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PHY 004: MODERN PHYSICS
LECTURE I
Drs. A. O. AKALA, K. OGUNGBEMI, &
N. E. ERUSIAFE
Atomic Structure
All matter is composed of atoms.
Understanding the structure of atoms is
critical to understanding the properties of
matter
Evolution of Atomic Theory
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.
HISTORY OF THE ATOM
1808
John Dalton
suggested that all matter was made up of
tiny spheres that were able to bounce around
with perfect elasticity and called them
ATOMS
DALTONS ATOMIC THEORY
16 X
+
8Y
8 X2Y
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.
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’s 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’s 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’s Model
• This surprised Thomson, because the
atoms of the gas were uncharged. Where
had the negative charges 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’s Model
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.
HISTORY OF THE ATOM
1898
Joseph John Thompson
found that atoms could sometimes eject a far
smaller negative particle which he called an
ELECTRON
A = alpha
B = gamma
C = beta
J.J. Thomson, measured mass/charge of e(1906 Nobel Prize in Physics)
Limitation of the Thompson’s
Model
It could not explain the scattering experiment
performed by Rutherford.
Charge of an electron
gold foil
helium nuclei
Millikan oil drop
experiment
Millikan’s Oil Drop Experiment
I.
Charge of electron – very important
application of uniform electric field
between two plate – Robert Millikan (18681953)
A.
Purpose: to find charge of an electron
1. Fine oil sprayed into air in top – gravity causes them to fall
2. A few enter the hole
3. Potential difference between plates is applied – exerts a force
on the charged drops
4. Top plate is positive enough that negative drops will rise
-
5. Potential difference adjusted to suspend (float) particle
E*q = m*g
6. Electric field was determined from potential difference between
two plates (E = V/d)
-
7. Found velocity of charge when field was turned off. Using velocity,
mg was found. Using E & mg, the charge could be calculated.
-
8. The drops had a variety of charges. So, he ionized the air, added or
removed electrons. The change in charge was always a multiple of 1.6 x 10-19 C. Thus, the charge on one electron.
-
9. Showed that charge is quantized – an object can only have charge
with a magnitude that is some integral of the charge of an electron.
-
HISTORY OF THE ATOM
1910
Ernest Rutherford
oversaw Geiger and Marsden carrying out his
famous experiment.
they fired Helium nuclei at a piece of gold foil
which was only a few atoms thick.
they found that although most of them
passed through. About 1 in 10,000 hit
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 Gold Foil
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.
Rutherford’s experiment.
Rutherford’s Gold Foil
Experiment
Rutherford’s Gold Foil
Experiment
• 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’s Model
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.
Limitations of the Rutherford’s Model
According to Rutherford’s model of an atom,
electrons revolve around the nucleus as
planets revolve around the sun. But, electrons
revolving in circular orbits will not be stable
because during revolution, they experience
acceleration. Due to acceleration, they will
lose energy in the form of radiation arid fall
into the nucleus. In such a cases the atom
would be highly unstable and would collapse.
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
orbit around the sun.
These orbits, or
energy levels, are
located at certain
distances from the
nucleus.
HISTORY OF THE ATOM
1913
Niels Bohr
studied under Rutherford at the Victoria
University in Manchester.
Bohr refined Rutherford's idea by adding
that the electrons were in orbits. Rather
like planets orbiting the sun. With each
orbit only able to contain a set number of
electrons.
Bohr’s Model of
the Atom (1913)
1.
e- can have only specific (quantized)
energy values
2.
light is emitted as e- moves from one
energy level to a lower energy level
En = -RH
1
( n2
)
n (principal quantum number) = 1,2,3,…
RH (Rydberg constant) = 2.18 x 10-18J
The Bohr Model of the Atom
The Bohr Model of the Atom:
Ground and Excited States
• In the Bohr model of hydrogen, the lowest amount
of energy hydrogen’s one electron can have
corresponds to being in the n = 1 orbit. We call this
its ground state.
• When the atom gains energy, the electron leaps to a
higher energy orbit. We call this an excited state.
• The atom is less stable in an excited state and so it
will release the extra energy to return to the ground
state.
– Either all at once or in several steps.
Line Emission Spectrum of Hydrogen Atoms
Every element has a unique emission spectrum
The Bohr Model of the Atom:
Hydrogen Spectrum
• Every hydrogen atom has identical orbits, so every
hydrogen atom can undergo the same energy
transitions.
• However, since the distances between the orbits in
an atom are not all the same, no two leaps in an
atom will have the same energy.
– The closer the orbits are in energy, the lower
the energy of the photon emitted.
– Lower energy photon = longer wavelength.
• Therefore, we get an emission spectrum that has a
lot of lines that are unique to hydrogen.
The Bohr Model of the Atom:
Hydrogen Spectrum
Bohr showed the energy a H
atom can have is equal to:
En = -RH
1
( n2
)
Ephoton = DE = Ef - Ei
1
Ef = -RH ( 2
nf
1
Ei = -RH ( 2
ni
1
DE = RH ( 2
ni
RH is the Rydberg constant
n is the principal quantum number
)
)
1
)
n2f
Light emission of sodium atom
Line spectrum
Limitations of Bohr’s Model
It is in violation of the Heisenberg Uncertainty Principle. The Bohr
Model considers electrons to have both a known radius and orbit,
which is impossible according to Heisenberg.
The Bohr Model is very limited in terms of size. Poor spectral
predictions are obtained when larger atoms are in question.
It cannot predict the relative intensities of spectral lines.
It does not explain the Zeeman Effect, when the spectral line is
split into several components in the presence of a magnetic field.
The Bohr Model does not account for the fact that accelerating
electrons do not emit electromagnetic radiation.
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.
Atomic Structure
Atoms are composed of
-protons – positively charged particles
-neutrons – neutral particles
-electrons – negatively charged particles
Protons and neutrons are located in the nucleus.
Electrons are found in orbitals surrounding the
nucleus.
HELIUM ATOM
Shell
proton
+
electron
N
N
+
-
neutron
Atomic Structure
Every different atom has a characteristic
number of protons in the nucleus.
atomic number = number of protons
Atoms with the same atomic number have
the same chemical properties and belong
to the same element.
ATOMIC STRUCTURE
Atomic number
the number of protons in an atom
Atomic mass
the number of protons and
neutrons in an atom
2
4
He
number of electrons = number of protons
ATOMIC NUMBER (Z) = number of protons in nucleus
MASS NUMBER (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
ISOTOPS are atoms of the same element (X) with different numbers of neutrons in the nucleus
Mass Number
A
Atomic Number
Z
1
1
H
235
92
2
1
U
X
Element Symbol
H (D)
238
92
3
1
U
H (T)
Atomic Structure
Atomic Structure
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.