Flashcards in Chapter 10 Deck (29):

1

## VESPR

### valence shell electron pair repulsion theory - repulsion between electron groups on interior atoms determines the geometry of the molecule; preferred geometry is one where electron groups have max separation

2

## effects on bond angles

### different types of electron groups exert slightly different repulsions; ex: double bond exerts more repulsion than single

3

## electron group

### can be lone pair, single electron, single bond, double, or triple. all = 1 group

4

## electron group repulsions

### lone pair - lone pair > lone pair - bonding pair > bonding pair - bonding pair

5

## what determines geometry of molecule

### number of electron groups on central atom or all interior atoms

6

## what determines # of electron groups

### Lewis structure - if more than 1 resonance structure, any can be used

7

## what determines geometry of electron groups

### minimizing their repulsions

8

## why do bond angles vary from ideal?

### double and triple bonds occupy more space than single and lone pair occupy more space than bonding groups

9

## symbols for drawing mol. geo. on paper

###
straight line = bond in plane of paper

hatched wedge = bond going into paper

solid wedge = hatch coming out of paper

10

## predicting shapes of larger molecules

### find geometries of each atom and add together in 1 molecule

11

## Why can dipole moments cancel each other

### They are vector quantities with a magnitude and direction. They can cancel if they directly oppose each other

12

## Process to determine if a molecule is polar

###
1. Draw Lewis structure and determine molecular geometry

2. Determine if it contains polar bonds

3. Determine if polar bonds add together to form a net dipole moment

13

## Why are polar molecules attracted to other polar molecules

### They have positive and negative ends and the polar end of one is attracted to the negative end of the other

14

## Like dissolves like

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Polar molecules will mix together and non polar molecules will mix together.

Polar and non polar will not mix together

15

## Valence bond theory

###
Electrons reside in quantum mechanical orbitals localized on individual atoms

Usually this is s p d f but could also be hybridized atomic orbitals

16

## When does a chemical bond form

### When two atoms approach each other the electrons and nucleus of one interact with the other and vice versa. If energy of system is lowered from the interactions a bond forms. If it's raised a bond will not form

17

## Equilibrium bond length

### If graphed, it's the lowest point which represents the lowest energy

18

## Summary of valence bond theory - 3 points

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1. Valence electrons lie in atomic orbitals which are s p d f or hybrids

2. Chemical bonds result from the overlap of two half filled orbitals with spin pairing of the two valence electrons or the overlap of a filled orbitals with an empty one.

3. The geometry of the overlapping orbitals determines the shape of the molecule.

19

## how does the concept of hybridization move valence bond theory forward

### orbitals in a molecule are not necessarily the same as the orbitals in an atom

20

## hybridization

### mathematical procedure that allows us to combine standard atomic orbitals to form new atomic orbitals called hybrid orbitals that correspond more closely to actual distribution of electrons in chemically bonded atoms

21

## compare/contrast standard atomic orbitals and hybrid orbitals

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both are localized on individual atoms, but have different shapes and energies

In valence bond theory a chemical bond is the overlap of two orbitals that together contain two electrons. In hybrid orbitals the electron probability density is more concentrated in a single directional lobe allowing greater overlap with orbitals of other atoms

22

## Why do molecules have hybrid orbitals

### Hybrid orbitals maximize orbital overlap in a bond and minimize energy of the molecule

23

## Relationship between number of bonds and hybridization

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The more bonds an atoms forms the greater the tendency to hybridize

Carbon always hybridizes because it tends to form four bonds in its compounds

24

## Rules for hybridization - 3 points

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1. Number of standard atomic orbitals added together always equals number of hybrid orbitals formed. Total number of orbitals is always conserved

2. Particular combinations of standard orbitals added together will determine the shape and energy of hybrid orbitals

3. Type of hybridization that occurs is the one that yields the lowest overall energy. Use electron geometries as determined by vsepr theory to predict type of hybridization

25

## Sp^3 hybridization

###
Mixture of 1 s plus 3 p orbitals

All hybrids orbitals have the same energy so there degenerate

Tetrahedral geometry with angles if 109.5 degrees which is the same as the resulting molecule geometry

Presence of a lone pair lowers the tendency of nitrogens orgitals to hybridize so 3 hybrids are involved with bonding with hydrogen atoms and the fourth contains a lone pair. This moves the bond angle to 107 degrees which is closer to the unhybridized p orbital

26

## Sp^2 hybrid orbital

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Hybridization of one s and two p orbitals results in 3 sp^2 hybrids and one unhybridized p orbital

3 hybrids orbitals are in trigonal planar geometry with 120 degree bond angles and the p orbital is perpendicular to all.

27

## Pi vs sigma bonds

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Pi bond is when P orbitals overlap side by side. The electron density is above and below the internuclear axis.

Sigma bond is when orbitals overlap end to end

Double bonds in Lewis structure always correspond to one pi and one sigma bond in valence bond theory

28

## Rotation around a bond

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Rotation around a double bond is highly restricted but not restricted around a single bond isn't

Restricted rotation can result in two forms with difrfor one molecule.

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