Bonding 2 Molecular Shape & Bonding Flashcards
Valence Bond Theory
- Based on the principle that
covalent bonds form from the overlap of atomic orbitals on the bonding atoms - Atoms come together to maximize the overlap between the orbitals (but not so close so that they repel each other)
Hybridization Defined
- A theoretical concept used to explain bonding, hybrid orbitals are not “real” and we never use them for isolated atoms
- The mixing of at least two nonequivalent atomic orbitals (i.e. s and p orbitals) – so the hybrid orbitals are not pure atomic orbitals,
and they have different shapes and orientations
Hybridization Steps
Step 1: Promotion (excitation) of electron(s)
Step 2: Mix the relevant orbitals
Step 3: Bond hybrid orbitals with
unpaired electrons on other atoms
Rules of Hybridization
1) Orbitals in = orbitals out
▪ So: 1 s + 3 p orbitals = four sp^3 orbitals in total
▪ 1s + 2p orbitals = three sp^2 orbitals
▪ 2d + 1s + 3p orbitals = six d^2sp^3 orbitals
2) The energy of hybrid orbitals is in-between the energies of the constituents
– Energy of the sp3 hybrids in carbon is between the energies of the carbon 2s and 2p orbitals
Hybridization with Lone Pairs
If the central atom contains one or more lone pairs, the procedure is exactly the same except that one or more of the hybrid orbitals will be
occupied by lone pairs
Hybridization in Water
- In water, two of the hybrid orbitals end up occupied by lone pairs
- VSEPR gives the shape (bent), two hydrogen atoms are located along two of the tetrahedral axes, and the two lone pairs extend along the other two
- H-O-H bond angle is smaller than the tetrahedral value of 109.5o because larger size lone pairs push the atoms together
Size of Lone Pair and Bonding Pair Orbitals
- Bonding pair orbitals are smaller than lone pair orbitals because they are more confined along the bond axis than the diffuse lone pairs
- Computational chemistry allows us to calculate the size of the orbitals
Hybridization and the Octet Rule
- No matter the hybridization type, the result is always four orbitals,
enough to hold eight electrons - The maximum electrons that a second-period atom can hold is eight
- To go beyond the octet rule, we need to use d orbitals for hybridization
Hybridization Using d Orbitals
- A total of six orbitals will be hybridized, resulting in six hybrid orbitals called sp3d2 hybrids
- After bonding with six fluorine atoms, sulfur will have twelve electrons around it (expanded octet)
- This cannot happen for second-period elements since there isn’t a 2d subshell, and it would require too much energy to excite electrons to the 3d subshell
Sigma Bonds
Sigma (σ) Bonds: Two orbitals overlap end-to-end
Has a greater overlap and is thus stronger than the pi bond
Pie Bonds
pi (π) Bonds: Orbitals overlap side-
by-side
Problems with Valence Bond Theory
– It approximates the orbitals in a molecule as belonging to individual atoms, without consideration of the molecular environment in which they reside
– There are also molecules for which valence bond theory fails, such as O2:
Molecular Orbital Theory
Describes covalent bonds in terms of
molecular orbitals
Molecular Orbitals
- Molecular orbitals come from the interaction of atomic orbitals of the bonding atoms
- They are associated with the entire molecule, not just one atom like atomic orbitals are
- Each molecular orbital can only hold two electrons – Pauli’s Exclusion Principle also applies in MO theory
Forming Molecular Orbitals
- When two atomic orbitals interact in MO theory, two new molecular orbitals are formed
- Are two ways of combining the two atomic orbitals – adding them or subtracting them
- This results in two very different molecular orbitals
Hybrid vs Molecular Orbitals
Hybrid Orbitals: Are the result of mixing atomic orbitals on the same atom
MOs: Are the result of mixing atomic orbitals on different atoms
Constructive Interference Defined
If the phases (signs of the wavefunction - positive or negative) are the same and the interaction is favorable
Destructive Interference Defined
If the phases are the opposite and the interaction is unfavorable
Bonding Orbitals Defined
If we get constructive interference, the energy of the molecular orbital is lower (more stable) than the energy of the atomic orbitals from which it was made
Antibonding Orbitals Defined
If we get destructive interference, the energy of the molecular orbital is higher (less stable) than that of the separated atomic orbitals
Naming Molecular Orbitals
First of all, classify the MO as sigma or pi:
– If the electron density in the orbital is centered symmetrically around a line connecting the two bonding
atoms, then the orbital is a sigma (σ) orbital
– In a pi (π) orbital, the electron density is concentrated above and below a line connecting the two bonding atoms
- The symbol (σ or π) gets a subscript indicating which atomic orbitals (on each atom) are involved in forming the MO
- Finally, if the MO is an antibonding orbital, it gets a superscript star
Stabilities of Molecules
- Diatomic molecules (and cations) are made from elements of the first row (H2, H2+, He2, He2+)
- In all cases, there are two molecular orbitals, formed by the interaction of 1s orbitals on the atoms (σ1s and σ*1s)
- We can measure the strength of a bond (and here the stability of the species) by calculating the bond order
Bond Order Formula
Bond order = ½ (bonding e- -antibonding e-)
Bond Order Characteristics
- A higher bond order indicates a more stable bond
- A bond order of zero (or below) indicates that the bond is unstable
- Bond order can only be used for qualitative purposes, not to get quantitative values for bond energies
(i.e. two different compounds with a bond order of 1 will not necessarily have the same bond length or strength)
