Arrangement of electrons in the atom

1. Arrangement of Electrons in the
Atom
Chapter 3

2. Location of the Electron
The next question to be
answered by atomic scientists
concerned the location of the
electrons!
Ernest Rutherford had a proposal
for this!

3. Ernest Rutherford proposed that the electrons
could orbit the nucleus like the planets orbit the
sun.
• This explanation however defied the laws of
physics
• Electrons revolving around the nucleus would
lose energy and would spiral into the nucleus

4. Rutherford’s Explanation was obviously wrong and another
model of the arrangement of electrons had to be devised
It was a young Danish
Physicist called Niels
Bohr who provided an
insight into the
arrangement of
Electrons in the Atom
that helped to solve
the problem

5. Bohr’s Study of Spectra
• Bohr developed his
theory about the
arrangement of
electrons in atoms by
studying what were
known as the Spectra of
Elements

6. White Light
• White light when
passed through a prism
is broken up into an
array of colours
• Such an array is called a
spectrum
• Thus the spectrum is
the spread of colours
that come out of the
prism

7. Continuous Spectrum
• This spread or array of
colours is known as a
continuous spectrum or
rainbow
• Rainbows are formed
when the water
droplets in the sky act
as prisms and separate
the sunlight into its
colours

8. But what happens if we use a type of light
that is not made from white light but only
part of it?
• If a glass tube is filled
with Hydrogen at low
pressure and an electric
current is passed
through it, a spectrum
is produced that is
different to that given
by white light

9. A Line Spectrum for Hydrogen
• Instead of seeing a Continuous spectrum of all
the colours of the rainbow Bohr saw a series of
narrow lines

10. Line or Emission Spectrum
• Since it consists of lines it is called a Line Spectrum
to distinguish it from a Continuous Spectrum ( or
Rainbow) produced using white light.
• It is more accurately called an Emission Spectrum
as it has been emitted when the electric current
was passed through the hydrogen gas

11. Emission Spectra of Different Elements
• When the Hydrogen in
the discharge tube was
replaced by other
elements like Sodium
and Mercury it was
found that these
elements also produced
spectra.
• Each of these spectra
were unique to that
element

12. Emission Spectra of Different Elements

13. Each Element has a unique Emission
Spectrum
• This meant that • Elements present in the
examining the spectrum sample of salt being
produced from light burnt could then be
from a burning sample identified by comparing
of a salt should give you them to known spectra
a unique pattern of stored on a computer
lines database.
• This pattern could then
be compared to known
spectra

14. Spectra of most elements of the
periodic table
• http://jersey.uoregon.edu/vlab/elements/Ele
ments.html

15. To Study Emission Spectra using a
Spectrometer
• Experiment 3.1
• Page 12
• Method A and Method B
• To be written up properly
following guidelines as
shown for homework

16. Flame Tests
• Experiment 3.2 Page 13
• When salts of certain
metals are heated in a
bunsen burner’s flame
the colours obtained
can be used to identify
the metals in unknown
compounds

17. Results for Flame Test
• Metal Present in the Colour
Salt being burnt
• Lithium • Crimson
• Potassium • Lilac
• Barium • Green
• Strontium • Red
• Copper • Blue-green
• Sodium • Yellow

19. But what has the study of spectra
got to do with the structure of the
atom?

20. Niels Bohr’s Insight
Bohr realised that any model of the Atom
needed to explain two observations
1) Why the Emission Spectra of the Elements
are Line Spectra rather than Continuous
Spectra,
2) Why the Emission Spectrum of each element
is unique to that element

21. Bohr’s Explanation for the Emission Line Spectrum of
Hydrogen
1) Electrons revolve around the nucleus in fixed paths called
Orbits or Energy levels
2) Electrons in any one orbit have a fixed or quantised
amount of energy
3) Electrons in an energy level do not gain or lose energy
4) When atoms absorb energy electrons jump from a lower
energy level to a higher energy level
5) At these higher levels the electrons are less stable and do
not remain there for long but fall back down
6) When an electron falls back to any energy level it loses
energy in the form of light

22. These photons of light because they have a fixed amount of energy also
have a specific frequency and thus colour

23. Bohr represented each energy level by the letter n
He called the lowest energy level the n = 1 level
The next highest the n = 2 level
and so on….

24. Hydrogen’s electron is
normally found at the
n = 1 level
This electron is said to
be in the ground state
(or unexcited state)
When heated and
after absorbing energy
it jumps to a higher
level or is said to have
an excited state

25. After remaining a short time
it drops down to a lower
energy level emitting a
definite amount of energy
This definite energy is equal
to the difference between
the two energy levels
This definite amount of energy appears as a line of a
particular colour in the emission spectrum
Each colour corresponds to a particular wavelength or
frequency of light

26. Spectra of most elements of the
periodic table
• http://jersey.uoregon.edu/vlab/elements/Ele
ments.html

27. The Types of Light or Energy Transitions (or Emission
Lines) produced ranged from Infra-Red to Visible to
Ultra-Violet Light
So some are visible and others are invisible.

28. A Mathematical relationship exists
between the energy emitted and
frequency of the light
E = hf
• E is the amount of energy emitted from the
atom
• h is just a number called Planck’s constant
• f is the frequency or wavelength of the light
emitted
[Shown as ∆E = hv in diagrams above where v = wavelength]

29. Each definite amount
of energy emitted gives rise to a line in the
emission spectrum
This can be calculated from the equation
E2 - E1 = hf
where E represents the n energy level
Since only definite amounts of energy are
emitted this implies that electrons can occupy
only definite energy levels
Therefore energy levels must exist in the atom

30. Bohr’s Theory
• Bohr examined the spectrum of hydrogen and
measured the wavelengths of the visible light
seen through the spectroscope
• He compared these values to those calculated
using his theory and found them to be an
exact match

31. Electromagnetic Spectrum
He predicted the existence of other series
of lines in the ultra-violet and infra-red
regions of the spectrum both of which are
invisible

32. An Element’s Unique Emission
Spectrum
• Bohr’s Theory also explained why each
electron had its own unique emission
spectrum.
• Since each element has its own particular
number of electrons then there will be
different numbers and types of transitions for
each element, thus giving rise to a different
emission line spectrum in each case.

33. Emission Series
Lymann Invisible - Ultra-violet
Balmer Visible
Paschen Invisible - Infra-Red

34. Absorption Spectra
• There is another type of spectrum apart from the
emission spectrum
• The Absorption spectrum is obtained when white light
is passed through a gaseous sample of an element and
analysed
• It is found that the light coming out has certain
wavelengths missing or dark lines present

35. Natural Sunlight’s Absorption
Spectrum

36. Emission vs. Absorption Spectra
Emission Absorption
• Produced when a hot gas •Produced when white light is
glows giving off light shone through a tube of gas
which absorbs some of the
light

37. Emission vs. Absorption Spectrums
Emission Spectrum Absorption Spectrum
• Consists of coloured lines • Consists of dark lines
against a dark background against a coloured
background
Atomic Absorption Spectrometry is therefore a very
useful analytical tool used by chemists to detect the
presence of certain elements and to measure the
concentrations of these elements

38. Energy Sub-levels
• As time went by the study of spectra became
ever more sophisticated
• Scientists now found that many lines which
appeared to be one were in fact made up
several lines close together
• For example what appeared to be a single
yellow line in Sodium’s emission spectrum was
found in fact to consist of two yellow lines
very close together

39. Sodium’s Spectrum
Full Spectrum showing a
single yellow line Zooming in on the yellow line

40. Energy Sub-Levels
• These two lines could not be due to electrons
dropping to two different energy levels as this
would give rise to lines much further apart
• In order to explain this observation scientist
proposed that -
Each main energy level except the first was
made up of a number of sublevels all of which
were close in energy

41. Energy Sub-Levels
It was discovered that
the number of sub-
levels was the same as
the value of n for the
main energy level
So the n =2 main
energy level had 2
sub-levels
The n = 3 main energy
level had 3 sub-levels
And the n = 4 main
energy level had 4
sub-levels

42. S, p, d & f sub-levels
These sub-levels were
now labelled
The sub-level of
lowest energy being
called the s sub-level
The next highest
called the p sub-level
The one above that as
the d sub-level
And the one of highest
energy was called the
f sub-level

43. Wave Nature of the Electron
• Bohr’s Theory works very well for hydrogen
but when his theory is applied to atoms with
more than one electron, it fails to account or
many of the lines in the emission spectra of
these atoms
• So for other elements

44. Louis de Broglie
In 1924 a French
Scientist called
Louis Le Broglie
suggested that all
moving particles
had a wave motion
associated with
them
This was called a
‘Wave Particle
Duality’

45. Wave-Particle Duality

46. Werner Heisenberg
If the electron has a
wave motion it clearly
is not travelling along
a precise path or
energy level predicted
by Bohr
A German Physicist
tackled this problem
mathematically and
put forward a very
famous principle
called Heisenberg
Uncertainty Principle

47. Heisenberg Uncertainty Principle
• “It is impossible to measure at the same
time both the velocity and the position of
the electron”

48. Improving Bohr’s Idea
• Bohr’s Model saw electrons moving with a
certain speed in orbits at fixed distances from the
nucleus.
• Heisenberg stated that you cannot say this about
the electron as you cannot measure both the
speed and the distance from the nucleus at the
same time
• This led to scientists to change Bohr’s idea to the
probability of finding an electron at a particular
position inside the atom and a new picture of the
atom

49. Atomic Orbitals
• Imagine taking hundreds of photos of an atom
of hydrogen and its electron spinning around
it and combining them all together. You would
get a picture like the one shown below