Advanced Bond Theories

Question 1

Valence Bond Theory

Answer 1

Occurs when orbitals on the two atoms overlap (the mechanism for which electrons go back forth). Energy of the system depends on the quality of overlap between orbitals - s orbitals have more effective overlap than d orbitals.

Question 2

Bond Energy Diagram

Answer 2

Compares distance between two orbitals and the energy. Minimum distance is where bond is strongest because lower energy is the most stable. If the distance is too close, repulsive forces between nuclei increases the energy of the system.

Question 3

Hybrid Orbitals

Answer 3

The idea of applying orbital overlap and valence-bond theory to polyatomic molecules. Mix/combine/hybridize (n) atomic orbitals to make (n) hybrid orbitals (same number). Number of letters corresponds to the number of electron region domains on an atom.

Question 4

How to write Hybrid Orbitals

Answer 4

Write ground state, then promote electron, then rewrite orbital diagrams to show hybrid orbitals.

Question 5

sp hybrid orbitals

Answer 5

Two degenerate orbitals are formed form hybridization of one (s) and one (p) orbital. Linear electron region geometry. (2 electron domains)

Question 6

sp^2 hybrid orbitals

Answer 6

Three orbitals are formed from hybridization of one (s) and two (p) orbitals. Trigonal planar region geometry. (3 electron domains)

Question 7

sp^3 hybrid orbitals

Answer 7

Four orbitals are formed from hybridization of one (s) and three (p) orbitals. Tetrahedral region geometry. (4 electron domains)

Question 8

sp^3d

Answer 8

Five orbitals formed from hybridization of one (s), three (p) orbitals and one (d) orbital. Trigonal bipyramidal geometry. (5 electron domains)

Question 9

sp^3d^2

Answer 9

Six orbital formed from hybridization of one (s), three (p) orbitals and two (d) orbitals. Octahedral geometry. (6 electron domains)

Question 10

Sigma bonds

Answer 10

Head-to-head overlap, stronger than (pi) bonds, have cylindrical symmetry of electron density about the internuclear axis (when they rotate, they are always the same). Single bonds are always (sigma) bonds because their overlap is greater, resulting in a stronger bond that is more energy lowering.

Question 10

Pi bonds

Answer 10

Each pi bond has 2 unhybridized (p) orbitals. Side-to-side overlap, occur in multiple bonds, weaker than (sigma) bonds. Electron density is above and below the internuclear axis (not symmetrical, can’t rotate without breaking pi bond).

Question 11

Multiple bonds

Answer 11

Always one (sigma) bond, the rest are (pi) bonds.

Question 12

Delocalized bonding

Answer 12

Occur in the presence of resonance structures.

Question 13

Molecular orbitals

Answer 13

Aspects of bonding not explained by Lewis structures, VESPR theory and hybridization. Can be used to explain the interaction of oxygen with magnetic fields, why some molecules are coloured, and metallic bonding (conduction.) Also explains why electrons in atoms have atomic orbitals and electrons in molecules have molecular orbitals.

Question 14

Molecular orbitals vs atomic orbitals

Answer 14

Each contain 2 electrons with opposite spin, each has definite (quantized) energy, and electron density distribution can be visualized with contour diagrams. However, molecular orbitals are associated with the entire molecule, unlike atomic orbitals.

Question 15

Formation of molecular orbitals

Answer 15

When (n) atomic orbitals (AOs) overlap, (n) molecular orbitals (MOs) form.

Question 16

Bonding molecular orbitals

Answer 16

Have electron density centered around the internuclear axis. Constructive combination of AOs results in a bonding molecular orbital that has a lower, more stable energy.

Question 17

Antibonding molecular orbitals

Answer 17

Have electron density centered around the internuclear axis. Destructive combination of AOs results in an antibonding molecular orbitals that has more energy. It forms a node, resulting in no bonding (antibonding.) (*) denotes an antibonding orbital. Less stable.

Question 18

Bond order

Answer 18

1/2 (# of bonding electrons - # of antibonding electrons)

Question 19

Attractive and repulsive forces

Answer 19

Attractive forces are between electron and nuclei, while repulsive forces are between electron and electron or nuclei and nuclei. Repulsive forces are ONLY greater in bonds that DO NOT FORM; to form a bond, attractive forces must be greater.

Question 20

Band theory

Answer 20

As MOs fill, spacing between orbitals gets smaller. Can help us understand the bonding in metals. Always as many MOs as AOs. Energy differences between orbitals within a band are tiny, so electrons can readily move through the metal. This is why metals are great conductors.

Question 21

Orbital overlap

Answer 21

Overlap in s-orbitals is better than the overlap in d-orbitals. Good overlap = wide (long) band, poor overlap = short band. For metals in groups 1 - 12, s & d bands overlap. S band can hold up to 2 electrons per atom, d up to 10.

Question 22

s-d band and metallic bonding

Answer 22

Moving from groups 1 - 12, s-d band fills with electrons. When the band is half full, all of the bonding orbitals are full and all of the antibonding orbitals are empty. Metal-metal bonding is maximized and metals are very hard. However, when the band is full, all of the bonding orbitals are full and all of the antibonding orbitals are also full, so the effects cancel each other out. Metal - metal bonding is minimized and metals are very soft. So, metals in group 6 (half-full) are the hardest and metals in group 12 (full) are the softest and can be liquid at room temperature. Tungsten, the hardest metal, has electrons that are further away from the nucleus and can be more involved with bonding.

Question 23

Conductors

Answer 23

Materials can conduct electricity if electrons can move easily from an occupied orbital to an unoccupied one. Metals do not have a gap between the valence band and conduction band.

Question 24

Semiconductors

Answer 24

Have a small band gap between the valence and conduction band, so there can be a small flow of electrons depending the the surrounding environment.

Question 25

Insulators

Answer 25

Have a large band gap between the valence and conduction band, so there is no flow of electrons because the gap is too large.