ch9 molecular geometry and bonding theory

Question 1

molecular architecture

Answer 1

Lewis structures do not show the overall shapes of molecules (molecular architecture), they only show number and types of bonds between atoms

Question 2

bond angles

Answer 2

The shape of a molecule is determined by its bond angles: angles made by the lines joining the nuclei of atoms in the molecules. Bond angles and bond lengths together accurately define the shape/size of a molecule

Question 3

5 geometric structures

Answer 3

for ABn (A is bonded to n B atoms)
linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral
When A is a representative element (from s or p block) we can predict shape of molecule using valence-shell electron-pair repulsion (VSEPR model)

Question 4

bonding pair

Answer 4

a region in which the electrons will most likely be found

Question 5

nonbonding/lone pair

Answer 5

electron domain that is located principally on one atom

Question 6

electron domain

Answer 6

each nonbonding pair, single bond, or multiple bond produces an electron domain around the central atom. The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them (basis of Valence Shell Eelectron Pair Repulsion model)

Question 7

electron domain geometry2

Answer 7

• arrangement of electron domains about the central atom of ABn
• The different shapes of ABn molecules/ions depend on the number of electron domains surrounding central A atom.

Question 8

molecular geometry

Answer 8

arrangement of only the atoms in the molecule/ion, does not describe nonbonding pairs of electrons

Question 9

ideal bond angles

Answer 9

angles for specific electron geometry when surrounding electrons and domains are identical

Question 10

factors for greater bond angle

Answer 10

Electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and thus tend to compress bond angles. Electron domains for multiple bonds exert a greater repulsive force on adjacent electron domains than do electron domains for single bonds (more nonbonding electrons or double/triple bonds=greater angle for that electron domain)

Question 11

molecules with 5 domains

Answer 11

• electron domain geometry: trigonal bipyramidal (2 trigonal pyramids sharing a base)
• (0 lone pairs) 2 of 5 domains: axial positions, makes 90º with equatorial bonds
• (0lone pairs) Other 3 domains: equatorial positions, makes 90º with axial bonds, 120º with other equatorial bonds, nonbonding pairs always occupy this position
• 1 lone pair: see saw, <120, <90
• 2lone pair: tshaped, <90
• 3 lone pair: linear, 180

Question 12

molecules with 3 domains

Answer 12

• 0 lone pair: trigonal planar, 120

- 1 lone pair: bent, <120

Question 13

4 electron domains

Answer 13

• 0 lone pair: tetrahedral, 109.5
• 1 lone pair: trigonal pyramidal, <109.5
• 2 lone pair: bent, «109.5

Question 14

2 electron domains

Answer 14

linear, 180

Question 15

6 electron domains

Answer 15

• electron domain geometry: octahedron (2 square pyramids sharing a base)
• 0 lone pairs: octahedron, 90
• 1 lone pair: square pyramidal, <90
• 2 lone pair: square planar, 90

Question 16

bond dipole

Answer 16

the dipole moment that is due only to the two atoms in that bond. Bond dipoles/dipole moments are vector quantities: they have magnitude and direction

Question 17

polar vs nonpolar electron geometry

Answer 17

if bond dipoles dont cancel out (if lone pairs), then polar.

Question 18

valence-bond theory

Answer 18

ORBITALS OVERLAP

• model of chemical bonding that combines lewis structures with orbitals. The buildup of electron density between two nuclei is visualized as occurring when a valence atomic orbital of one atom merges with that of another atom. The orbitals are said to share a region of space, aka to overlap.
• As the distance between the atoms decreases, the overlap region increases, so there is an increase in electron density between nuclei, and potential energy of the system decreases, and strength of the bond increases.
• However, if the nuclei become too close, energy increases due to electrostatic repulsion between nuclei

Question 19

internuclear distance

Answer 19

The internuclear distance at minimum potential energy is the bond length. Bond length is the distance at which the attractive forces between unlike charges are balanced by the repulsion forces between like charges

Question 20

hybrid orbitals

Answer 20

to explain geometries, assume that the atomic orbitals on an atom mix to form new hybrid orbitals. Hybridization: mixing atomic orbitals as atoms approach each other to form bonds. Number hybrid orbitals=number of atomic orbitals mixed

Question 21

internuclear axis

Answer 21

The electron density in covalent bonds is concentrated symmetrically about the line connecting the nuclei (the internuclear axis)

Question 22

sigma bonds

Answer 22

In single bonds, the line joining the two nuclei passes through the middle of the overlap region. These bonds are all called sigma () bonds. (the overlap region is sigma bond)

Question 23

pi bonds

Answer 23

In multiple bonds, the sideways overlap of p orbitals (perpendicular to internuclear axis) produces as pi () bond. In pi bonds, the overlap regions lie above and below internuclear axis and there is no probability of finding the electron on the internuclear axis. Pi bonds are weaker than sigma bonds because the overlap is less in pi than sigma

Question 24

sigma and pi bonds equal multiple bonds

Answer 24

Single bond=sigma bond, double bond=one sigma bond, one pi bond, triple bond=one sigma bond, two pi bonds

Question 25

delocalized2

Answer 25

• In most molecules, Bonding electrons are localized: the sigma and pi electrons are associated totally with the two atoms that form the bond. In molecules with resonance structures with pi bonds: not entirely localized.
• Delocalized: cannot describe the pi bonds as individual electron-pair bonds between neighboring atoms (ex: benzene because it has two resonance structures)

Question 26

molecular orbital theory5

Answer 26

• alternative model to explain bonding that describe the electrons in molecules by using specific wave functions called molecular orbitals (MO)
• An MO can hold max two electrons with opposite spins, has definite energy, and has electron density distribution. However, unlike atomic orbitals, MOs are associated with the entire molecule, not with a single atom
• Whenever two atomic orbitals overlap, two molecular orbitals form
• atomic orbs combine with other atomic orbs of similar energy
• as overlap increases, the energy of the bonding MO decreases and the energy of the antibonding MO increases

Question 27

bonding MO

Answer 27

lower-energy MO concentrates electron density between two nuclei. Sausage shaped. Summed the two atomic orbitals into one. Electron is more stable (lower energy) because it is attracted to both nuclei. This is covalent bond

Question 28

antibonding MO

Answer 28

higher-energy MO has very little electron density between the nuclei. Atom orbitals cancel each other in the region between the nuclei, greatest electron density on opposite sides of nuclei. An electron in this MO is repelled from boning region and is less stable (higher energy)

Question 29

energy level diagram

Answer 29

shows interaction between two 1s atomic orbital and molecular orbitals that result. Bonding electrons=electrons occupying a bonding molecular orbital.

Question 30

bond order

Answer 30

stability of covalent bond, =half the difference between the number of bonding electrons and the number of antibonding electrons. Bond order= ½ (bonding electrons – antibonding electrons). Bond order of 1=single bond, 2=double bond, 3=triple bond, 0=no bond exists.

Question 31

first row diatomic molecules

Answer 31

only 1s atomic orbitals used and only sigma MOs produced

Question 32

second row diatomic molecules5

Answer 32

• 2s, 2px, 2py, 2pz atomic orbitals used.
• 2s atomic orbs form sigma2s MO
• 2pz forms sigma2p MO
• 2px and 2py form pi2p MO
• The sigma2p MO has lower energy (more stable) than pi2p MO because there is greater overlap for sigma MOs

Question 33

doubly degenerate

Answer 33

Both pi2p and pi*2p MO are doubly degenerate: two orbitals (same energy) MOs of each type.

Question 34

second row MO using sp interactions2

Answer 34

• LARGE INTERACTION= HOPSCOTCH: B2, C2, and N2, the sigma2p MO is above the pi2p MOs in energy.
• SMALL INTERACTION= SIG, PI,PI, SIG: For O2, F2, and Ne2, the sigma2p MO is below the pi2p MOs in energy because of 2s-2p interactions

Question 35

paramagnetism2

Answer 35

• molecules with one or more unpaired electrons are attracted into a magnetic field. The more unpaired electrons= stronger force of attraction
• gains weight if in magnetic field

Question 36

diamagnetism2

Answer 36

• substances with no unpaired electrons are weakly repelled from a magnetic field. Weaker effect than paramagnetism
• loses weight if in a magnetic field

Question 37

bond order/length/enthalpy trend

Answer 37

As bond orders increase=bond distance decrease and bond enthalpies increase

Question 38

heteronuclear diatomic molecules2

Answer 38

• two atoms are not the same, (NO). if the atoms in heteronuclear diatomic molecule don’t differ too greatly in electronegativities, MO will be similar to homonuclear diatomics, except the atomic energies of the more electronegative atom will be lower in energy than those of the less electronegative element.
• An MO will have a greater contribution from the atomic orbital to which it is closer in energy. Not an equal mixture between the two atoms