Electron Configuration,
Periodicity, Bonding, and
Molecular Geometry
By XX
Bonding and Molecular
Geometry
Table of Contents
Section
Topics
1.
Key terms &
principles
Ÿ Ionic versus Covalent Bonding
Ÿ Electronegativity
Ÿ Polarity
2.
Octet Rule and the
Exceptions to the
rule
Ÿ Octet Rule
Ÿ Exception to the Octet Rule: Electron deficient molecules
Ÿ Exception to the Octet Rule: Expanded Octet
3.
Hybrid Lewis
Structures
Ÿ Lewis structures
Ÿ Formal Charge
4.
Resonance
structures
Ÿ Resonance structures
5.
Molecular geometry
Ÿ Molecular geometry of molecules
Ÿ Unshared electronic pairs
6.
Atomic orbitals
Ÿ Hybridization
Ÿ Sigma & Pi bonds
3
1. Key terms & principles
Ionic versus covalent bonding
Ionic Bond
Covalent Bond
Ÿ Electrons
transferred from
one atom to
another
Ÿ Electrons shared
between atoms
(single or
multiple)
Ÿ Opposite
charges hold the
molecule
together
Ÿ Sharing creates
strong bond
1. Key terms & principles
Electronegativity
Electronegativity:
tendency for atom to
attract a bonding pair
of electrons
1. Key terms & principles
Which is the most electronegative element?
?
a.
Fr (Atomic number 87)
b.
Cs (Atomic number 55)
c.
Br (Atomic number 35)
d.
F (Atomic number 9)
e.
C (Atomic number 6)
1. Key terms & principles
Polarity: electrons pulled towards one atom resulting in partial charge
H2O molecule: Electrons are “closer” to Oxygen
Opposite ends of molecules have
a slight positive or negative
charge because of an
asymmetrical distribution of
electrons
2. Octet Rule and the exceptions to the rule
Octet Rule: Atoms want to fill their valence shells
Octet rule: Atoms
want to gain, lose or
share electrons to
have 8 electrons in
valence shell (outer
shell) where
possible
Electrons needed to
complete valence shell
1 (no octet since
can only hold 2 )
4
3
2
2. Octet Rule and the exceptions to the rule
Exceptions to the Octet Rule: Electron deficient molecules
NO molecule with 11 valence electrons
Ÿ Some species do not follow
the octet rule such as when
there are odd number of
valence electrons (e.g.,
Nitrogen has odd number of
electrons in NO and NO2)
Ÿ Unpaired electrons are put
on other atoms to get
formal charge of zero
Octet is NOT filled –
Octet is filled – 8
7 electrons (needs 1
electrons
more)
BUT formal charge is 0 (remember for later)
2. Octet Rule and the exceptions to the rule
Exceptions to the Octet Rule: Expanded Octet
PCl5 Molecule
Ÿ When there are “too many”
electrons, some elements
can expand their octet and
hold more than 8 electrons
in their valence shell by
utilizing the d orbitals
Ÿ Elements with atomic
numbers greater than or
equal to that of Si are able
to have an expanded octets
(e.g., Si, P, S, Cl)
Ÿ
Ÿ
P has 5 bonds resulting in 10 electrons in the valence shell
Formal charge is 0 (remember for later)
3. Hybrid Lewis Structure
Hybrid Lewis Structures
Ÿ Diagram that show the
bonding between atoms of
a molecule
Ÿ Single lines represent
single bonds, two lines
represent double bonds,
etc.
Ÿ Dots represent lone pairs
and unshared electrons
3. Hybrid Lewis Structure
Formal Charge – helps you decide between multiple options
𝑪𝑶𝟐 Example
3. Hybrid Lewis Structure
Formal charge (FC) – an elegant technique to help you decide
Ÿ FC is the charge an
atom would have if
valence electrons
distributed evenly
Ÿ FC assigned to each
atom to help inform
how to balance
electrons across atoms
in given molecule (try
to get to “0” for each
atom)
Ÿ Sum across all atoms
to minimize formal
charge across atoms
(e.g., ions)
Formal Charge (FC) = [# of valence electrons on atom] – [non-bonded electrons + number of bonds]
3. Hybrid Lewis Structure
Steps to Follow for Drawing Lewis Dot Structure:
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the center
(except hydrogen)
3.
Draw the skeleton structure w/ single bonds
4.
Determine the remaining number of valence
electrons
5.
Apply octet rule
6.
Formal charge - balance center atom first
7.
Leftover electrons/ lone pairs go to the center
atom
8.
Beware of the number of electron clouds - lone
pairs count as electron clouds
3. Hybrid Lewis Structure
Hybrid Lewis Dot Structure: Practice Problem
Draw Hybrid
Lews Structure
for Methane
(CH4)
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
Ÿ CH4 has 8 valence electrons (from periodic
table)
‒ Carbon has 4
‒ Hydrogen has 1 each (4 total)
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
The Carbon is the center atom
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
3.
Draw the skeleton structure w/ single bonds
Skelton structure looks as follows:
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
3.
Draw the skeleton structure w/ single bonds
4.
Determine the remaining number of valence
electrons
There are no remaining valence electrons, all 8
are used
2 e-
2 e-
2 e2 e-
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
3.
Draw the skeleton structure w/ single bonds
4.
Determine the remaining number of valence
electrons
5.
Check to make sure that all atoms are abiding by
the octet rule (hydrogen only needs two electrons
for its valence shell):
2 e-
Apply octet rule
2 e-
2 e2 e-
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
3.
Draw the skeleton structure w/ single bonds
4.
Determine the remaining number of valence
electrons
5.
Apply octet rule
6.
Formal charge - balance center atom first
Ÿ Check to ensure atoms are
balanced via formal Charge
Ÿ Carbon has formal charge of 0
Ÿ Hydrogen has full valence shell of
2 electrons
2 e-
2 e-
2 e2 e-
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
3.
Draw the skeleton structure w/ single bonds
4.
Determine the remaining number of valence
electrons
5.
Apply octet rule
6.
Formal charge - balance center atom first
7.
There are no leftover electrons or lone pairs,
hence no need to put on center atom
2 e-
2 e-
2 e2 e-
Leftover electrons/ lone pairs go to the
center atom
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for Methane (CH4)
1.
Determine total number of valence electrons
2.
The less electronegative atoms are in the
center (except hydrogen)
3.
Draw the skeleton structure w/ single bonds
4.
Determine the remaining number of valence
electrons
5.
Apply octet rule
6.
Formal charge - balance center atom first
There are four electron clouds/areas of density, no
electron clouds from “lone” pairs
2 e-
2 e-
2 e2 e-
7.
Leftover electrons/ lone pairs go to the
center atom
8.
Beware of the number of electron clouds lone pairs count as electron clouds
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for NOCl (Nitrosyl chloride)
Use the steps to draw
the lewis structure for
NOCl
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for NOCl (Nitrosyl chloride)
Ÿ Nitrogen is the least
electronegative out of the three,
so you know that it will be in the
middle
Ÿ Ensure the octet rule (which
makes you realize that Nitrogen
needs double bond for octet)
Ÿ Check that formal charges are
correct
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for NO3- (Nitrate)
Use the steps to draw
the lewis structure for
NO3- (Nitrate)
3. Hybrid Lewis Structure
PRACTICE: Hybrid Lewis Dot Structure for NO3- (Nitrate)
Ÿ Nitrogen is the least
electronegative (will be in the
center) with Oxygen at the
terminals
Ÿ All atoms satisfy the octet
rule
Ÿ The formal charge is not
zero (it is an ion)
4. Resonance Structures
Drawing Resonance Structures
Approach for “resonance structures”
Ÿ First select one correct Lewis structure for
the molecule
?
What happens when there are
multiple Lewis Dot Structures
that are possible for the electron
configuration of a molecule?
Ÿ Make resonance structures off the selected
structure - Rotate through the different
positions the double, triple, etc
Ÿ Make sure that when you move the position
of the bonds that the molecule maintains its
original properties (formal charge,
hybridization etc.)
4. Resonance Structures
Drawing Resonance Structures
4. Resonance Structures
Drawing Resonance Structures
4. Resonance Structures
Drawing Resonance Structures
4. Resonance Structures
PRACTICE: Create resonance structure for two molecules
1. NCO-
NCO-
2. CO32-
4. Resonance Structures
PRACTICE: resonance structure for NCO-
NCO-
4. Resonance Structures
PRACTICE: resonance structure for NCO-
CO32-
5. Molecular geometry
Guideline for molecular geometry of molecules
Predicting the shape of molecules
Ÿ Once the Lewis structure has been
completed, the key is to arrange the
electron pairs around the central atom in
way that minimizes repulsion (as far part
as possible)
Ÿ The electronic repulsion means the
electron configuration outcome may be
different than the molecular geometry of
the molecule
Ÿ The table on the left provides a guide for
what the molecular geometry looks like
5. Molecular geometry
Unshared pairs determine molecular geometry
Ÿ Even though all three
molecules have 4 electron
clouds (electronic
configuration), the
molecular geometry is
different
Ÿ The lone pair of electrons
repel bonded pairs (and
hence become closer) as
lone pairs are more
“spread out”
Bond angle
(degrees)
109.5
<109.5
(slightly less)
105
6. Atomic orbitals
Sigma and Pi bonds
Ÿ Single bonds are sigma bonds,
which means their atomic orbitals
overlap during the covalent bonding
process (strong)
Ÿ The remaining bonds, are pi bonds
because of unhybridized p orbitals
above and below the sigma bond
Ÿ NOTE: one pi bond and consists of
two lobes (one lobe above and one
below the sigma bond)
6. Atomic orbitals
Sigma and Pi bonds C3H4 (Propyne)
Ÿ The two triple bonded
Carbons have 1 sigma
bond and 2 pi bonds
Ÿ Total of 6 sigma bonds
and 2 pi bonds
6. Atomic orbitals
PRACTICE: Determine sigma and pi bonds
?
How many
sigma bonds
and pi bonds
are there?
6. Atomic orbitals
PRACTICE: Determine sigma and pi bonds
?
3 sigma bonds
and 1 pi bond
6. Atomic orbitals
Hybridization of Oribitals for bonding
Ÿ Orbitals mix into new hybrid
orbitals (with different energies,
shapes, etc., than component
orbitals), to allow electrons to pair
and form bonds (sigma bonds)
Ÿ Example: carbon atom which forms
four single bonds the valence-shell s
orbital combines with three valenceshell p orbitals to form four equivalent
sp3 mixtures which form tetrahedral
Ÿ Remaining bonds, are pi bonds
because of unhybridized p orbitals
above and below the sigma bond
6. Atomic orbitals
Hybridization of Oribitals for bonding: oribital table and geometries
Ÿ Hybrid orbitals are
based on number of
atomic orbitals that
are mixed
Ÿ Reminder: Hybrid
orbitals are for sigma
bonds and unshared
pairs; extra bonds
are pi bonds
6. Atomic orbitals
PRACTICE: Hybridization of Oribitals
Question: What is the hybridization of nitrogen in following molecules?:
NO2-
NH3
N2
6. Atomic orbitals
PRACTICE: Hybridization of Oribitals
Question: What is the hybridization of nitrogen in following molecules?:
NO2-
NH3
sp2 (one unshared
pair, 1 single bond, 1
of the electron pairs in
the double bond)
N2
sp3 (one unshared
pair, 3 single bonds)
sp (one unshared pair,
, 1 of the electron pairs
in the double bond)
Electron Configuration
and Periodicity
Table of Contents
Section
Topics
1.
Quantum model of
electrons
Ÿ Quantum Numbers
Ÿ Orbitals (s, p, d, f)
Ÿ Periodic element “blocks”
2.
Electron
configurations
Ÿ Aufbau Process
Ÿ Hunds Rule
Ÿ Gas Core configuration
3.
Isoelectric species
Ÿ Isoelectric species examples
4.
Periodicity
Ÿ Atomic size
Ÿ Ionization energy (first and subsequent)
46
1. Quantum model of electrons
Electron behavior is more complicated than we think
Bohr model (Classical
Mechanics)
Electrons are “particles”
orbit the nucleus
Electron Cloud Model
(Quantum Mechanics)
Electrons are “waves” that
exist in certain “cloud” or
“orbital” patterns
1. Quantum model of electrons
Four quantum numbers describe electron behavior
1
2
4
MAIN energy
level of the
electrons –
higher n means
electrons have
more energy;
eletcrons are
further away
from the
nucleus
The ℓ value (Azimuthal quantum
number) determines the SHAPE of
the orbital (e.g., s, p, d, f)
3
“SPIN” of
the electrons
– clockwise
or counter
clockwise
Each ℓ level has sub-orbitals mℓ which
determines the direction or ORIENTATION
in space of the electron cloud (e.g., 1, 3, 5,
7 sub orbitals)
1. Quantum model of electrons
The s Orbital (ℓ quantum number of 0)
`
`
`
`
The ℓ value or Azimuthal number of
0 means it is an s orbital and its
SHAPE is a “sphere”
`
`
`
`
As quantum
number (n)
increases from 1
to 3, the energy
level goes up –
electrons further
away from the
nucleus
1. Quantum model of electrons
The p Orbital (ℓ quantum number of 1)
Ÿ Quantum number ℓ of
1 means p orbitals and
its SHAPE is a “bow
tie”
Ÿ mℓ quantum number
of -1, 0, +1 means
three sub-orbitals with
its ORIENTATION
along three axes
Ÿ Each sub-orbital holds
2 electrons (6 total)
1. Quantum model of electrons
The d Orbital (ℓ quantum number of 2)
Ÿ Quantum number ℓ of 2
means d orbitals its SHAPE
is a “clover”
Ÿ mℓ quantum number of -2, -1,
0, +1, +2 means five suborbitals with its
ORIENTATION as shown
here
Ÿ Five sub-orbitals each holding
2 electrons (10 total)
1. Quantum model of electrons
f Orbital (ℓ quantum number
of 3) with seven sub-orbitals
Ÿ Quantum number ℓ of 3
means d orbitals its SHAPE
is a “tetrahedral”
Ÿ mℓ quantum number of -3, 2, -1, 0, +1, +2 , +3 seven
sub-orbitals with its
ORIENTATION as shown
Ÿ Seven sub-orbitals each
holding 2 electrons (14 total)
Ÿ High energy – usually
unstable
1. Quantum model of electrons
Periodic table “blocks” guide which valence electron orbitals are
important for reactions
2. Electron configurations
Aufbau Process to determine
electron configuration
Ÿ Aufbau Process helps determine the
electron configuration of an atom keeping electrons at their lowest
possible energy (close as possible
to the nucleus)
Ÿ Algorithm for allocating electrons:
‒ Determine number of electrons
present (atomic number)
‒ Fill lowest energy orbitals first and
work way up as detailed in exhibit)
2. Electron configurations
Hund’s rule of maximum municipality
Nitrogen example – Atomic number (7)
Ÿ Electrons are placed into each
orbital before any pairing takes
place (e.g., three 2p electrons go
into each sub-orbital versus
pairing)
Ÿ Definitions for solid forms:
‒ Paramagnetic: created when
unpaired electrons are present
and attracted to magnetic field
‒ Diamagnetic: occurs when all
electrons are paired and slightly
repelled by magnetic field
2. Electron configurations
PRACTICE: problems on Orbital Diagrams
Hint: Make sure
you know the
number of
electrons you
have based on
periodic table
2. Electron configurations
PRACTICE: problems on Orbital Diagrams
2. Electron configurations
Writing gas core diagrams
Ÿ Noble gases have full octets and
therefore can be used as reference
points to show orbital diagrams
for subsequent elements – avoids
writing long orbital notations
Ÿ Example: Na will has 1 more
electron than Ne, hence you can
write as [Ne]3s1; the one extra
electron goes into next orbital
2. Electron configurations
PRACTICE: Gas core diagram for Ca
Draw the gas core
diagram for calcium
(hint: what is closest
noble gas)
2. Electron configurations
PRACTICE: Gas core diagram for Ca
3. Isoelectric species
Isoelectronic species
Examples of Isoeletronic species
Isoelectronic
species have
same electronic
number and
configuration
4. Periodicity
Periodicity
Key Characteristic
Atomic radius
Periodicity refers to
trends or recurring
variations in element
properties with their
placement on the
periodic table (e.g.,
atomic number, group,
period)
Ionic radius
(not covered in
exhibit)
Ionization energy
Description
Ÿ The size of the atom – the distance between center of
atoms of an element
Ÿ Smaller the size, the closer and more stable are the
electrons (closer to the nucleus)
Ÿ The size of the radius of the ions
Ÿ Cations have smaller radii (less electrons, so
remaining are closer to nucelli)
Ÿ Anions have larger radii since more electrons
avoiding each other (further from nucleus)
Ÿ Energy needed to unlock electron from gaseous atom
Ÿ The closer the electrons to the nuclei (lower energy
states), the more energy is needed to dislodge
7/31/23
Electronegativity
(covered in
previous section)
Ÿ Measure of ability of atom to attract (pull in) electrons
during covalent bodning
62
4. Periodicity
Atomic size: Electrons closer to the nucleus (stronger pull) have smaller
atomic size and hence are more stable
The n value
increases as you
move down each
period, and hence
electrons are further
away from the
nucleus (higher
atomic radius)
4. Periodicity
Ionization energy
Ÿ Amount of energy needed to
lose the most loosely bound
e- of an atom in its gaseous
state
Ÿ “Zig-zag” effect: as you
move further right (to noble
gas) higher ionization
because valence shells more
filled, move down less energy
as electrons further out
4. Periodicity
Successive ionization energies are MUCH higher
Example of Aluminum Ionization Energy (Kj/mol)
Ÿ The energy needed to
remove subsequent
electrons (closer to the
nucleus) jumps a LOT
Ÿ Note that the Y axis is
logarithmic which
means that to remove
the 12th electron
requires almost 350X
the energy as the first
electron
~350X