Download Atomic and Nuclear Physics

ATOMIC AND NUCLEAR
PHYSICS
Topic 7.1 The Atom
Courtesy:
https://sites.google.com/site/ndhssciencerevisionstorage/ib-physics-summarypowerponts
Physics for the IB Diploma by Tsokos
ATOMIC STRUCTURE_ DALTON’S MODEL
John Dalton said that atoms were tiny
indivisible spheres, but in 1897 J. J.
Thomson discovered that all matter
contains tiny negatively-charged particles.
 He showed that these particles are smaller
than an atom.
 He had found the first subatomic
particle - the electron.

ATOMIC STRUCTURE_ THOMSON’S MODEL
Scientists then set
out to find the
structure of the
atom.
 Thomson thought
that the atom was a
positive sphere of
matter and the
negative electrons
were embedded in
it as shown here
 This `model' was
called the
`plum-pudding'
model of the atom.

RUTHERFORD’S EXPERIMENT
Ernst Rutherford decided to probe the atom
using fast moving alpha (α) particles.
 He got his students Geiger and Marsden to fire
the positively-charged α-particles at very thin
gold foil and observe how they were scattered.
 Simulation

RUTHERFORD’S EXPERIMENT_ FINDINGS


Most of the α-particles passed straight through
the foil, but to his surprise a few were
scattered back towards the source at very
large scattering angles.
Rutherford said
that this was
rather like firing
a gun at tissue
paper and
finding that
some bullets
bounce back
towards you!
INTERPRETING THE RESULTS

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The very large deflection was indicative of a
large repulsive force between the alpha
positive particles and the carrier of the
positive charge in the atom.
Such force could not be produced if the
positive charge was distributed over the entire
atomic volume (Thomson’s model).
Rutherford concluded that the positive charge
in the atom must be very tiny so that the
alpha particles could come very close, and
massive; otherwise, the atom would have
recoiled.
Thus, the alpha particles can approach the
positive charge at very small distance and the
electric repulsive force, being proportional to
the inverse of the square of the distance,
would be very large This explains the large
deflection.
Most α-particles are hardly deflected because
they are far away from the nucleus and the field
is too weak to repel them much.
 The electrons do not deflect the α-particles
because the effect of their negative charge is
spread thinly throughout the atom.

RUTHERFORD’S MODEL OF THE ATOM_NUCLEAR
MODEL
Rutherford nuclear
model of the atom:
 All of an atom's positive
charge and most of its
mass is concentrated in
a tiny core.



The electrons orbit the
nucleus, like planets
orbit the sun


Rutherford called this the
nucleus.
Planetary model
The atom is mainly
empty space!



Using this model Rutherford
calculated that the diameter
of the gold nucleus could
not be larger than 10-14 m.
This diagram is not to scale.
With a 1 mm diameter
nucleus the diameter of the
atom would have to be
10000 mm or 10 m!
The nucleus is like a pea at
the centre of a football pitch.
DIFFICULTIES WITH THE RUTHERFORD’S MODEL

Atoms emit certain discrete characteristic
frequencies of electromagnetic radiation


The Rutherford model is unable to explain this phenomena
According to Maxwell’s theory of electromagnetism,
an accelerated charge should radiate
electromagnetic waves and thus lose energy.
The electrons in Rutherford’s model move in circular orbits
and are subject to centripetal acceleration.
 Accordingly, they must radiate and lose energy, and then
spiral into the nucleus.

 This
is not the case!
BOHR’S MODEL OF THE ATOM
Bohr was the first to address the problems
with the Rutherford Model.
 He benefitted from earlier studies done on the
emission spectrum of gases.

CONTINUOS SPECTRUM OF WHITE LIGHT
The light coming from the sun, from a candle,
or from an incandescent lamp is called white
light because it contains all the wavelengths of
the visible light.
 If you pass this white light through a prism or a
through a spectroscope, then the white light is
broken down into all the visible colors. We
obtain a continuous spectrum

LIGHT SPECTRA
Not all spectra of light are continuous.
 If you pass the light coming from a Fluorescent
lamp through a prism, the obtained spectrum is
not continuous.
 There are three types of spectra:

 Continuous
spectrum
 Emission spectrum
 Absorption spectrum
EMISSION SPECTRA OF GASES



Thomas Melville was the first to study the light
emitted by various gases.
He used a flame as a heat source, and passed
the light emitted by the heated gas through a
prism.
Melville discovered that the pattern produced by
light from heated gases is very different from the
continuous rainbow pattern produced when
sunlight passes through a prism.
EMISSION SPECTRA OR LINE SPECTRA



The new type of spectrum consisted of a series of bright
lines separated by dark gaps. This spectrum became known
as a line spectrum.
Melville also noted the line spectrum produced by a
particular gas was always the same.
The line spectra are produced from atoms of gases that have
been excited in some way, either by heating or by an
electrical discharge.
ABSORPTION SPECTRA

Suppose that you pass white light (containing
all visible light wavelengths) through a cool gas
element. If you analyze the light that passes
through the cool gas, you will notice that the
spectrum of the light is not continuous. It is
missing some wavelengths. This spectrum is
called absorption spectrum because the cool
gas absorbed these wavelengths.
EMISSION VERSUS ABSORPTION SPECTRA



The dark lines on an absorption spectrum will fall in exactly the
same position as the bright lines on an emission spectrum for a
given element.
For example, the emission spectrum of sodium shows a pair of
characteristic bright lines in the yellow region of the visible
spectrum.
An absorption spectrum will show 2 dark lines in the same
position.
Emission Spectra
Absorption Spectra
BOHR’S MODEL OF HYDROGEN



In 1913 Bohr provided an explanation of atomic
spectra that includes some features of the currently
accepted theory
His model includes both classical and non-classical
ideas
His model included an attempt to explain why the
atom was stable
THE BOHR THEORY OF HYDROGEN
In 1913 Bohr provided an explanation of atomic
spectra that includes some features of the
currently accepted theory
 His model includes both classical and nonclassical ideas
 His model included an attempt to explain why
the atom was stable

BOHR’S ASSUMPTIONS FOR HYDROGEN

The electron moves in circular
orbits around the proton under
the influence of the Coulomb
force of attraction


The Coulomb force produces the
centripetal acceleration
Only certain electron orbits are
stable


These are the orbits in which the
atom does not emit energy in the
form of electromagnetic radiation
Therefore, the energy of the atom
remains constant and classical
mechanics can be used to
describe the electron’s motion
BOHR’S MODEL OF HYDROGEN

The hydrogen single
electron can only exist in
specific orbits of specific
energy (Potential energy).
These orbits are called
energy levels.
 The
electron energy is
discrete (quantum) and not
continuous
 The electron has its lowest
energy state when it is in
the ground level (least PE).
 When the electron is in an
energy level other than the
ground level, it is said to be
in an excited state of
energy
BOHR’S MODEL OF HYDROGEN




The electron can lose energy only
when it jumps from one energy
state level to a lower one.
The emitted energy is equal to the
difference in energy between the
initial and final states.
Similarly, an electron can jump
from one energy state level to a
higher one only if it receives an
amount of energy equal to the
difference in energy between the
final and initial states.
In this sense, Bohr’s model
provided an explanation for the
emission and absorption spectra
NUCLEAR STRUCTURE
MASS NUMBER

The total number of protons and neutrons in
the nucleus is called the mass number (or
nucleon number).
NUCLEON
Protons and neutrons are called nucleons.
 Each is about 1800 times more massive than
an electron, so virtually all of an atom's mass is
in its nucleus.

ATOMIC NUMBER

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All materials are made from about 100 basic
substances called elements.
An atom is the smallest `piece' of an element you can
have.
Each element has a different number of protons in its
atoms:
it has a different atomic number (sometimes called
the proton number).
The atomic number also tells you the number of
electrons in the atom.
ISOTOPES
Every atom of oxygen has a proton number of
8. That is, it has 8 protons (and so 8 electrons
to make it a neutral atom).
 Most oxygen atoms have a nucleon number of
16.
 This means that these atoms also have 8
neutrons.
 This is 168O.

Some oxygen atoms have a nucleon number of
17.
 These atoms have 9 neutrons (but still 8
protons).
 This is 178O.
 168O and 178O are both oxygen atoms.
 They are called isotopes of oxygen.

There is a third isotope of oxygen 188O.
 How many neutrons are there in the nucleus of
an 188O atom?


Isotopes are atoms with the same proton
number, but different nucleon numbers.
Since the isotopes of an element have the
same number, of electrons, they must have the
same chemical properties.
 The atoms have different masses, however, and
so their physical properties are different.

EVIDENCE FOR NEUTRONS

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The existence of isotopes is evidence for the existence of
neutrons because there is no other way to explain the
mass difference of two isotopes of the same element.
By definition, two isotopes of the same element must
have the same number of protons, which means the
mass attributed to those protons must be the same.
Therefore, there must be some other particle that
accounts for the difference in mass, and that particle is
the neutron.
INTERACTIONS IN THE NUCLEUS

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Electrons are held in orbit by the force of
attraction between opposite charges.
Protons and neutrons (nucleons) are bound
tightly together in the nucleus by a different kind
of force, called the strong, short-range nuclear
force.
There are also Coulomb interaction between
protons.
Due to the fact that they are charged particles.