Chemical Bonding Flashcards
Topic 3, Lectures 9-12 - Ben Ward
Octet rule
Atoms bond together in order to achieve a stable electron configuration—one that has a full outer shell.
Octet rule weaknesses
There is no reason as to why noble gas confiduration is the most stable; it is based on observation alone. Oftentimes, the Octet rule doesn’t apply.
Ionic and Covalent character
While bonding is usually dealt with as either ionic or covalent, in reality many compounds exhibit bonding in between the two etremes. Therefore, sometimes compounds are said to have a covalent or ionic character.
Ionic bonding
Bonding is based purely on electrostatic attraction. Most common in groups 1, 2, 17, 18. This is because as you remove electrons, ionisation energy increases, so for groups 13–15, ionisation energy becomes so high it is no longer favourable.
Colvalent bonding
When the ionisation energy is too high to favour ionic bonding, atoms share electrons instead.
What two models explain covalent bonding?
Lewis model and molecular orbital theory
Lewis (1916)
Thought chemical bonds consisted of a shared pair of electrons between two atoms. For each atom bonded to atom X, X gains one electron; the number of bonds an atom can make is therefore determined by its number of valence electrons. If an atom requires several bonds, it can bond to several atoms or form double/triple bonds.
Issues with the Lewis model
Assumes that all valence electrons are equal, which isn’t true as bonds are formed by the overlap of orbitals. It is difficult to deal with excited states, e.g., some bonds change lengths when e- are excited. Some experimental observations are impossible to explain, e.g., O2 has lone pairs and therefore appears blue, despite the Lewis model claiming it has a full octet.
Hybridisation
Atoms mix their orbitals in order to orient themselves in a direction that benefits bonding. Forms stronger and more stable bonds.
sp3 hybridisation
Combination of 1 s orbital and 3 p orbitals to form 4 equivalent orbitals. CH4 is an example. Forms a tetrahedral shape.
sp2 hybridisation
1 s orbital and 2 p orbitals combine 3 orbitals. C2H4 is an example. Forms a trigonal planar shape.
sp hybridisation
1 s and 1 p orbital combine to form 2 orbitals. C2H2 is an example. forms a linear shape.
Breakdown of the octet rule
PCl5: forms 5 bonds (not 4) and has more than 8 electrons in the outer shell. P has 5 valence electrons; if it only used 3s and 3p orbitals, it could only form 3 bonds. In order to form 5 orbitals, the orbitals hybridise with the 3d orbital as they are energetically similar. Cannot use 3d orbitals with 2 orbitals as the energy gap is too large; a large amount of energy would be required to mix the orbitals.
Molecular orbitals
Similar to atomic orbitals but surround the whole molecule rather than the nucleus.
Linear Combination of atomic orbitals
Method of constructing molecular orbitals from constituent atomic orbitals. If there are X AOs, there must be X MOs. LCAO method states that the wavefunction of a molecular orbital is the product of the wavefunction of bonding atomic orbitals.
Bonding MO
The product of an in phase combination of the AOs. Contributes towards bonding between nuclei and has a large amount of electron density (wavefunction value) between nuclei.
Anti-bonding MO
Product of an out-of-phase combination of AOs. It reduces bonding interactions and will have a node at the midpoint between nuclei.
MO energies
Electrons that have a high probability of being found between nuclei will be more stabilised as they are attracted to both nuclei simultaneously
Order of MO energies
Since greater electron density = stable = low energy, ABMO (highest), non-interacting AOs, BMO (lowest).
MO energy level diagrams
Electrons are placed into BMOs first as lowest energy fills first
Magnetic properties
All electrons paired: repelled by a magnetic field = diamagnetic. One or more unpaired electrons: attracted to a magnetic field = paramagnetic.
Bond order
Difference between bonding and antibonding electrons in a molecule, indicating strength of bond. (Bonding - antibonding)/2
MO nomenclature
Orbitals are called sigma if their phases are unaffected by rotating 180 about the z axis. They are called pi if their phases are reversed by rotating 180 about the z axis.
s/p orbital mixing
In elements such as F2 and O2, 2s and 2p orbitals are far enough in energy to not interact; however, in N2, s and p orbitals are closer in energy, which causes a change in order of the MOs in the energy level diagram.
Atomic orbital energies
Energy tends to decrease as you move across the periodic table. This is due to increasing Zeff having greater interaction with electrons. When constructing homonuclear diatomic MO diagrams, AOs must have the same energies, which is not the case with heteronuclear systems. The element with the greater electronegativity will have lower energy atomic orbitals.
HOMO and LUMO
Highest occupied molecular orbital, lowest unoccupied molecular orbital, involved in chemical reactions.
