape of the orbital and is sometimes called
orbital-shape quantum number
Angular Momentum QN (l = 0 to n 1)
• The magnetic quantum number (ml) is an integer from –l
through 0 to +l. it prescribes the orientation of the orbital in the
space around the nucleus and is sometimes called the orbital-
orientation quantum number
Magnetic QN (ml = l to l including 0)
45. Quantum Numbers (QN)
• Spin Quantum Number, (s)This led to a fourth quantum
number, the spin quantum number, ms.
Electron Spin QN has only 2 allowed values: +1/2 and
−1/2.
46. Pauli Exclusion Principle
• No two electrons in the same atom can
have exactly the same energy.
• No two electrons in the same atom can
have identical sets of quantum numbers.
(n, l, ml, ms).
• Therefore, an orbital can hold only two
electrons, and they must have opposite
spins.
