Light
• White light is made up a continuous
spectrum of colours
• Light is made up of
waves.
•Light is
a form
of
energy
Niels Bohr
• Danish physicist
• University of Copenhagen
• Repeated the experiment but used light
from a hydrogen discharge tube instead
of white light
• A Hydrogen discharge tube is simply a gas
tube filled with hydrogen at low pressure
through which an electric current is passed
Niels Bohr
• Using a hydrogen discharge tube he saw a
series of narrow lines of light emitted instead
of a continuous spectrum of light
• Since the spectrum consists of lines , it is called
an emission line spectrum
Niels Bohr
• Replacing the hydrogen in the discharge
tube by other elements like sodium or
mercury
• These elements also produced line
spectra
• Each element has its own unique emission
line spectrum
• This spectrum is like a ‘fingerprint’ for the
element
Each element has its own unique emission
line spectrum
Like a fingerprint!
Emission line spectrum
• Often used by chemists to analyse
materials for the presence of certain
elements
e.g. an emission spectrometer is used to
analyse steel for different metals
Different elements emit different
colours
• Sodium street lamps
are yellow
• Mercury vapour
lamps are blue
• Neon advertising
signs are red
Spectrums
(White light)
Spread of colours
Narrow coloured lines
On black background
Narrow black lines
on coloured background
Introduction
• When metal ions are heated
strongly in a Bunsen flame
• or excited by high voltage electric
currents
• characteristic colours are
produced.
Na
Li
K
•Each metal produces
a different colour.
EXPT: Flame tests with salts of
lithium, sodium, potassium,
barium, strontium and copper
The colours imparted to a flame
by the given elements
Element
Colour
Lithium Li
Crimson
Sodium Na
Yellow
Potassium K
Lilac
Barium Ba
Green
Copper Cu
Blue-green
Strontium Sr
Red
Fireworks!
• Strontium and barium flames used in
firework displays
• E.g.
Strontium nitrate gives red colour
to fireworks
Barium nitrate gives green colour
•Each metal produces
a different colour.
•How are these
colours produced?
Explaining the evidence!
Where are electrons located?
• Rutherford
proposed that the
electrons were
revolving around
the nucleus like the
planets revolve
around the sun.
• This can’t happen as it
defies the laws of physics.
• Electrons revolving around
the nucleus would lose
energy and would spiral into
the nucleus.
The
atom
would
collapse
Niels Bohr
• Came up with a new idea of
how the electrons might be
arranged around the
nucleus.
• He used Emission spectra.
• Emission spectra
known for hundreds of
years
• Niels Bohr unlocked
their secret
• He concluded that
electrons revolve
around the nucleus in
fixed paths called
orbits
• Won the 1922 Nobel
Prize for Physics
Emission Spectrum
Emission Spectrum
• Coloured lines against a black
background.
• Why the emission spectra are line
spectra rather than continuous?
• Why the emission spectra of each
element is unique?
How are the
lines of the
line spectrum
formed?
Bohr’s Theory
 Bohr used the emission
spectrum of Hydrogen as
evidence for his theory.
1. Electrons revolve around the
nucleus in fixed paths called
orbits (energy levels). (false!)
Bohrs theory
2. Each electron in an
orbit has a definite
amount of energy Energy level.
Therefore Orbits are also called energy levels
Bohrs theory
• An energy level is the fixed
amount of energy that an
electron can have. (definition)
• Bohr used n to represent
energy level.
Bohrs theory
3. Once electrons
stay in an energy
level, they will not
lose or gain energy.
How did Bohr use
emission line spectra to
explain his theory???
BLUE
GREEN
YELLOW
RED
Emission line spectrum
• Each coloured line has a
different frequency• representing a different
amount of energy
• and therefore is different
colour.
Niels Bohr
4. When an atom absorbs energy,
electrons jump from a lower energy
level to a higher energy level. Electrons
are less stable in the higher energy levels
and do not remain there for long.
5. Energy is lost when an electron falls
from a higher energy level to a lower
energy level.
1
Energy absorbed when
Electron jumps from n = 1
To n = 2 energy level
n=2
Orbit of energy E1
n=1
Fixed
Energy
Levels
Orbit of energy E2
2
Electron
Energy emitted when
Electron falls from n = 2
To n = 1 energy level
Energy emitted as light
Spectrum
Excited Staten=4
UV
Excited Staten=3
Excited State unstable
and drops back down
Excited State
n=2
But only as far as
n = 2 this time
•Energy released as a photon
•Frequency proportional
to energy drop
V
i
s
i
b
l
e
IR
n=1
Ground State
Summary
• Electron normally in Ground State
• Energy supplied [ as heat or electricity]
• Electron jumps to higher energy level
• Now in Excited State
• Unstable
• Drops back to a lower level
• Energy that was absorbed to
make the jump up is now
released as a photon of light
• As the electron falls a definite
amount of energy is emitted in the
form of light.
• The energy emitted is the difference
(E2 –E1).
Energy is given out in the form of
light when an electron falls back
from the excited state to the ground
state
E2
Excited state
Energy
Light emitted
E1
Ground state
• This energy appears as a line
of colour in an emission
spectrum.
• The energy of the emitted line
can be calculated as
h x f = E2 –E1.
• h = planks constant
• f is the frequency of light
•Each line of light
has a specific
frequency (f).
• Since each line has a definite
frequency,
• and represents a definite amount
of energy (E2-E1)
• (the energy an electron is losing
as it is falling to a different energy
level)
•This means that the
electrons are
occupying definite
energy levels.
n
• Associated with each energy
level is an integer, n, called the
principal quantum number
• i.e. Bohr represented the energy
levels by the letter n
• When electron falls to
n = 1 level gives UV Range
n = 2 level gives Visible
Range
n = 3,4 or 5 levels gives IR
Range
n = 1 level gives Lymann
series
n = 2 level gives Balmer
series
n = 3,4 or 5 levels gives
Paschen series
If Hydrogen atoms have only one
electron,
why is it that so many electron
transitions are possible?
( lines in an emission spectrum)
•In a sample of
Hydrogen there are
millions of atoms of
hydrogen.
• When energy is given to the
sample of hydrogen,
• not all the atoms receive the
same amount of energy.
• Electrons will move to
different energy levels 1, 2
or 3 etc.
• When electrons are falling
back down
• they are falling from
different energy levels ,
• so different lines will be
formed in the emission
spectrum.
Why does each element
have a unique line
spectrum?
• As each element has a
different number of electrons,
• there will be a different colour
and number of lines in a line
spectrum for each element.
Uses
• Fluorescent light strips make
use of electron transitions
• Contain mercury vapour at low
pressure
• Electrons excited by current
• Electrons fall back to lower
energy levels and emit UV light
Uses
• Lasers make use of electron transitions
• Light Amplificaton by Stimulated
Emission of Radiation
• Contain helium and neon gas
• Flashlamps inside laser excite
electrons in gases
• Results in intense flash of light
emerging from laser
Atomic absorption spectrometry
• We saw how atoms emit light when
electrical or heat energy is
supplied to them  emission
spectrum
• Scientists also found that atoms
absorb light
Atomic absorption spectrometry
• If white light passed through
gaseous sample of element
• Light coming out has certain
wavelengths missing
• i.e. dark lines are observed in the
spectrum
• Called Atomic absorption
spectrum
Atomic absorption spectrometry
Narrow black lines
on coloured background
Atomic absorption spectrometry
• The dark lines represent missing
wavelengths and can tell us what
elements are present
• Absorption spectrum is like a
photographic negative of an emission
spectrum
Atomic absorption spectrometry
Atomic absorption spectrometry
• A very useful tool used by chemists to
detect the presence of certain elements
• And to measure the concentration of
these elements
• In practice chemists do not use white
light as the source of radiation
Atomic absorption spectrometry
• Instead, lamps which emit the radiation
of the element being analysed are used.
• E.g. a sodium lamp is used to measure
the amount of sodium present in a
sample being analysed
Atomic absorption spectrometry
• The sample containing the sodium
is held in a flame in the instrument
• Atomic absorption spectrometer
Atomic absorption spectrometry
• The intensity of light transmitted
through the flame before the sample is
introduced and after sample is
introduced is measured
• Instrument indicates amount of light
absorbed and from this
• Calculates the concentration of sodium
in sample
Uses
• Analysis of water for heavy metals
like
• Lead
• Mercury
• Cadmium
Applications of emission and
absorption spectra
Emission spectra
Atomic absorption
spectra
1. ID of elements
1. ID of elements
2. Sodium street lamps 2. Conc of elements
3. Giving particular
colours to fireworks
3. Conc of heavy
metals in water
Back to emission spectra
Coloured Lines formed due to electrons
falling from higher energy levels and
emitting photons of light
Energy Sublevels
• Lines in an emission
spectrum were found to
consist of a small number of
finer lines very close
together.
Energy sublevels
• The lines were too close
• to represent two different energy
levels
• as this would give rise to lines
much further apart in the
spectrum.
Energy sublevels
• This proved that energy
levels consisted of sub
levels.
• They are s, p, d and f
sublevels.
Energy sublevels
Lowest
energy
Highest
energy
• First energy level – 1 s sublevel
• Second energy level – 2 sub levels. 1 s
sublevel and 1 p sublevel.
• Third energy level – 3 sub levels. 1 s
sublevel and 1 p sublevel and 1 d
sublevel.
• Fourth energy level – 4 sub levels. 1 s
sublevel and 1 p sublevel, 1 d sublevel
and 1 f sublevel.
Energy sub-levels
• d sub-level (holds ten electrons) increasing
• p sub-level (holds six electrons)
energy
• s sub-level ( holds two electrons)
• The number of sub-levels in a main energy
level is the same as the no. of that level
e.g. the n = 3 energy level has 3 sub-levels
Energy sublevels
Energy level
n=1
n=2
n=3
n=4
Sub-level
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
Capacity
2
2
2
6
8
2
6
18
10
2
6
32
10
14
4f
n=4
4d
4p
Energy
3d
4s
n=3
3p
3s
n=2
2p
2s
n=1
Main energy level
1s
Sublevel
Note: 4s sublevel is lower in energy than the 3d sublevel
Bohr’s atomic theory
was later modified.
Give one reason why
this theory was
updated.
• Bohr’s theory worked perfectly to
explain the emission spectrum of
Hydrogen.
• When it was applied to atoms with
more than one electron,
• it failed to account for many of the
lines in the emission spectra.
• The term orbit was used by
Bohr to describe the fixed path
of an electron
• traveling with a certain velocity
• at a precise distance from the
nucleus.
• This is not true.
De Broglie
• 1927 De Broglie
discovered that
electrons traveled
around the nucleus
with a wave
motion.
• This idea that a particle can
behave like a wave is often
referred to as
‘Wave-particle duality’
Bohr’s Theory
 Bohr used the emission
spectrum of Hydrogen as
evidence for his theory.
 Electrons revolve around the
nucleus in fixed paths called
orbits (energy levels). ( false)
•This meant that
Bohr’s Theory could
not be true.
Heisenberg
see book
• stated that it was impossible to
measure at the same time
• the speed of an electron and
• its distance from the nucleus.
• (A beam of light is used to find the
position of the electron, but this
beam would change the speed of
the electron.)
Heisenberg’s
Uncertainty principle
• stated that it was impossible
to
• measure at the same time
• the speed of an electron
and
• its distance from the
nucleus.
• The notion of
electrons having a
fixed path around the
nucleus was changed
to that of the orbital.
An atomic orbital
• is the region in space
• within which
• there is a high probability
of finding an electron.
• Do not confuse ‘orbit’ with
‘orbital’
• Bohr = orbit
• Schrodinger = orbital
Schrödinger
• Calculated the
probability that an
electron would be at a
particular spot in the
orbital but not know for
sure.
• He found the shapes of
4 types of orbitals.
• Spdf
•An orbital can hold
two electrons of an
opposite spin.
• An s orbital
can hold two
electrons
• S orbital is
spherical
shaped.
All s orbitals are spherical in shape
1s
2s
3s
A p orbital can hold 6
electrons
P orbital is dumb- bell
shaped
• Each p sublevel consists of three
parts
• Px
• Py
• Pz
• x, y and z stand for the 3 main axes
to represent their direction
• 3 p orbitals are at right angles to
each other
Dumb-bell shaped
Shape & orientation of orbitals
Energy
level
n=1
Orbital
Shape
Orientation
1s
Spherical
n=2
2s
2px, 2py and 2pz
Spherical
Dumb-bell
At right angles
to each other
n=3
3s
3px, 3py and 3pz
Spherical
Dumb-bell
At right angles
to each other
Spherical
Dumb-bell
At right angles
to each other
Five 3d orbitals
n=4
4s
4px, 4py and 4pz
Five 4d orbitals
Seven 4f orbitals
Recheck
•
•
•
•
•
•
•
Bohr’s study of spectra
Explaining the evidence: the Bohr theory
Atomic absorption spectrometry
Energy levels
Energy sub-levels
Wave nature of the electron
Atomic orbitals
Shorthand Configuration
A neon's electron configuration (1s22s22p6)
B
third energy level
[Ne] 3s1
C
D
one electron in the s orbital
orbital shape
Na = [1s22s22p6] 3s1
electron configuration
• D Orbital
• www.uwgb.edu/dutchs/PETROLGY/Wh
atElmsLookLike.HTM
The Aufbau
Principle
• states that when building up
the electronic configuration of
an atom in its ground state• the electrons occupy the
lowest available energy level.
Hunds Rule of Maximum
Multiplicity states
• that when two or more
orbital’s of equal energy are
available,
• the electrons must occupy
them singly before filling
them in pairs.
3px1 3py1 3pz1
3px2 3py1 3pz1
The Pauli Exclusion principle
states
• that no more than two
electrons
• may occupy an orbital and
• they must have opposite spin.
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Hydrogen
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
H = 1s1
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Helium
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
He = 1s2
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Lithium
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
Li = 1s22s1
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Carbon
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
C = 1s22s22p2
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Nitrogen
4f
Bohr Model
N
Hund’s Rule “maximum
number of unpaired
orbitals”.
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
N = 1s22s22p3
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Fluorine
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
F = 1s22s22p5
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Aluminum
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
Al = 1s22s22p63s23p1
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Argon
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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Fe La
Ar = 1s22s22p63s23p6
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
Iron
4f
Bohr Model
N
2s
2p
1s
Electron Configuration
Fe = 1s22s22p63s23p64s23d6
NUCLEUS
H He Li C N Al Ar F
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Fe La
Arbitrary Energy Scale
Energy Level Diagram
6s
6p
5d
5s
5p
4d
4s
4p
3d
3s
3p
4f
Lanthanum
Bohr Model
N
2s
2p
1s
Electron Configuration
NUCLEUS
H He Li C N Al Ar F
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La = 1s22s22p63s23p64s23d10
Fe La 4s23d104p65s24d105p66s25d1