Module 5: Exam 3 Flashcards
2 electron groups, 0 lone pairs
linear
3 electron groups, 0 lone pairs
trigonal planar
3 electron groups, 1 lone pair
bent
4 electron groups, 0 lone pairs
tetrahedral
2 atoms in opposite planes
4 electron groups, 1 lone pair
trigonal pyramidal
2 atoms in opposite planes
4 electron groups, 2 lone pairs
bent
atoms are not in opposite planes, technically the lone pairs are, but we do not show this
5 electron groups, 0 lone pairs
trigonal bypyramidal
2 atoms in opposite planes
6 electron groups, 0 lone pairs
octahedral
4 atoms in opposite planes
Linear ideal angle
180
Trigonal planar ideal angle
120
Tetrahedral ideal angle
109.5
Trigonal bypyramidal ideal angle
Distance between equatorial: 120
Distance between axial and equatorial: 90
*drawing helps understand this
What are molecular geometries based on?
Electron electron repulsion in ALL electron groups (bonded and lone pairs)
Octahedral ideal angle
90
Why do lone pairs squish bonds together?
Because they only are attracted to 1 nucleus, instead of 2, this gives them less force of attraction
Therefore, they have more space to spread out and repel other domains
What repels more single bonds or double bonds?
Double bonds since they have greater electron density
When is it okay for a 1st period element to have a 2- or 2+ formal charge?
When it is in a resonance structure
When just drawing Lewis structures this is NOT allowed
Do resonance structures truly exist?
No
They are just representations
All bonds are the same length between the same atoms surrounding a central atom
Why is hybridization and resonance structures favored/
They are often more stable than an individual structure that does not exist as a resonance molecule
What defined electron group geometry?
The number of areas of electron density surrounding the central atom
How do hybrids add up?
If there is 2 unhybridized orbitals, they will come together to form 2 hybridized orbitals
This is due to the law of conservation of energy
sp3
commonly used with 4 electron domains (tetrahedral)
sp2
commonly used with 3 electron domains (trigonal planar)
Where does resonance happen?
pi-bonds
When there is a leftover p-orbital, where is it?
remember that p-orbitals can be in 3 planes. and if you are using two of them for a sp2 orbital, then the leftover is in the z-plane
this makes it perpendicular
Orientation of pi-bonds
perpendicular
because they are the leftover p-orbitals
What constitutes an electron domain?
a lone pair = 1
a single bond = 1
a double bond = 1
a triple bond = 1
Why do electron domains like to be as far apart as possible?
Since electrons repel eachother, want to be far apart to minimize force of repulsion and lower energy
Valence bond theory
atoms share electrons when an atomic orbital on one atom overlaps with an atomic orbital on the other
What holds electrons together in covalent bond?
mutual attraction to the nuclei
What happens when 2 single bonded orbitals overlap?
they both become doubly occuppied
Which theory explains why a bond forms?
valence bond theory
When will a covalent bond form?
the potential energy of the resulting molecule is lower than the potential energy of 2 separate atoms
What happens energetically when bonds are formed?
energy is released
force of attraction increasing, energy dropping
Hybridization
the mixing of atomic orbitals
an extension of valence bond theory for when valence bond theory fails
What does an sp orbital look like?
large lobe and small lobe
50% s-characteristics
50% p-characteristics
sp
commonly used with 2 electron domains (linear)
What do we use hybrid orbitals for?
not to predict geometries, that is vsepr, rather we use them to EXPLAIN geometries that we already know
sp3d2
commonly used with 6 electron domains (octahedral)
Unique about the 3rd period and beyond in terms of hybridization
electrons can be promoted to the d-orbital
octets can be overloaded
3 steps to evaluate the importance of resonance structures
- determine if all the atoms have a complete octet. full octet>not full
- count # of atoms with formal charges. less formal charges=greater importance
- if same # of atoms with formal charges, see where they lie. more electronegative atoms want the negative charge
What happens to ionization energy as you go across p table?
increases
What happens to ionization energy as you go down p table?
decreases
What happens to atomic radius as you go accross?
decreases
What happens to atomic radius as you go down?
increases
Delocalized bond
pi-bond is spread out over the entire molecule
double bonds are not in one location because each bond surrounding a central atom connected to the same other atom, has the same length
therefore, bonds are not in one spot. they are constantly changing
Are pi-bonds stronger than sigma?
No. Sigma bonds are stronger because they involve the most direct overlap
Pi bonds are more of a cloud over bonds
Why is a double bond not twice the strength of a single bond?
Because it is not two sigma bonds
it is a sigma bond and a pi bond. the pi bond is not as strong.
What is stronger delocalized bond or localized?
delocalized
Why are delocalized bonds stronger?
they increase stability through resonance
there is a loss of internal energy which makes more stable
since electrons repel eachother, by distributing them further apart over multiple bonds, energy is reduced
Why is cyclic geometry less stable in a triangle?
Forces electrons into 60º angles which makes them very close
HOWEVER, it depends the shape. An aromatic ring like benzene it very stable because angle is still 120º
Which is more stable sp2 or sp3? Why?
sp2 because it has a greater percentage of s-character
sigma bonds are strong because they are direct overlap of orbitals
What is a possible reason the bond angle might not be predicted by VSEPR?
it could be due to a resonance structure that gives a geometry that is better predicted by vsepr
What does VSEPR stand for?
valence shell electron pair repulsion
Who invented vsepr?
Gillespie and Nyholm
What are the limitations of vsepr?
fails to take into account differences in isoelectric species
fails for transition metals
Formula for bond order
number of bonds / number of atoms sharing them
Relationship between bond order and bond length
as bond order increases, bond length decreases
