1. Bonding and Properties
Unit 4

2. Atomic bonding Metallic bonds (metals)
Positive ions in a “sea” of electrons
Generally strong bonds
Held together by attraction of positive ions for sea of electrons
Electrons are delocalized (not attached to specific atom)
Properties
High conductivity
Malleablilty and ductility
High MP
Network (non-metals)
Involves directional covalent bonding and van der Waals forces
Atomic bonding makes giant molecule of single element (carbon)
General properties
Strong directional bonds
Brittle
No conductivity of heat or electricity
Group 8A (Noble gases)
Weak dispersion forces that increase with # of e-
Increased dispersion forces as you go down group (explains MP, BP)

3. Molecules: Intramolecular
Ionic
Transfer of e- between atoms and formation of crystal lattice through electrostatic attraction of oppositely charged ions
Properties
Brittle
Non-conductive as a solid
Conductive when dissolved in water
High MP
Soluble in water
Covalent
Sharing e- in a directional bond (between 2 specific atoms)

4. Shapes of molecules, VSEPR, polarity
VSEPR = Valence Shell Electron Pair Repulsion Theory
Three postulates
Valence electrons are involved in bonding
Lone electron pairs and shared electron pairs on the central atom repel each other
Lone pairs repel more than shared pairs
notice the shapes with tetrahedral electron arrangements
trigonal planar has a bond angle of 107 rather than 109.5 because the lone pair repels the bonding pairs more than they repel each other
Based on the Lewis structures the shape of the molecule can be determined
the shape is based on the location of the nuclei in a molecule
So double and triple bonds count as one shared pair when determining the shape of the molecule
locate the shared pairs and lone pairs on the central atom
Determine the shape based on the chart on the next slide

6. Bond strength, length Bond length: distance between bonded nuclei
Bond strength: energy required to break bond
Multiple bonds = shorter, stronger than single
single
double
triple
Decreasing length
Increasing strength

7. Polarity of molecules
Determined by shape and difference in electronegativity
Ionic: difference of more than 1.8
Polar covalent: difference of about 0.3 to 1.8
Non-polar covalent: difference of about 0
Indicate polarity of bonds using partial + (δ+) partial – (δ-)
More electronegative atom is partial – (δ-)
Depending on shape and individual bond polarity will determine whether the molecule itself is polar

8. Molecules: Intermolecular forces(IMF)
Network covalent
Ceramics, glasses, some metal oxides bond this way
Same as atomic network, same kind properties
Dispersion forces/ Van der Waals Forces
Weak forces from induced dipoles
Induced dipoles
Random movement of electrons can result in a bunch of electrons located on one side of the molecule which gives that side of the molecule a slightly negative charge. A molecule near it will have a positive charge induced on it as the electrons in it move away from the negative charge.
Larger molecule, stronger forces b/c increased e- and mass

9. Dipole – dipole forces Strength: Forces between permanent dipoles
Hydrogen bonds special case
Involve hydrogen attached to O, N, F
Stronger than dipole- dipole
Is NOT a bond but an attraction between 2 molecules
Strength:
intramolecular > hydrogen bond > dipole-dipole > dispersion forces

10. This is NOT a hydrogen bond. This is a polar covalent bond.
These are hydrogen bonds. They are attractions between a partially positively charged atom on one molecule and a partially negatively charged atom on another molecule.
Coventional bonds make molecules.
Hydrogen bonds do not make molecules. They are attractions between molecules that are already made.

11. Determining IMF
Ionic – intermolecular forces is kind of a misnomer since there isn’t an ionic molecule but the forces that hold the formula units together are ionic bonds
Polar molecule
intermolecular forces will be dipole-dipole forces and
may be Hydrogen bond
polar molecules have permanent dipoles
Non-polar molecule
Intermolecular forces will be dispersion forces
Dispersion forces result from a temporary induced dipole
The larger the molecule the more important dispersion forces become
Larger molecule means more electrons moving around so there is more chance for induced dipoles.

12. IMFs Affect on physical properties
Stronger forces, more sticking together
Increase
MP (more energy to break IMF)
BP (more energy to separate to gas phase)
Viscosity (stronger attraction, more resistant to moving)
Surface tension (stronger IMF, more effort to break apart surface molecules)
Decrease
Volatility (weaker forces, more easily escape to gas phase)
Vapor pressure ((more gas molecules, higher vapor pressure)
Dependent on temp
Equilibrium between states of matter

13. Solubility
Like dissolves like
Polar + polar = miscible
Nonpolar + nonpolar = miscible
Polar + ionic = miscible
Polar + nonpolar = immiscible
Nonpolar + ionic = immiscible
Why? Strength of IMF
-Ionic attractive to break attraction between polar solvent molecules
-Polar compounds attractive to break attraction between polar solvent molecules
-Non-polar molecules NOT attractive enough to break attraction between polar solvent
Polar and nonpolar molecular parts
Emulsifiers, soaps, organic molecules with OH groups

14. Dissolution (solute dissolving in solvent)
Energy change (energy required to break IMF, energy released when new IMF form)
Increase in entropy
Large non-polar molecules cannot make as many strong IMF as polar can Energy change for non-polar solute dissolving in polar solvent is unfavorable
Colligative Properties (for solutions)
Boiling point elevation, freezing point depression
Lower vapor pressure
Vapor pressure lowered, fewer solvent molecules escaping to gas phase when solute is present
Fewer molecules able to escape to vapor phase b/c strong attraction between solute and solvent molecules