 held tight by strong electrostatic forces in between cations and anions  non-volatile, high mp, high bp  solid at room temp  hard and brittle because of a 3-D lattice/crystalline structure  conductivity › non-mobile e- as solid = no › ions move freely when melted = yes › cations and ions separate when dissolved in water = yes  soluble in polar solvents like water

 strong intramolecular forces, weak intermolecular forces, usually liquids or gases at room temp or soft solid › strength of polarity and strength of London forces determine mp and bp  greater polarity = higher mp and bp  greater van der Waals’ = higher mp and bp  often dissolve in nonpolar solvents but not in strong polar solvents like water  do not conduct electricity

 high mp and bp › decreases going down the periodic table  harder for cations to attract the sea of electrons › increases going across the periodic table  atomic radii becomes smaller, easier to attract the sea of electrons  low volatility  not soluble in most solvents (polar or non-polar)  conduct electricity well because of moving sea of electrons

 from highest to lowest 1. metallic bonds 2. ionic bonds (cations and anions) 4. hydrogen bonding (strong δ+ or δ-)  very strong when H is bonded with NOF (nitrogen, oxygen, or fluorine) 5. dipole - dipole *δ+ or δ-) 6. London forces (weak, temporary δ+ or δ-)

 generally speaking › the greater the intermolecular force (IMF) between the molecules, the higher the melting point, boiling point, and volatility (evaporate)  more electrons help increase the van der Waals’ forces and keep the substance in the liquid state  molecules that can stick together better remain a liquid at higher temps. boiling point increases this flat shape allows it to stick to one another better these round shapes do NOT allow them to stick to one another

hydrogen bonding can occur here which is the strongest type of dipole : dipole intermolecular force only normal dipole : dipole bonding can take place ethanol - higher BP dimethyl ether - lower BP Exampe: two Lewis structures for the formula C 2 H 6 O. Compare the boiling points of the two molecules.

 “like dissolves like” › polar substances tend to dissolve in polar solvents (ex. water dissolves ionic compounds)  dissociation of salt YouTube (:53) dissociation of salt YouTube (:53) › non-polar substances tend to dissolve in non-polar solvents (ex. alcohol dissolves covalent molecules)

++ –– ++ ++ –– ++ ++ –– ++ ++ –– ++ ++ –– ++ ++ –– ++ ++ –– ++ ++ –– ++ ++ –– ++ The dipoles of water attract, pushing the oil (with no partial charge) out of the way: attractions win out over the tendency toward randomness.

 substances must possess Freely Moving Charged Particles › this occurs in…  metals with their “sea of electrons”  YouTube (1:05) YouTube (1:05)  molten ionic compounds (+ and – ions can move)  chemistry/bonding/bonding5.htm chemistry/bonding/bonding5.htm  ionic compounds in aqueous solution (dissolved in water)  water pulls apart + and – ions and allows them to move

Type of Bonding Melting Point Boiling Point Volatilit y Electrical Conductivity Solubility in Non- polar Solvent Solubilit y in Polar Solvent Non- polar Low HighNoYesNo Polarvaries No Yes Hydroge n bonding varies No Yes Ionic Bonding high lowYes (molten or aqueous) NoYes (most) Metallic Bonding high lowYesNo Covalentvaries No