Chem. 17 - Chpt. 4 Reading Notes

Front 1

What is the origin of lines in the atomic spectra?

Back 1

Electrons falling back to the ground state.

Front 2

Can you have a continuous range of energies in the atomic spectra?

Back 2

No, you can only have certain definite energies.

Front 3

What do the wave functions of Schrodinger's equation describe?

Back 3

The energy and position of the electrons, which are thought of as waves

Front 4

What are the four quantum numbers?

Back 4

n: principle quantum number

l: angular momentum quantum number

ml: magnetic quantum number

ms: spin angular momentum quantum number

Front 5

What does the principle quantum number represent?

Back 5

The atomic shells

Front 6

What does the angular momentum quantum number represent?

Back 6

The shape of the orbital

Front 7

What does the magnetic quantum number represent?

Back 7

The spatial orientation of the magnetic number - the direction along the x, y, and z axes

(in other words, there is only one sort of s-orbital, there are three types of p-orbital, there are five types of d-orbital, etc.)

Front 8

Construct a diagram showing the relationship of the quantum numbers.

Back 8

n = 1
l = s
ml = 0
ms = +0.5, -0.5

n= 2
l = s,p
ml = 0 for the s, 010 for the p
ms = +0.5, -0.5 for each

Front 9

What is the Pauli exclusion principle?

Back 9

No more than two electrons may occupy any orbital.

Front 10

What is Heisenberg's uncertainty principle?

Back 10

We cannot know both the exact position and the exact momentum of an electron simultaneously

Front 11

What is the nodal surface?

Back 11

Region between the 1s and the 2s orbital where there is no electron density at all

Front 12

What is the nodal plane?

Back 12

Region without any electron density between the 2p orbitals.

Front 13

Can an electron be found in the middle of hte p orbital?

Back 13

No, never

Front 14

Is the 1s orbital the middle part of the s2 orbital?

Back 14

No

Front 15

What is the Aufbau principle?

Back 15

Electrons are put in orbitals until they are full and then electrons fill up the orbital with the next lowest energy.

Front 16

Are the 2s and 2p orbitals degenerate?

Back 16

Yes

Front 17

What type of energy does is required to put two electrons into the same orbital with their spins paired?

Back 17

Pairing energy

Front 18

What are molecular orbitals?

Back 18

Combine the atomic orbitals of the atoms to make up the molecule

Front 19

What is the bonding molecular orbital?

Back 19

Combines orbitals in-phase to lower their energy relative to the 1s atomic orbital

Front 20

What is the antibonding molecular orbital?

Back 20

Combines electrons that are out-of-phase in the orbital, hindering the bonding and creating higher energy than the 1s orbital.

Front 21

What is bond order calculated?

Back 21

number of electrons in bonding molecular orbitals - number of electrons in antibonding molecular orbitals / 2

Front 22

How are sigma orbitals, (giving rise to sigma bonds), connected?

Back 22

End-on overlap to give cylindrical symmetry

Front 23

What does the side-on overlap of two 2p atomic orbitals give rise to?

Back 23

The 2p pie bonding molecular orbital

Front 24

Is more overlap possible when the atomic orbitals overlap end-on or side-on?

Back 24

End-on

Front 25

Is the p-sigma or the p-pie bond lower in energy?

Back 25

The p-sigma

Front 26

What is a single bond made of?

Back 26

A sigma bond

Front 27

What is a double bond made of?

Back 27

A sigma bond and a pie bond.

Front 28

What is a triple bond made of?

Back 28

A sigma bond and two pie bonds

Front 29

Is F or Li more electronegative?

Back 29

F

Front 30

If the energies of the atomic orbitals are both the same, what type of bond is formed?

Back 30

A purely covalent bond

Front 31

If the energies of the atomic orbitals are both the same, what is it easiest to break the bond into?

Back 31

Two radicals, by homolytic fission

Front 32

If the atomic orbital on atom Bi s a little lower in energy than the atomic orbital on atom A, then what type of bond is likely to form?

Back 32

A covalent bond with some electrostatic, (ionic), attraction

Front 33

If the atomic orbital on atom Bi s a little lower in energy than the atomic orbital on atom A, what is it easiest to break the bond into?

Back 33

Two ions, A+ and B-

Front 34

If the atomic orbital on atom B is a lot lower in energy than the atomic orbital on atom A, what type of bond is likely to form?

Back 34

Ionic bond

Front 35

If the atomic orbital on atom B is a lot lower in energy than the atomic orbital on atom A, what is it easiest to break the bond into?

Back 35

Compounds already exist as ions A+ and B-, so ions

Front 36

Do p-orbitals overlap better end-on, (sigma bond), or side-on, (pie bond)?

Back 36

End-on, forming a sigma bond

Front 37

Can 2px combine with any of the other 2p orbitals?

Back 37

No, because they are orthogonal

Front 38

Describe the different hybridizations of the carbons.

Back 38

Carbon is linear - sp hybridized

Carbon is trigonal - sp^2 hybridized

Carbon is tetrahedral - sp^3 hybridized

Front 39

What does isoelectronic mean?

Back 39

Total number of bonding electrons

Front 40

What is the hybridization of alkenes?

Back 40

sp^2, because they have a planar trigonal framework

Front 41

What is the hybridization of imines?

Back 41

sp^2, because they have a planar trigonal framework from the carbon atom as well as the nitrogen atom

Front 42

What type of hybridization do carbonyl compounds have?

Back 42

sp^2, because they have a planar trigonal framework

Front 43

What are the hybridizations of single, double, and triple bonds?

Back 43

Single: s

Double: sp^2

Triple: sp