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Chem 150 > Module 5: Exam 3 > Flashcards

Module 5: Exam 3 Flashcards

(58 cards)

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1
Q

2 electron groups, 0 lone pairs

A

linear

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2
Q

3 electron groups, 0 lone pairs

A

trigonal planar

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3
Q

3 electron groups, 1 lone pair

A

bent

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4
Q

4 electron groups, 0 lone pairs

A

tetrahedral

2 atoms in opposite planes

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5
Q

4 electron groups, 1 lone pair

A

trigonal pyramidal

2 atoms in opposite planes

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6
Q

4 electron groups, 2 lone pairs

A

bent

atoms are not in opposite planes, technically the lone pairs are, but we do not show this

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7
Q

5 electron groups, 0 lone pairs

A

trigonal bypyramidal

2 atoms in opposite planes

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8
Q

6 electron groups, 0 lone pairs

A

octahedral

4 atoms in opposite planes

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9
Q

Linear ideal angle

A

180

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10
Q

Trigonal planar ideal angle

A

120

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11
Q

Tetrahedral ideal angle

A

109.5

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12
Q

Trigonal bypyramidal ideal angle

A

Distance between equatorial: 120
Distance between axial and equatorial: 90

*drawing helps understand this

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13
Q

What are molecular geometries based on?

A

Electron electron repulsion in ALL electron groups (bonded and lone pairs)

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14
Q

Octahedral ideal angle

A

90

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15
Q

Why do lone pairs squish bonds together?

A

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

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16
Q

What repels more single bonds or double bonds?

A

Double bonds since they have greater electron density

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17
Q

When is it okay for a 1st period element to have a 2- or 2+ formal charge?

A

When it is in a resonance structure

When just drawing Lewis structures this is NOT allowed

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18
Q

Do resonance structures truly exist?

A

No

They are just representations

All bonds are the same length between the same atoms surrounding a central atom

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19
Q

Why is hybridization and resonance structures favored/

A

They are often more stable than an individual structure that does not exist as a resonance molecule

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20
Q

What defined electron group geometry?

A

The number of areas of electron density surrounding the central atom

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21
Q

How do hybrids add up?

A

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

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22
Q

sp3

A

commonly used with 4 electron domains (tetrahedral)

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23
Q

sp2

A

commonly used with 3 electron domains (trigonal planar)

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24
Q

Where does resonance happen?

25
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
26
Orientation of pi-bonds
perpendicular because they are the leftover p-orbitals
27
What constitutes an electron domain?
a lone pair = 1 a single bond = 1 a double bond = 1 a triple bond = 1
28
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
29
Valence bond theory
atoms share electrons when an atomic orbital on one atom overlaps with an atomic orbital on the other
30
What holds electrons together in covalent bond?
mutual attraction to the nuclei
31
What happens when 2 single bonded orbitals overlap?
they both become doubly occuppied
32
Which theory explains why a bond forms?
valence bond theory
33
When will a covalent bond form?
the potential energy of the resulting molecule is lower than the potential energy of 2 separate atoms
34
What happens energetically when bonds are formed?
energy is released | force of attraction increasing, energy dropping
35
Hybridization
the mixing of atomic orbitals an extension of valence bond theory for when valence bond theory fails
36
What does an sp orbital look like?
large lobe and small lobe 50% s-characteristics 50% p-characteristics
37
sp
commonly used with 2 electron domains (linear)
38
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
39
sp3d2
commonly used with 6 electron domains (octahedral)
40
Unique about the 3rd period and beyond in terms of hybridization
electrons can be promoted to the d-orbital octets can be overloaded
41
3 steps to evaluate the importance of resonance structures
1. determine if all the atoms have a complete octet. full octet>not full 2. count # of atoms with formal charges. less formal charges=greater importance 3. if same # of atoms with formal charges, see where they lie. more electronegative atoms want the negative charge
42
What happens to ionization energy as you go across p table?
increases
43
What happens to ionization energy as you go down p table?
decreases
44
What happens to atomic radius as you go accross?
decreases
45
What happens to atomic radius as you go down?
increases
46
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
47
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
48
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.
49
What is stronger delocalized bond or localized?
delocalized
50
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
51
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
52
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
53
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
54
What does VSEPR stand for?
valence shell electron pair repulsion
55
Who invented vsepr?
Gillespie and Nyholm
56
What are the limitations of vsepr?
fails to take into account differences in isoelectric species fails for transition metals
57
Formula for bond order
number of bonds / number of atoms sharing them
58
Relationship between bond order and bond length
as bond order increases, bond length decreases

Chem 150 (10 decks)

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