4Electron and Electron Configuration.pptx

© Boardworks Ltd 2008
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The nucleus is:
 made up of protons and neutrons
 positively charged because of the
protons
 dense – it contains nearly all the
mass of the atom in a tiny space.
Electrons are:
 very small and light, and negatively charged
 able to be lost or gained in chemical reactions
 found thinly spread around the outside of the nucleus,
orbiting in layers called shells.
Protons, neutrons and electrons

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Protons, neutrons and electrons
positively charge
lost and gained in a
reaction
found thinly spread
outside the
nucleus
dense, contains the
mass of the atom

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Quantum Shells

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Electrons in the First Four Quantum Shell

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s Orbitals

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s Orbitals

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p Orbitals

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d Orbitals

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Orbitals

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How are electrons arranged?
Electrons are not evenly spread, but exist in layers called
shells. (The shells can also be called energy levels).
The arrangement of electrons in these shells is often
called the electron configuration.
Note that this diagram is not drawn to scale – the atom is
mostly empty space. If the electron shells are the size
shown, the nucleus would be too small to see.
1st shell
2nd shell
3rd shell

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How many electrons per shell?
Each shell has a maximum number of electrons that it can
hold. Electrons will fill the shells nearest the nucleus first.
1st shell holds
a maximum of
2 electrons
2nd shell holds
a maximum of
8 electrons
3rd shell holds
a maximum of
8 electrons
This electron arrangement is written as 2,8,8.

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Bohr Models
• Used to represent a model of an atom.
• To draw a Bohr model follow these steps:
(We will use Helium as an example)

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Making a Bohr Model Using Helium
1. Look to the periodic table and determine
how many protons, neutrons and electrons
are in 1 atom of helium.
P=____ N=_____ E=_____
2. Draw a circle and label the # of P and N in
the inside of the circle
P= 2
N= 2
2
2 2

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Making a Bohr Model Using Helium
3. Draw your 1st electron shell.
4. Draw up to 2 electrons in the 1st shell.
5. If you need to add more electrons, you need to
add more electron shells! Remember…2, 8, 8!!!
P= 2
N= 2
P= 2
N= 2

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Filling the shells and orbitals
• The most stable electronic configuration of an atom is the
one that has the lowest amount of energy.
• The order in which the subshells are filled depends on their
relative energy.
• The subshell with the lowest energy, the 1s, is therefore
filled first, followed by those that are successively higher in
energy.

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• also called the aufbau rule, states that in the ground
state of an atom or ion, electrons fill subshells of the
lowest available energy, then they fill subshells of higher
energy
• For example, the 1s subshell is filled before the 2s
subshell is occupied.
Aufbau Principle

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Calculate electron configurations

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Electronic configurations
• Representing electronic configurations
A detailed way of writing the electronic configuration of
an atom that includes information about the number of
electrons in each subshell is shown above for
hydrogen.

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• The electronic
configurations of some of
the elements after argon
are shown in Table 3.6.
• In this table part of the
electronic configuration of
each element is
represented by [Ar].
• This ‘noble gas core’
represents the electronic
configuration of argon:
1s2 2s2 2p6 3s2 3p6.
• This method is a
shorthand way of writing
electronic structures of
atoms with many
electrons.

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■ Electronic configuration of potassium
 Potassium has the electronic structure 1s2 2s2 2p6 3s2 3p6
4s1. The outer electron goes into the 4s subshell rather than
the 3d subshell because the 4s is below the 3d in terms of its
energy
■ Filling the 3d subshell
 After calcium, a new subshell becomes occupied.
 The next electron goes into a 3d subshell rather than a 4p
subshell.
 So scandium has the electronic configuration [Ar] 3d1 4s2.
 This is because electrons occupy the orbitals with the lowest
energy – the 3d subshell is just above the 4s subshell but
below the 4p subshell.
 This begins a pattern of filling the 3d subshell ending with
zinc.
 Zinc has the electronic configuration [Ar] 3d10 4s2.

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■ Chromium and copper
 The electronic configurations of chromium and copper do
not follow the expected pattern.
 Chromium has the electronic configuration [Ar]3d5 4s1
(rather than the expected [Ar]3d4 4s2).
 Copper has the electronic configuration [Ar]3d10 4s1 (rather
than the expected [Ar]3d9 4s2).
■ Gallium to krypton
 The electrons add to the 4p subshell because this is the next
highest energy level above the 3d.

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DA:
1. a. Name the three types of orbital present in the third
principal quantum shell. (3 pts)
b. State the maximum number of electrons that can be
found in each subshell of the third quantum shell. (3 pts)
2. Use 1s2 notation to give the electronic configurations of
the atoms with the following atomic numbers: (3 pts)
a. 16
b. 9
c. 20

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Which element?

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Valence Electrons are:
 The electrons in the outermost shell
 Responsible for atomic bonding
 Equal to the last digit of the group number
 How many valence electrons in this atom? What
group would it be in?
Valence electrons

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LEWIS (DOT) SYMBOLS FOR THE ELEMENTS
A Lewis dot structure for an atom consists of
the symbol for the element and one dot for
each valence electron.

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1) Find your element on the periodic table.
2) Determine the number of valence
electrons by looking at the group
(column)
3) This is how many electrons you will draw.

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1) Write the element
symbol.
2) Carbon is in the 14th
group, so it has 4
valence electrons.
3) Starting at the right,
draw 4 electrons, or
dots, counter-
clockwise around the
element symbol.

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1) Check your work.
2) Using your periodic
table, check that
Carbon is in the 4th
group.
3) You should have 4
total electrons, or
dots, drawn in for
Carbon.

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What would the Lewis Dot
Structure for Phosphorus
look like?

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Spin-Spin Pairing & Electronic Configuration
Hund’s Rule and Pauli Exclusion Principle
Hund’s rule states that electrons will occupy orbitals singly
before pairing takes place.
The Pauli Exclusion Principle states that two electrons cannot
occupy the same orbitals unless they have opposite spins.

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Spin-Spin Pairing & Electronic Configuration

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Exercise

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Ions
• Ions
How do ions form?

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Ions
•Positive and negative ions form when
electrons are transferred between atoms.

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Ions
•Positive and negative ions form when
electrons are transferred between atoms.

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Ions
• Some compounds are composed of
particles called ions.
– An ion is an atom or group of atoms that
has a positive or negative charge.
– A cation is an ion with a positive charge.
– An anion is an ion with a negative charge.

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Electron configuration of ions
When writing the electron
configuration of ions, it is
important to add or subtract
the appropriate number of
electrons.
For negative ions
add electrons.
Example: what is the electron structure of O2-?
8
1. Count number of electrons in atom
1s22s22p6
3. Fill sub-levels as for uncharged atom
8 + 2 = 10
2. Add or remove electrons due to
charge
For positive ions
remove electrons.
For non-transition metals, the
sub-levels are then filled as for
atoms.

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Electron configuration of ions
When writing the electron
configuration of ions, it is
important to add or subtract
the appropriate number of
electrons.
For negative ions
add electrons.
Example: what is the electron structure of Al3+?
13
1. Count number of electrons in atom
1s22s22p6
3. Fill sub-levels as for uncharged atom
13 - 3 = 10
2. Add or remove electrons due to
charge
For positive ions
remove electrons.
For non-transition metals, the
sub-levels are then filled as for
atoms.

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When transition metals form ions, it is the 4s electrons
that are removed before the 3d electrons.
Example: what is the electron structure of Ni2+?
28
1. Count number of
electrons in atom
1s22s22p63s23p64s23d8
2. Fill sub-levels,
remembering 4s is filled
before 3d
1s22s22p63s23p63d8
4. Remove electrons starting
with 4s
Electronic configuration of transition metal ions
2
3. Count number of
electrons to be removed

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What is ionization energy?
Ionization is a process in which atoms lose or gain
electrons and become ions.
The first ionization (I1) energy of an element is the energy
required to remove one electron from a gaseous atom.
M(g) → M+
(g) + e-
(g)
Looking at trends in ionization energies can reveal useful
evidence for the arrangement of electrons in atoms and
ions.
The second ionization (I2) energy involves the removal of
a second electron:
M+
(g) → M2+
(g) + e-
(g)

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Ionization energy definitions

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Plotting the successive ionization energies of
magnesium clearly shows the existence of different
energy levels, and the number of electrons at each
level.
Successive ionization
energies increase as more
electrons are removed.
Evidence for energy levels
Large jumps in the ionization
energy reveal where
electrons are being removed
from the next principal
energy level, such as
between the 2nd and 3rd, and
10th and 11th ionization
energies for magnesium.
electron removed
ionization
energy
2
3
4
5
6

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More evidence for energy levels
The first ionization energies of group 2 elements also
show evidence for the existence of different principal
energy levels.
Even though the nuclear
charge increases down
the group, the first
ionization energy
decreases.
element
first
ionization
energy
(kJ
mol
-1
)
500
600
700
800
900
Be Mg Ca Sr Ba
This means electrons
are being removed from
successively higher
energy levels, which lie
further from the nucleus
and are less attracted to
the nucleus.
400

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Trends in first ionization energies

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Factors Affecting the Energy of an Electron
• The orbital in which the electron exists
• The nuclear charge of an atom
• The repulsion (shielding) experienced by electron from
all the other electrons present

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Trends in Ionization Energy
• Size of the nuclear charge
The nuclear charge increases with increasing atomic
number, which means that there are greater attractive
forces between the nucleus and electrons, so more
energy is required to overcome these attractive forces
when removing an electron

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• Distance of outer electrons from the nucleus
Electrons in shells that are further away from the
nucleus are less attracted to the nucleus - the nuclear
attraction is weaker - so the further the outer electron
shell is from the nucleus, the lower the ionisation
energy

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• Shielding effect of inner electrons
The shielding effect is when the electrons in full inner
shells repel electrons in outer shells, preventing them
from feeling the full nuclear charge, so the more shells
an atom has, the greater the shielding effect, and the
lower the ionisation energy

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• Spin-pair repulsion
Electrons in the same atomic orbital in a subshell repel
each other more than electrons in different atomic
orbitals which makes it easier to remove an electron
(which is why the first ionization energy is always the
lowest)

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Group and Periodic Trends in Ionization Energy
• First ionization energy tends to increase from
bottom to top within a group and increase from left
to right across a period.
6.3

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6.3

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Ionisation energy across a period
The ionisation energy over a period increases due to the
following factors:
• Across a period the nuclear charge increases
• This causes the atomic radius of the atoms to decrease,
as the outer shell is pulled closer to the nucleus, so the
distance between the nucleus and the outer electrons
decreases
• The shielding by inner shell electrons remain reasonably
constant as electrons are being added to the same shell
• It becomes harder to remove an electron as you move
across a period; more energy is needed
• So, the ionisation energy increases

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There is a rapid decrease in ionisation energy between
the last element in one period, and the first element in
the next period because:
• There is increased distance between the nucleus and
the outer electrons as you have added a new shell
• There is increased shielding by inner electrons
because of the added shell
• These two factors outweigh the increased nuclear
charge

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There is a slight decrease in IE1 between beryllium and
boron as the fifth electron in boron is in the 2p
subshell, which is further away from the nucleus than
the 2s subshell of beryllium
• Beryllium has a first ionisation energy of 900 kJ mol-1
as its electron configuration is 1s2 2s2
• Boron has a first ionisation energy of 800 kJ mol-1 as
its electron configuration is 1s2 2s2 2px
1

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There is a slight decrease in IE1 between nitrogen and
oxygen and phosphorus due to spin-pair repulsion in the
2px orbital of oxygen
• Nitrogen has a first ionisation energy of 1400 kJ mol-1
as its electron configuration is 1s2 2s2 2px
1 2py
1 2pz
1
• Oxygen has a first ionisation energy of 1310 kJ mol-1 as
its electron configuration is 1s2 2s2 2px
2 2py
1 2pz
1

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Ionisation energy down a group
The ionisation energy down a group decreases due to the
following factors:
• The number of protons in the atom is increased, so the
nuclear charge increases
• But, the atomic radius of the atoms increases as you
are adding more shells of electrons, making the atoms
bigger
• So, the distance between the nucleus and outer electron
increases as you descend the group
• The shielding by inner shell electrons increases as
there are more shells of electrons
• These factors outweigh the increased nuclear charge,
meaning it becomes easier to remove the outer electron
as you descend a group
• So, the ionisation energy decreases

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Ionisation energy trends across a period
& going down a group table

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Trends in the Atomic Radius
• The atomic radius is the distance between the
nucleus and the outermost electron of an atom
• The atomic radius is measured by taking two atoms
of the same element, measuring the distance between
their nuclei and then halving this distance
• In metals this is also called the metallic radius and in
non-metals, the covalent radius

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Atomic radii of Period 3 elements table
• Across the period, the atomic radii decrease
• This is because the number of protons (the nuclear charge)
and the number of electrons increases by one every time
you go an element to the right
• The elements in a period all have the same number of shells
(so the shielding effect is the same)
• This means that as you go across the period the nucleus
attracts the electrons more strongly pulling them closer to
the nucleus
• Because of this, the atomic radius (and thus the size of the
atoms) decreases across the period

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• Ionic radius
• The ionic radius is the distance between the nucleus
and the outermost electron of an ion
• Metals produce positively charged ions (cations)
whereas nonmetals produce negatively charged ions
(anions)
• The cations have lost their valence electrons which
causes them to be much smaller than their parent
atoms
• Because there are less electrons, this also means
that there is less shielding of the outer electrons
• Going across the period from Na+ to Si4+ the ions get
smaller due to the increasing nuclear charge
attracting the outer electrons in the second principal
quantum shell nucleus (which has an increasing
atomic number)

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• Ionic radius
• The anions are larger than their original parent atoms
because each atom has gained one or more electrons
in their third principal quantum shell
• This increases the repulsion between electrons,
while the nuclear charge is still the same, causing the
electron cloud to spread out
• Going across P3- to Cl- the ionic radii decreases as
the nuclear charge increases across the period and
fewer electrons are gained by the atoms (P gains 3
electrons, S 2 electrons and Cl 1 electron)

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Ionic radii of ions of Period 3 elements
table

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a.The first ionisation energies of four consecutive elements in the
Periodic Table are:
sodium = 494 kJ mol−1 magnesium = 736 kJ mol−1
aluminium = 577 kJ mol−1 silicon = 786 kJ mol−1
i. Explain why aluminium has a lower first ionisation energy than
magnesium.
ii. Explain the general increase in ionisation energies from sodium
to silicon.
b. The first ionisation energy of fluorine is 1680 kJ mol−1 whereas
the first ionisation energy of iodine is 1010 kJ mol−1
Explain why fluorine has a higher first ionisation energy than iodine
despite fluorine having a smaller nuclear charge.
c. Explain why the values for the ionic radii for the negative ions in
Period 3 are higher than the ionic radii for the positive ions.
• SURPRISE!!!!

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a. i. The distance between the nucleus and the outer electrons
increases from Mg to Al. The shielding by inner shells increases.
These two factors outweigh the increased
nuclear charge.
ii. From sodium to silicon, the nuclear charge increases. The
distance between the nucleus and the outer electron remains
reasonably constant. The shielding by inner shells remains
reasonably constant. Ionisation energy increases to match
increase in attraction from the nucleus with an increased nuclear
charge.
• SURPRISE!!!! - Answer Key
b. The distance between the nucleus and the outer electron
increases from F to I. The shielding by inner shells increases.
These two factors outweigh the increased nuclear charge.
c. As there more electrons than a proton, the electrons repel
each other causing expansion in the size of an ion. Thus anion
has a larger size than its parent atom.

4Electron and Electron Configuration.pptx

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