Electron Arrangement
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2.3 Electron arrangement
Assessment statement
2.3.1 Describe the electromagnetic spectrum.
Teacher’s notes
Students should be able to identify the
ultraviolet, visible and infrared regions, and to
describe the variation in wavelength, frequency
and energy across the spectrum.
2.3.2 Distinguish between a continuous spectrum
and a line spectrum.
Students should be able to draw an energy level
2.3.3 Explain how the lines in the emission
spectrum of hydrogen are related to electron diagram, show transitions between different
energy levels and recognize that the lines in a
energy levels.
line spectrum are directly related to these
differences. An understanding of convergence is
expected. Series should be considered in the
ultraviolet, visible and infrared regions of the
spectrum. Calculations, knowledge of quantum
numbers and historical references will not be
assessed.
2.3.4 Deduce the electron arrangement for atoms
and ions up to Z = 20.
For example, 2.8.7 or 2,8,7 for Z = 17.
Electron Arrangement
2.3.1
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Assessment statement
Teacher’s notes
Describe the electromagnetic spectrum.
Students should be able to identify the ultraviolet,
visible and infrared regions, and to describe the
variation in wavelength, frequency and energy across
the spectrum.
• The study of the emission of light by atoms and ions is the most
effective technique for determining the structure of atoms.
Blackbody Radiation:
(Heating metal to glow RED
Or even WHITE)
Max Planck, 1900:
Energy, like matter, is discontinuous.
E = h
Prentice-Hall © 2002
General Chemistry: Chapter 9
Slide 4 of 50
The Nature of Energy
• Einstein used this
assumption to explain the
photoelectric effect.
• He concluded that energy is
proportional to frequency:
E = hf
where h is Planck’s constant,
6.63  10−34 J-s.
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The Photoelectric Effect
• Light striking the surface of certain metals
causes ejection of electrons.
• Depends on WAVELENGTH, not intensity
Prentice-Hall © 2002
General Chemistry: Chapter 9
Slide 6 of 50
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The Nature of Energy
• Therefore, if the wavelength of
light is known, the energy of
that light (one photon, a
packet) can be calculated, :
E = hf
(packet = quanta)
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Wavelength
• The color of light is sometimes defined by its wavelength (λ)(lambda)
rather than its frequency ()(nu).
• The two are related by the equation c =  λ
• All electromagnetic radiation travels at the same velocity: the speed of
light (c) =2.997925 x 108 m s-1 (3.00  108 ms-1 ) (meters per second)
• Frequency () (nu) in Hertz—Hz or s-1.
• Wavelength (λ) (lambda) in meters—m.
• cm
m
nm
Angstrom
(10-2 m) (10-6 m)
(10-9 m) (10-10 m)
Equations:
c = λ
λ = c/
(c = lambda nu)
= c/λ
pm
(10-12 m)
Wavelength
• The greater the velocity (the higher the frequency), the shorter the
wavelength.
Low 
High 
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Wavelength
Prentice-Hall © 2002
General Chemistry: Chapter 9
Slide 11 of 50
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ROYGBIV
Red
Orange
Yellow
700 nm
450 nm
Green
Blue
Indigo
Violet
Prentice-Hall © 2002
General Chemistry: Chapter 9
Slide 12 of 50
Continuous vs Line Spectrums
Assessment statement
2.3.2
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Teacher’s notes
Distinguish between a continuous spectrum and a
line spectrum.
• Electrons surround the nucleus in specific orbitals or energy levels.
• When an element is excited, it will often emit a light of a
characteristic color.
• Gases can be excited by passing an electrical discharge at low
pressure.
• When this excited gas or excited elements light is viewed through a
spectroscope a continuous spectrum is not observed, but a series
of very bright lines of specific colors with black spaces in-between
instead.
•
• This is known as a line spectrum.
Continuous vs Line Spectrums
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In a continuous spectrum, all wavelengths of visible light are
represented (like a rainbow) as is observed with “white” light.
Continuous vs Line Spectrums
In a line spectrum, only specific wavelengths are
represented. It comprises very bright lines of specific
colors with black space in between. This is from an
emission (energized samples “glowing”.)
Line spectra are unique for
every element and are used to
identify atoms (much like
fingerprints are used to identify
people).
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Continuous vs Line Spectrums
In a Absorption spectrum, only specific wavelengths are
missing from a continuous spectrum. It comprises colors
with black lines. This is an absorption (a source of white
light is being filtered by a sample)
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Continuous vs Absorption/Emission *
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Atomic Emission Spectra
• When an atom is excited, its electrons
gain energy and move to a higher energy
level.
• In order to return to the lower energy
levels, the electron must lose energy.
• It does this by giving out light.
Atomic Line Spectra
1. An electron in the atom
gains (absorbs) energy
from heating.
2. Electron jumps up an
energy level.
3. Electron is now unstable
(unwelcome) in this level
and is “kicked out”.
4. When the electron loses
the energy and come
back to the original
level, light is emitted.
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Energy Levels
2.3.3
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Assessment statement
Teacher’s notes
Explain how the lines in the emission spectrum of
hydrogen are related to electron energy levels.
Students should be able to draw an energy level
diagram, show transitions between different energy
levels and recognize that the lines in a line spectrum are
directly related to these differences. An understanding
of convergence is expected. Series should be considered
in the ultraviolet, visible and infrared regions of the
spectrum. Calculations, knowledge of quantum
numbers and historical references will not be assessed.
Energy Levels
• Because there are only certain allowed energy levels in the atom, there
are a limited number of amounts of energy that an electron can lose.
• The energy levels of atoms are not evenly spaced, like the rungs of a
ladder, but that the higher the energy level, the smaller the difference
in energy between successive energy levels becomes.
• This means that only certain frequencies can emit light and this is why
we get the line spectrum.
• This means that the lines of a spectrum will converge (get closer
together with increasing energy).
• The limit of this convergence indicates the energy required to
completely remove the electron from the atom (to ionize it) and so
may be used to determine the ionization energy.
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Energy Levels
The Atomic Emission Spectrum of
Hydrogen
• The emission spectrum of hydrogen is the simplest emission spectrum
because there is no electron – electron repulsion which will cause the
principle energy levels to split into different sub levels.
• When a voltage is applied across hydrogen gas, electromagnetic
radiation is emitted.
• This is not uniform, but concentrated into bright lines, indicating the
existence of only certain allowed electron energy levels.
• The study of this spectra allowed for the electron structure of
hydrogen to be deduced.
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The Atomic Emission Spectrum of
Hydrogen
The Atomic Emission Spectrum of
Hydrogen
• The spectrum is divided into three distinct series which occur in
different spectral regions (visible, IR, UV).
• Each series corresponds to transitions in which the electron falls to a
particular energy level.
• The reason why they occur in different spectral regions is that as the
energy level increases, they converge (they get closer together in
energy).
• This means that all transitions to the n=1 level include the large n=1 to
n=2 energy difference and so they are all high energy transitions found
in the UV region.
• For similar reasons, all transitions to the n=2 level are in the visible
region.
• All transitions to the n=3 level are in the infrared region.
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Electron Structure and the Periodic Table
2.3.4
Assessment statement
Teacher’s notes
Deduce the electron arrangement for atoms
and ions up to Z = 20.
For example, 2.8.7 or 2,8,7 for Z = 17.
• The most stable energy levels, or shells, are those closest to
the nucleus and these are filled before electrons start to fill
the higher levels.
• There is a maximum number of electrons that each energy
level can hold.
– The first can hold 2
– The second can hold up to 8
– Beyond this the situation gets more complex
Electronic Structure
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•
The number of electrons in each orbital is known as the electronic structure of the
atom.
•
Example Aluminum
– Aluminum has 13 electrons
• 2,8,3 is the electronic structure.
• 2 in the first, 8 in the second and 3 in the third
•
It is the electrons in the outermost (valence) shell that determine the physical and
chemical properties of the element.
•
Elements with three or less electrons in the valence shell are metals (except boron).
•
The electronic structure is closely related to the position on the periodic table.
•
The period from the periodic table gives us the number of energy levels that contain
electrons.
•
The group on the period table gives us the number of electrons in the valence shell.
Electronic Structure of the Periodic
Table (up to Z=20)
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Electronic Structure of the Periodic Table *
Electronic Structure of the Periodic Table *