Learning Outcomes
• Emission and absorption spectra of the hydrogen
atom .
• Balmer series in the emission spectrum as an
example.
• Line spectra as evidence for energy levels.
• Energy sub-levels.
• Viewing of emission spectra of elements using a
spectroscope or a spectrometer.
Atomic structure
Spectra
Spectroscope
In a light spectroscope,
light is focused into a thin
beam of parallel rays by a
lens, and then passed
through a prism or
diffraction grating that
separates the light into a
frequency spectrum.
Continuous Spectrum
Emission Spectra
Continuous spectrum
A Spectrum in
which all
wavelengths are
present between
certain limits.
Emission Spectrum
Emission spectrum
Spectrum lines
When light from an unknown
source is analyzed in a
spectroscope, the different patterns
of bright lines in the spectrum
reveal which elements emitted the
light. Such a pattern is called an
emission spectrum.
Absorption spectrum
Emission Spectrum
• Shows that atoms can emit
only specific energies
(discrete wavelengths,
discrete frequencies)
hypothesis: if atoms emit only
discrete wavelengths, maybe
atoms can have only discrete
energies
Balmer Series
• Balmer analysed the hydrogen
spectrum and found that
hydrogen emitted four bands of
light within the visible spectrum:
• Wavelength (nm) Color
• 656.2
red
• 486.1
blue
• 434.0
blue-violet
• 410.1
violet
Flame Test
• Flame Test
The following metals emit certain colours of light when their atoms
are excited.
• Metal
Colour
• Sodium (Na)
Yellow
• Lithium (Li)
Pink/Red
• Potassium (K)
Purple
• Copper (Cu)
Green
• Calcium (Ca)
Pink
• Barium (Ba)
Yellow/Green
• Strontium (Sr)
Red/Orange
Learning Outcomes
• Energy levels in atoms.
• Organisation of particles in atoms of
elements nos. 1–20 (numbers of electrons in
each main energy level).
• Classification of the first twenty elements in
the periodic table on the basis of the number
of outer electrons.
Bohr
Bohr’s theory
• Electrons revolve around nucleus in
orbits
• Electron in orbit has a fixed amount
of energy
• Orbits called energy levels
• If electron stays in level it neither
gains nor loses energy
Bohr
• Atom absorbs energy
• Electron jumps to higher level
• Atom unstable at higher levels. Electron falls back
to a lower level
• Atom loses or emits energy of a particular
frequency.
quantisation
• Electrons can have
only certain particular
values of energy
EVIDENCE FOR ENERGY
LEVELS
• In Hydrogen electron in lowest (n=1) level;
ground state
• Energy given; electron jumps to higher
level excited state
• Falls back and emits a definite amount of
energy
• Energy appears as a line of a particular
colour
colours
• Energy emitted
depends on the
jumps
• Different jumps
emit different
amounts of energy
and hence different
colours
Bohr Diagram
Atomic structure 2
Learning Outcomes
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Energy sub-levels.
Heisenberg uncertainty principle.
Wave nature of the electron. (Non-mathematical treatment in
both cases.)
Atomic orbitals. Shapes of s and p orbitals.
Building up of electronic structure of the first 36 elements.
Electronic configurations of ions of s- and p-block elements only.
Arrangement of electrons in individual orbitals of p-block atoms.
Heisenberg
• We cannot know both the position and
speed of an electron
• Therefore we cannot describe how an
electron moves in an atom
.
Einstein
• .
•
De Broglie
• Matter has
wave
characteristics
Electrons were both particles and waves
Same for all sub-atomic particles
Matter exists as particles and waves at the same time.
The electron as a wave
Orbital
• A region in
space where the
probability of
finding an
electron of a
particular is
high
Electrons moving
Electron paths
Main levels AND THE
NUMBER OF ELECTRONS
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1 = 2e
2 = 8e
3 = 18e
4 = 32e
Sub-levels
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Each main level has sub-levels
1has s sub-level only
2 has s and p sub-levels
3 has s,p and d sub-levels
4 has s,p,d and f sub-levels
Energy of sub-levels spd
1s
2s
2p
3d
Electrons in sub-levels
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s = 2e
p = 6e
d = 10e
f = 14e
Sub-levels
• 1 = s(2e)
• 2 = s(2e) + p(6e) = 8e
• 3 = s(2e) + p(6e) + d(10e) = 18e
The "p" orbital is dumb belled shaped and each
P sub level is made of three "p" orbitals (because
the P sub level can hold 6 electrons and every
orbital holds 2 electrons)
P-orbitals
P-orbitals
Electrons in orbitals
• S holds 2e
• 3 p orbitals each holds only 2e
• 5 d orbitals each holds only 2e
Pauli’s exclusion principle
• Orbital can only hold
2electrons and these
electrons must have
opposite spins
Pauli's exclusion principle
Aufbau principle
• Electrons fill levels in a specific
order.
• 1s 2s 2p 3s 3p 4s 3d 4p
AUFBAU
Hunds rule
• When filling up
the orbitals in a
sublevel
electrons fill
then singly at
first.
5 electrons
6 electrons Hund’s rule
Electron Configurations
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He, 2, helium : 1s2
Ne, 10, neon: 1s2 2s2 2p6
Ar, 18, argon : 1s2 2s2 2p6 3s2 3p6
Kr, 36, krypton : 1s2 2s2 2p6 3s2 3p6 4s2 3d10
4p6
Exceptions to Electron
configuration rules
• Cr
• Half-filled orbitals give greater stability
• 1s2 2s2 2p6 3s2 3p6 3d4 4s2 1s2 2s2 2p6 3s2 3p6
3d5 4s1
• Cu
• Full 3d sub-level gives greater stability
• 1s2 2s2 2p6 3s2 3p6 3d9 4s2  1s2 2s2 2p6 3s2 3p6
3d10 4s1
Electron Configurations (ions)
• F-, 10, Flouride: [1s2 2s2 2p6 ]• Cl-, 18, Chloride : [1s2 2s2 2p6 3s2 3p6]• Na+, 10, Sodium ion: [1s2 2s2 2p6 ]+