Atomic spectra&models

CLIL Physics/English Daniela Aprile Liceo Linguistico “G. Verga” ModicaCLIL Physics/English Daniela Aprile Liceo Linguistico “G. Verga” Modica
Atomic Spectra and
Models of the Atom
LEARNING OBJECTIVE
• To know the relationship between atomic
spectra and the electronic structure of
atoms.
• To understand how scientific models
develop.

Atomic Spectra and Models
of the Atom

Do you know…
• … what is an atom?
• … what are atoms made of?
• … how/when were the atom components
discovered?
• …what is the light spectrum?

FIRST DISCOVERIES
• 1897 Joseph John Thomson (1856-1940)
discovered the electron, experimentally
determining its charge to mass ratio
e/m =1.7591011 C/kg and concluded that they
are a component of every atom.
• 1909 Robert Millikan (1868-1953) - Oil drop
experiment - determined the charge and the mass
of an electron.
e =1.602 10-19 C m = 9.11 10-31 kg

CLIL Physics/English Daniela Aprile Liceo Linguistico “G. Verga” Modica
HISTORY OF THE ATOM
TIMELINE

ATOMIC MODELS
Thomson (1902)
Rutherford (1911)
Bohr-Sommerfeld (1913)
Shrödinger (1926)

THOMSON’S MODEL (1902)
Plum Pudding Model
Electrons are surrounded by a
soup of positive charge to
balance the electrons negative
charges,
like negatively charged "plums"
surrounded by positively
charged "pudding"

The Gold Foil Experiment
Rutherford, Geiger & Marsden (1911):
“ It was almost as incredible as if you fired
a 15-inch shell at a piece of tissue paper
and it came back and hit you.”

RUTHERFORD MODEL (1911)
An atom contains a
small dense positive
charge surrounded by
negatively charged
electrons revolving
around the positive
charge just like planets
orbit the sun.

But…
An accelerating charged particle emits
electromagnetic radiation
The orbiting electron has centripetal acceleration
It looses energy
The orbit radius decreases
The electron falls onto the nucleus in about 10-11s
The atom is not stable

What do you know so far?
• Who discovered the electron?
• What is the plum-pudding model?
• Why was the Rutherford’s “gold foil”
experiment outcome surprising?
• What is new in Rutherford’s model?
• Why wasn’t it a suitable model?

FROM CLASSICAL TO QUANTUM
PHYSICS
When you try to apply the classical mechanics
(Newton’s) laws to microscopic objects –
such as atoms and the e.m. waves they emit –
you come to contradictions with experience.
Therefore the fundamental dynamics’ and
electromagnetism’ laws have to be changed
so that the new laws correctly describe
microscopic phenomena, but reduce to
classical laws when applied to macroscopic
objects.
“Correspondence Principle” – Niels Bohr

BOHR MODEL
The electrons move about
the nucleus in "stationary
states" which are stable,
NOT radiating energy.
Each orbit can hold no
more than 2 electrons.
(Pauli’s Exclusion
Principle, 1925)

ELECTRONIC TRANSITIONS
An electron is able to
absorb or emit energy
in order to “jump” from one to the other
orbit.
But what is it exchanging energy with?

The photon (1905)
Some metals emit electrons
when light shines on them:
the brighter the light,
the more electrons.
Einstein (1905) explained the
effect by describing light as a
bundle of particles, or
photons.
A photon gives a quantum of
energy depending on its
frequency/wavelenght
(colour)
E = hf = hc/.
Jsh 34
1063.6 


COMPTON EFFECT (1923)
If a photon has high energy (X-rays), the electron is given
enough energy to be completely ejected from its atom,
and a photon containing the remaining energy (smaller
frequency) is scattered in a different direction from the
original.
LIGHT = PARTICLE (PHOTON)

ABSORPTION
An atom can absorb a photon
of light, so that an electron
“jumps” to an orbit that has a
higher energy.
This can only happen if the
photon energy equals the
difference in energy between
the two states
DE = h f
Jsh 34
1063.6 


EMISSION
When an atom in an excited
state undergoes a transition
to the ground state, it loses
energy by emitting a
photon whose energy
corresponds to the
difference in energy
between the two states
DE = h f
Jsh 34
1063.6 


What do you know so far?
• What is new in Bohr’s model?
• What is the light made of?
• How can an atom absorb/emit energy?
• How much energy can the atom
absorb/emit?

HYDROGEN BOHR’S ATOM
• Allowed orbits are labeled
with the integer n
(quantum number) from
the inner to the outer one
• Orbits have increasing
radius, proportional to the
integer number n2
• The electron energy is
proportional to 1/n2

HYDROGEN SPECTRUM
• Because each atom has a unique set of orbits,
• it has a unique set of energy differences and so
• it emits or absorbs a unique set of wavelengths,
• that is called the fingerprint of the element.

HYDROGEN SPECTRUM
• Each colour of the spectrum has different wavelength,
• so it is due to transitions between different energy
levels

GAS SPECTRA
Each element has characteristic emission and
absorption spectra, so scientists can use such
spectra to analyze the composition of matter.

ACTIVITY
• Play the game!
Site link:
http://spiff.rit.edu/classes/phys301/lectures/spec_lines/Atoms_Nav.swf
• Read “how to”
• Do the exercises

SPECRTOSCOPY
Spectra are used to provide information about the composition of a
substance or an object.
In particular, astronomers use emission and absorption spectra to
determine the composition of stars and interstellar matter.
We now know that the sun contains large amounts of hydrogen,
iron, and carbon, along with smaller amounts of other elements.

But…
• Bohr’s model only works for one electron
(hydrogen-like) atoms
• It can’t explain multi-electron atoms spectra
• It can’t explain multiplets of spectral lines
• Sub-levels

Next lesson… over Bohr
• De Broglie's matter waves (wave–particle
duality)
• Schrödinger wave function
• Quantum numbers
• Heisenberg's Uncertainty Principle (1927)
• Probability wave
• Orbital, shell, sub-shell
• Electronic configuration
• Aufbau

Atomic spectra&models

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