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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

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

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

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

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

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

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

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

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.

29

Sp hybridization

Hybridization of one s and one p orbital results in two sp hybrids and two leftover unhybridized p orbitals

Hybrid orbitals are linear with 180 degree bond angles. Unhybridized p orbitals are oriented in a plane perpendicular

Triple bond consist of one sigma bond (overlapping sp orbitals) and two pi bonds (overlapping p orbitals)