Chapter 26 – Electronic Structure of the Atom
Section A – Energy levels
Bohr Model
Neils Bohr proposed a theory to explain the arrangement of ___________ around the nucleus. In
this theory, called the Bohr Model, he suggested that electrons orbit around the nucleus in
_______________________.
Energy level
*** The first energy level is closest to the nucleus and is filled with electrons first. When
this energy level is full the second energy level will be filled and so on.
Evidence for energy levels
Bohr developed his theory by studying elements with a spectroscope (an instrument which analyses
the light emitted by elements)
Continuous spectrum
When white light is passed through a ___________ it will separate into a band (spectrum) of colours
which can be viewed with a spectroscope and produces a _______________spectrum (no breaks or
gaps in the spectrum)
Line Spectrum
If electricity is passed through hydrogen gas in a cathode ray tube it will also emit light. When the
light emitted is viewed with a spectroscope it produces a _____________ spectrum (a series of
________ lines against a ___________ background)
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Chapter 26 – Electronic Structure of the Atom
Bohr’s explanation of energy levels using spectra
1.
The electron in the hydrogen atom must occupy an energy level. It usually occupies the n = 1
level – the “_____________________”
2.
When the electron absorbs ______________ it will move to a ____________
energy level – an “________________state”. The amount of energy it absorbs is
exactly the energy difference between the two _____________ levels.
3.
The electron is _______________ at this higher energy level .The electron falls
back down to the lower energy level and gives out _____________. The amount of
energy it emits is exactly the energy difference between the two levels.
E2 – E1 = hf
4. The electron emits a photon - a ________________ amount of energy corresponding to the
difference in energy between those two levels.
5. We can see the definite amount energy that is emitted as a line in the emission spectrum of
Hydrogen.
The Hydrogen emission line spectrum
When the hydrogen atoms emit light some of it is visible (can be
seen with the naked eye), and some of it is invisible to the eye
(ultraviolet and infrared)
***Even though each hydrogen atom only has 1 electron in a
sample of the gas there may be millions of different atoms, and
each may absorb different amounts of energy causing different
lines to form in the emission spectrum from the various
transitions that occur.
The Balmer series contains all of the lines that are in the visible
region of the hydrogen spectrum. These lines are produced when
an electron falls from higher energy levels to the n= 2 energy
level.
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Chapter 26 – Electronic Structure of the Atom
Emission of light by other elements
 The same cathode tube experiment can be repeated with other
elements and each element will give a unique emission line
spectrum,
which
can
be
used
for
__________________purposes.
Line spectra are ________________ to each element because
in atoms of that element have unique electronic configurations
so different transitions can occur
 We sometimes see this phenomenon in everyday life. For example streetlights are simply
discharge tubes filled with ___________gas, which glow with an ______ light when electricity
flow through them.
 Some elements also emit light when ignited - Chemists have used this to make_______________.
We will observe this by doing some flame tests.
Mandatory experiment – Flame tests
The energy differences between energy levels in metal atoms vary from metal to metal. Using
energy from the flame, electrons in the metal atoms move to higher energy levels, and then return
to lower energy levels, emitting light whose energies in each case is equal to the energy difference
between the higher energy level and the lower energy level.
Procedure



Use a damp wooden splint and dip in the salt to be tested. If too much salt is used it may fall
into the Bunsen burner and cause a prolonged colour tinge in the flame.
Gently and quickly pass the splint through the ____________ flame of the Bunsen burner.
To avoid cross __________________use a new splint each time a new salt into the flame.
Metal present
Colour
Atomic absorption spectra
An absorption spectrum shows the light that you see with a
spectrometer after white light has been passed through an
element. It is a series of ___________ lines against a
_______________ background.
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Chapter 26 – Electronic Structure of the Atom
The emission and absorption spectra of the same element are like a
matching pair –the coloured lines in the emission spectrum match
the black lines in the absorption spectrum
Application: The __________________________________(AAS) is
used to measure the amount of light absorbed by a sample and a
chemist can thus determine the concentration of a particular
element in a sample e.g. lead in a blood sample, or mercury in a
water sample.
Limitations of Bohr’s Model
1. Bohr’s theory could not be used to explain the emission spectrum of elements other than
Hydrogen ( with more than _______ electron).
2. It did not account for Wave particle duality.
*De Broglie was a French physicist, who discovered that electrons have properties of
_____________ as well as particles. This led to the realisation that electrons are NOT really be
travelling at a precise distance from the nucleus in orbits, as Bohr had previously suggested.
3. It was conflicting with Heisenberg’s uncertainty principle
Heisenberg’s uncertainty principle
Section B – Atomic Orbitals
Since it is impossible to determine the exact position of an electron at any one moment it is better to
refer to the ___________________ of finding an electron in a region of space within the atom.
Atomic orbital
First energy level ( n =




) Here there is only one orbital in which electrons can be found.
The orbital is called a ________orbital.
It is on the first energy level so it is called “1s”
It is __________________ in shape – there is the same probability of finding the electrons
any direction form the nucleus.
The orbital can hold ___________ electrons.
In total this energy level can fit _______ electrons
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Chapter 26 – Electronic Structure of the Atom
Second energy level ( n = )
Here there are __________ orbitals in which electrons can be found.


There is one _____ orbital which can hold _____ electrons. It is refered to as “2s” because it
is an s orbital on the second energy level.
There are also 3 _______ orbitals which can hold _____ electrons EACH. They are
_______________ shaped and are named _____, _____ and ____.
In total this energy level can fit _______ electrons
Third energy level ( n = )
Here there are __________ orbitals in which electrons can be found.

There is one _____ orbital which can hold _____ electrons refered to as “____”.

There are also 3 _______ orbitals which can hold _____ electrons EACH. (referred to as
“______”,“______” and “______”.

There are five d orbitals called “3d”, which can hold two electrons EACH
In total this energy level can fit _______
electrons
Fourth energy level ( n = )
Here there are even more orbitals in which electrons can be found! This includes an s orbital, 3 p
orbitals, 5 d orbitals, and further orbitals called f orbitals which can hold even more electrons.
Energy sublevel
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Chapter 26 – Electronic Structure of the Atom
Section C – Electronic configurations
. The electronic configuration of an atom or ion tells how the electrons are arranged in their energy
levels, sublevels and will also tell which types of orbitals are present.
An energy level diagram will show this in a diagram form. Here is the format for a diagram like this:
1.
2.
3.
4.
Start to fill the electrons into the shells starting at the _____________ level and then going up.
A maximum of ___ electrons can fit into each orbital.
Where there is more than 1 orbital on an energy level fill them singly before pairing them up.
Note that the ____ energy sublevel has lower energy than the 3d level and must be filled first.
Example 1 – Write out an energy level diagram for an atom of Silicon, and use this to write out its
electronic configuration
Electronic configuration:
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Chapter 26 – Electronic Structure of the Atom
Exceptions to the rules:
*Copper has 29 electrons, and its electron configuration is :
This unexpected configuration is because there is extra stability given with a _________ 3d sublevel!
**Chromium has 25 electrons, and its electron configuration is :
This unexpected configuration is because there is extra stability given with a _______3d sublevel!
Electronic configuration of ions
These work in the same way, but you must remember than an ion is an atom that has lost or gained
electrons that it originally had –
Each + charge means it has _______ an electron
Each – charge means it has _______ an electron
Example 2 – Write out an energy level diagram for a Cl- ion and use this to write out it’s electronic
configuration
Electronic configuration:
Arrangement of electrons within individual p orbitals
Arrows are used to represent the spin which an electron has in an orbital, and if two electrons are in
the same orbital then their spins will be in ______________ directions
Sulfur
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Chapter 26 – Electronic Structure of the Atom
Section D- Trends in atomic radii
Atomic radius
As you go down the periodic table the atomic radii ______________
1. The number of ____________ increase
2. The screening effect increases - the electrons in the outer shells are being screened from the
extra nuclear charge by the inner shells of electrons.
As you go across the periodic table the atomic radii ______________
1. The _______________________________ charge increases, which pulls the electrons closer
to the nucleus
2. There is no extra ______________ effect because there are no extra shells of electrons
Section E – Ionisation Energies
Ionisation energy
This can be represented as:
Because energy must be taken in for this to happen it is called an ___________ reaction.
Trends in ionisation energies
Generally as you go down the periodic table the ionisation energies ______________
1. The atomic radii ______________________ - the furthermost electrons are further
from the nucleus
2. The screening effect increases - the electrons in the outer shells are being screened
from the extra nuclear charge by the inner shells of electrons.
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Chapter 26 – Electronic Structure of the Atom
Generally as you go across the periodic table the ionisation energies ______________
1. The nuclear charge ________________ - pulling the electrons closer to the nucleus
2. The atomic radii ______________________ - the furthermost electrons are closer to
the nucleus
***However there are some exceptions to this general trend as can be seen in the diagram:
For example going across the second period the elements _____________ and
________________ have higher ionisation energies than you might expect.
 This is due to their electronic configurations:
For Beryllium (_______________) it has a full
______ sublevel which makes it more __________
and more _____________ is needed to remove the
electron from it.
 For Nitrogen ( ________________) it has a ______
full __________ sublevel, making it more stable and
more energy is needed to remove the electron from
it.
(Similar exceptions can be seen with _____________ and _______________ in third period
– make sure you can explain these too!)
Successive ionisation energies
Second Ionisation energy
This can be represented as:
 The second ionisation energy is always ___________ than the first because
1. in an ion the atomic radius___________ and the electrons are closer to the
nucleus
2. There is greater attraction between the nucleus and the electron in a
positive ion
 If an electron is removed from a full energy level than the ionisation energy
increases_________________. This is because of the extra stability which Is
disturbed when the electron is removed. This same effect will be seen to a
lesser extent when an electron is removed from a full _________________.
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Evidence for
the existence
of energy
levels
Page 9
Chapter 26 – Electronic Structure of the Atom
Section F – More trends in the Periodic table
Group 1 – The alkali metals
Going down the group reactivity ___________ because:
1. The atomic radius ____________
2. Although the _________ charge increases the ____________ effect
cancels this out
It becomes easier to remove the outermost electron ( ionisation
energy decreases)
Group 7 – The halogens
Going down the group reactivity ___________ because:
1.
The atomic radius ____________
2.
Although the _________ charge increases the ____________ effect
cancels this out
It becomes harder for the atom to attract an electron from another
atom (electronegativity ______________)
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.
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Chapter 26 – Electronic Structure of the Atom
Check your learning of Electronic
structure of an Atom
Green
Orange
Red
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
HL only: Emission and absorption spectra of the
hydrogen atom – Balmer series in the emission
spectrum as an example.
Mandatory experiment 1.1* - Flame tests
HL only: Line spectra as evidence for energy levels.
HL only: Atomic absorption spectrometry ,fireworks,
streetlights
HL only: Energy sub-levels.
HL only: Heisenberg uncertainty principle.
HL only: Wave nature of the electron.
HL only: Atomic orbitals.
HL only: Shapes of s and p orbitals.
HL only: Building up of electronic structure of the first
36 elements.
HL only: Electronic configurations of ions of s- and pblock elements only.
HL only: Arrangement of electrons in individual
orbitals
Atomic radii
Explanations for general trends in values: (i) down a
group(ii) across a period
HL only: First ionisation energies.
HL only: Explanations for general trends in values: (i)
down a group (ii) across a period (main group elements)
and for exceptions to the general trends across a period.
HL only: Second and successive ionisation energies.
HL only: Evidence for energy levels provided by
successive ionisation energy values.
Dependence of chemical properties of elements on their
electronic structure.
Explanations in terms of atomic radius, screening effect
and nuclear charge for general trends in properties of
elements in groups I and VII.
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Chapter 26 – Electronic Structure of the Atom
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