Bonding and Properties Unit 4

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)

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)

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

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

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

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

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

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.

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.

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

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

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