2.1

Structure and Bonding

Ionic Bonding

• large difference in electronegativity ( >1.9)
• complete transfer of electrons
• ionic compounds normally exist as crystals… high melting points due to strong attractive forces
• not directional (no orbital shapes)

Coulomb’s Law:

where k= coulombs constant, Q_1,2 are the values of the charged ions, and r is the distance between the ions

• • the bigger the atoms, the smaller the force
  • smaller atoms have greater melting points

Covalent Bonding

• directional (determines orbital shapes)
• sharing of e^-
• can be covalent or polar covalent

Lewis Structures

• applies only to stable species
• put the valence e^- around the outside of the atom
• each atom in a molecule shares valence e^- to attain the e^- configuration of the noble-gas atom of the same period
• Can Not Over Fill – Can only bond to the number of valence e^- it has available

Formal Charges

• Formal Charge = Valence e^- ⎻ Non-bonding e^- ⎻ Bonds
• The # of e^- and bonds (each bond= 1e^-) compared to the number of valence e^-
• Each bond counts for 1 e^-
• The best structure has minimized formal charges and must be consistent with the electronegativities of the atoms

Octet Rule

• Can Not Over Fill & H MUST be surrounded by 2e^-

Resonance Structures

• Differ in the position of the e^-
• The actual structure is an average of the 2 resonance structures – called a resonance hybrid
• Allows one to estimate the energy of a molecule relative to others

Bond Order

• Single bond has a bond order of 1, double bond = 2, triple bond = 3 …

Line Diagrams

• solid line is a covalent bond
• C atoms are at the intersection of 2 line unless stated otherwise
• H atoms surround the C atom
• Non-bonding e^- are often omitted
• Not C atoms (heteroatoms) must be drawn with the H atoms surrounding it

2.2

VSEPR Theory

• Provides information about the 3D shape
• , where m is the # of X atoms bonded to the central atom, and n is the # of lone pairs
• Look at lone pairs and bonded atoms to determine electronic arrangement
• Ignore lone pairs for molecular arrangement
• SEE VSEPR SHAPES

Line-Wedge Symbols

• Line: represents a bond in the plane of the page
• Dashed Wedge: a bond going away from you (into the page)
• Solid Wedge: a bond going toward you (out of the page)

Molecule Polarity

• Polar: molecules with a dipole moment > 0, asymmetrical shapes
• Non-polar: molecules with a dipole moment = 0, symmetrical shapes

2.3

Valence Bond Theory

• Covalent bonds are formed by overlap of atomic orbitals (s, p, d, hybrid orbitals)
• e^- are localized in bonds
• bonds are the result of bond overlap of orbitals on the same axis (head-on)
• bonds are the result of side-to-side overlap of p orbitals, bonds form above and below the central atom due to the shape of p orbitals
• double bonds have a bond
• triple bonds have bonds

Hybridization

• fixed combinations (mixtures) of 2/more valence atomic orbitals
• hybrid orbitals were invented to make VB theory consistent with observed molecular shape (Linus Pauling)
• sp = 2 regions of density, sp²= 3 region, sp³= 4 regions, sp³d= 5 regions, sp³d²= 6 regions
• lone pairs are included as regions of density
• ex. PCl₃ has 4 regions of e^- density

Boron Hydrides

• compounds containing boron-hydrogen bonds called boranes
• boron has an e^- deficient bond which doesn’t allow it to form BH₃ only B₂H₆ (diborane)
• 3-centre, 2-electron bond: where one electron pair is shared between 3 atoms

2.4

Molecular Orbital Theory

• MO theory made because…
  • • VB theory incorrectly states that e^- are localized, but resonance structures are evidence that this is untrue
    • VB doesn’t provide the relative energies of the electrons
    • VB doesn’t determine magnetism
• Molecular orbitals in MO theory are formed by combining the wavefunctions of atomic orbitals (wave-particle duality)

Constructive Combination

• When 2 atomic orbitals are in the same phase & interact
• No new nodes are formed between the nuclei

Destructive Combination

• 2 orbitals of opposite phase combine
• generates a new node between the nuclei
• doesn’t lead to bond formation… antibonding

MO Energies

• the MO formed by constructive combination is stabilized therefore, lower in energy
• the antibonding MO creates a new node and exposes the positively charged nuclei to each other, it is destabilized therefore higher in energy

Principles of MO theory

1. # of MO’s formed = total # of atomic orbitals combined
2. Bonding MO’s are always lower in energy than antibonding or zero bonding
3. Must obey Pauli exclusion principal (2 e^- with opposite spins in each orbital) and Hund’s rule (half fill all before doubling up)

MO Bond Order & Stability of Molecules

• species with higher bond orders are more stable, a bond order of 0 means not net e^- holding the molecule together

Paramagnetism & Diamagnetism

• Paramagnetism: due to unpaired magnetism (at least 1 unpaired electron)
• Diamagnetism: even spin (not magnetism)

* The lewis structure for O₂ incorrectly predicts it to be diamagnetic

Diatomic (Homonuclear) Second Row Elements

• Valence electrons in the n=2 orbital
• 8 MOs formed (2s, 2p_x, 2p_y, 2p_z) each with 1 bonding & 1 antibonding
• ‘ * ‘ indicates that they are antibonding orbitals
• sideways overlap MOs are due to in-phase (constructive) or out-of-phase (destructive) combinations
• π MOs are formed from constructive combining, & π* MOs are formed from destructive combination (for 2p_y + 2p_y or 2p_z + 2p_z)
• if the orbital lies on the plane it is xy, yz, or xz, if it lies on the axis subscript is the same as the axis x², y², z²

Relative Energies

• varies depending on molecule
• π_2py & π_2pz MOs are degenerate (same energy) because they are formed by the same kind of overlap
• σ_2px has a different energy because it was formed by head-on overlap

Semiconductors & Band Theory

• used to explain the insulting, semiconducting, and conducting properties of various materials
• conductor conducts electricity, insulator does not, and semiconductor has properties of both
• as the # of atoms increases, the energy gap between the bonding and antibonding decreases
• the upper half of the energy band is the conducting band, and lower half is the valence band
• Band Gap: energy gap ΔE_g between the valence and conductor band

If ΔE_g < 0.5 eV = conductor, ΔE_g = 0.5-3 eV = semiconductor, ΔE_g > 3 eV = insulator

* Insulators have larger band gaps, conductors have 0 band gap

• as temperature increases, conductivity increase

because more thermal energy is added to the system

2.5

Intermolecular Forces: between molecules to hold them together

Types (in order of decreasing strength)

• • Van der Waals Forces “Dispersion”: temporary cloud of e^-, surface area is a factor to determine strength (larger SA = larger dispersion)
  • Dipole-Dipole Interactions: permanent dipoles of polar molecules
  • Hydrogen Bonding: when H is bonded to highly electronegative atom (N,O, Cl, F), attaches through a lone pair to a nearby molecule

Boiling Points & Melting Points

• Stronger IM forces = higher b.p. (for similar size molecules)
• Boiling points increase with molecular mass (due to dispersion forces)
• Trends similar for melting points, highly symmetric molecules have higher melting points than irregularly shaped molecules of the same type and mass
• Intermolecular forces are broken