Advanced Bond Theories Flashcards
Valence Bond Theory
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.
Bond Energy Diagram
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.
Hybrid Orbitals
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.
How to write Hybrid Orbitals
Write ground state, then promote electron, then rewrite orbital diagrams to show hybrid orbitals.
sp hybrid orbitals
Two degenerate orbitals are formed form hybridization of one (s) and one (p) orbital. Linear electron region geometry. (2 electron domains)
sp^2 hybrid orbitals
Three orbitals are formed from hybridization of one (s) and two (p) orbitals. Trigonal planar region geometry. (3 electron domains)
sp^3 hybrid orbitals
Four orbitals are formed from hybridization of one (s) and three (p) orbitals. Tetrahedral region geometry. (4 electron domains)
sp^3d
Five orbitals formed from hybridization of one (s), three (p) orbitals and one (d) orbital. Trigonal bipyramidal geometry. (5 electron domains)
sp^3d^2
Six orbital formed from hybridization of one (s), three (p) orbitals and two (d) orbitals. Octahedral geometry. (6 electron domains)
Sigma bonds
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.
Pi bonds
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).
Multiple bonds
Always one (sigma) bond, the rest are (pi) bonds.
Delocalized bonding
Occur in the presence of resonance structures.
Molecular orbitals
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.
Molecular orbitals vs atomic orbitals
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.
Formation of molecular orbitals
When (n) atomic orbitals (AOs) overlap, (n) molecular orbitals (MOs) form.
Bonding molecular orbitals
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.
Antibonding molecular orbitals
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.
Bond order
1/2 (# of bonding electrons - # of antibonding electrons)
Attractive and repulsive forces
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.
Band theory
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.
Orbital overlap
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.
s-d band and metallic bonding
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.
Conductors
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.
Semiconductors
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.
Insulators
Have a large band gap between the valence and conduction band, so there is no flow of electrons because the gap is too large.
