CHEM 110 Chapter Notes - Chapter 10: Octet Rule, Orbital Hybridisation, Chemical Polarity
10 May 2018
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Chapter 10
- VSEPR Theory:
o Valence shell electron pair repulsion theory
▪ Based on the idea that electron groups which we defined as lone pairs,
single bonds, multiple bonds repel one another through coulombic forces
▪ Electron group are attracted to the nucleus but VSEPR focuses on the
repulsions
▪ Repulsions between electron groups on interior atoms of a molecule
determine the geometry of the molecule
▪ For molecules having one atom, molecular geometry depends on the
number of electron groups around the central atom
• How many of those electron groups are bonding groups and how
many are lone pairs
• Constitute the five basic shapes of the molecules
o Two electron Groups:
▪ Consider the Lewis Structure of BeCl2
• Two electron groups, and two single
bonds about central atom
• The geometry of BeCl2 is determined by the repulsion between the
two electron groups
o Maximize their separation by assuming a 180 bond angle or
linear geometry
o Experimental measurements of geometry of BeCl2 indivate
the molecule is linear
• Molecules that only form two single bonds, are rare because they
do not follow octet rule
o Consider CO2 which has two electron groups around the
central carbon atom
o Double bond counts as one electron group
• Two double bonds repel each other, resulting in a linear geometry
for CO2
o Three Electron Groups
▪ BF3 has three electron groups around the central atom
• Maximize their separation by assuming 120
bond angles in a plane
o Trigonal planar geometry
• Experimental observations of the structure are in agreement with
predictions of VSEPR theory
o Another molecule with three electron groups, has one
double bond and two single bonds
o Four Electron Groups
▪ For molecules with four or more electron groups around the central atoms,
the geometries are three-dimensional and are more difficult to imagine and
draw
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▪ If you tie four balloons together, however, they assume a three-
dimensional tetrahedral geometry
▪ Verticals pointed form a tetrahedron
o Five electron Groups
▪ Assume a trigonal bipyramidal geometry
▪ Three of the groups lie in a single plane, in the trigonal planar
configuration
▪ The angles between equatorial positions are 120, while angle between
axial and trigonal are 90
o Six electron Groups
▪ Assume a octahedral geometry
▪ Four of the groups lie in a single plane, with fifth group above the plane
and another below it
- VSEPR Theory: Effect of Lone Pairs
o Four Electrons with Lone Pairs
▪ The central nitrogen atom has four electron groups
• Repel one another
• If we do not distinguish between the bonding electron groups and
the lone pairs, we find the electron geometry, the arrangement of
the electron groups is still tetrahedral
• The electron geometry is relevant to the molecular geometry
• Since it has four electron groups, its electron geometry is also
tetrahedral, but molecular geometry is bent
• Lone pairs compress the H2O bond angle to an even graeter extent
than in NH3
o Lone Pair > Lone-Bonding pair > Bonding Pair
o Five electron with Lone Pairs
▪ The electron geometry due to the five electron groups is trigonal
bipyramidal
▪ The line pair should occupy the position that minizes the interaction with
the bonding pairs
▪ Lone pair occupies an equatorial position
• Resulting geometry is seesaw
o Called an irregular tetrahedron
▪ When two of the five electron groups around the central atom are lone
pairs, the lone pairs occupy two of the three equatorial positions
• Resulting geometry is T-Shaped
o Six electron groups with Lone Pairs
▪ The electron geometry due to the six groups is octahedral
▪ All six positions are equivalent and the lone pair can be situated in any of
the positions
• Square Pyramidal
▪ When two of the six electron groups around the central atom are lone
pairs, the resulting geometry is Square Plannar
o The geometry is determined by the number of electron groups on the central atom
o The number of electron groups can be determined from lewis structure
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▪ If the Lewis Structure contains resonance structures, use any one of the
resonance to determine the number of electron groups
o Each of the following counts as single electron group
▪ Lone pair, single bond, double bond, triple bond
o Bond angles can vary from idealized angles because double and triple bonds
occupy more space than single bonds
▪ Lone pairs occupy more space than bonding groups
▪ Presence of lone pairs will make bond angles smaller than ideal angle for
particular geometry
- VSEPR Theory: Predicting molecular geometries
o Step 1: Draw a lewis structure
o Step 2: Determine the total number of electron groups
o Step 3: Determine the number of bonding groups and number of lone pairs around
the atom
o Step 4: Determine the electron and molecular geometry
▪ Large molecules have two or more interior atoms
- Molecular Shape and Polarity
o Molecules can eb polar depending on their shape and nature of their bonds
▪ In an electron density model,
• Yellow indicates high electron density
• red indicates very high electron density
• Blue indicates low electron density
o The presence of polar bonds may or may not result in a polar molecule
▪ If the molecular geometry is such that dipole moments of individual polar
bonds sum together to a net moment, than it will be polar
• Example: CO2
o CO2 is polar because oxygen and carbon have different
electronegativities
o Since CO2 is linear molecule, the polar bonds oppose on
another and dipole moment of one bond exactly opposes
the dipole of another
▪ Sum to zero and the molecule is nonpolar
▪ Dipole can cancel
eachother because
they are vector
quantities
o In CO2 we have two vectors
pointing in the opposite directions
▪ Yellow shows regions of moderately high electron
density positioned symmetrically on either end of
the molecule with a region of low electron density
(blue) in the middle
• Example H2O
o O—H bonds are polar
o Oxygen and hydrogen have electronegativities of 3.5 and
2.1
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Document Summary
Four electrons with lone pairs: the central nitrogen atom has four electron groups, repel one another. Molecular shape and polarity: molecules can eb polar depending on their shape and nature of their bonds. In an electron density model: yellow indicates high electron density red indicates very high electron density, blue indicates low electron density, the presence of polar bonds may or may not result in a polar molecule. If vector sum to zero (nonpolar) if vector sum to net vector (polar) Valence bond theory: the simpler of the two bonding theories is valencce bond theory. In valence bond theory, electrons reside in orbitals localized on individual atoms. In some cases, the orbitals are hybridized atomic orbitals: an application of a more general quantum-mechanical method called. In valence bond theory, we calculate the effect of these interaction on the energies of the electrons in the atomic orbitals. If energy is lowered because of interactions, a bond is formed.
