The Intermolecular Forces (forces between molecules) In general, the weaker the intermolecular forces, the less energy which is required for the substance to undergo a physical change ( change in state). – Thus the physical properties of melting point and boiling point gives a rough measure of the intermolecular forces which bind the molecules to each other. • The higher the melting and boiling points, the higher the strength of the intermolecular forces which bind the molecules together. • The intermolecular forces: 1. Forces between covalent molecules – London dispersion forces (Van der Waals forces) – Dipole-dipole attraction – Hydrogen bonding 2. Ionic bonding 3. Metallic bonding 4. Network covalent bonding 1. Covalent molecule forces a) London Dispersion Forces: are the dominant forces between covalently bonded, non-polar molecules. Ex. N2, O2, CH4 • formation of “ instantaneous dipoles” • weak and act over a short distance. • Force increases with the increase in the number of electrons. London Dispersion Forces • Small molecules generally have a low melting and boiling points ( molar mass of less than 50, generally will be a gas at room temp) • -covalent compounds are usually gases, liquids or low melting solids. Force varies with the shape of the molecule. • A long molecules has many electrons exposed and therefore a relatively high boiling point. The structure is flexible and does not stack well, so it has a low melting point. • A more compact molecule has fewer exposed electrons and will have a low boiling point but stacks better, so its melting point will be higher. Ex Normal pentane (C5H12) mp -130°C bp 36°C Neopentane (C5H12) (2,2 dimethylpropane) mp -20°C bp 9°C b) Dipole-Dipole Attraction: is a force which acts between polar molecules (ex: H2S ). It results from the attraction of opposite poles of permanent dipoles. - polar molecules will generally have higher melting and boiling points than nonpolar molecules. c) Hydrogen Bonding: occurs when a lone pair on a O, N or F atom forms a “bond” to the partial positive charge of a H atom (part of a polar covalent bond). - molecules will have high boiling and melting points. Figure 4.4 The hydrated proton Relative magnitudes of forces The types of bonding forces vary in their strength as measured by average bond energy. Strongest Covalent bonds (400 kcal) Hydrogen bonding (12-16 kcal ) Dipole-dipole interactions (2-0.5 kcal) Weakest London forces (less than 1 kcal) 2. Ionic bonding: generally occurs in compounds of metals and nonmetals. Ex. NaCl, K2SO4, NH4Cl. It is the result of the attraction of the opposite charges. -formation of a crystal lattice. -ionic solids are called crystals and are very brittle and hard -ionic solids have very high melting and boiling points. 3. Metallic bonding: Most metals have only one or two valence electrons and low ionization energies. The valence electrons do not seem to belong to any individual atom but move easily from one atom to another. Metals can be thought of as positive ions immersed in a “sea” of mobile electrons. The attractive forces that bind metals together are called metallic bonds. Ex: Hg, Cu, Au, Fe, alloys. -high electrical and thermal conductivity -shiny surfaces are due to the de-excitation of the electrons (the result of valence electrons absorb and re-emit light energy) -strength varies with the nuclear charge of the metal atoms and the number of electrons in the metal’s sea of electrons -metals can be flattened out (malleable) or stretched out into a wire (ductile) because the electrons can move into other positions without breaking up the essential structure. -Very low to very high melting point -Very low to very high boiling point 4. Network Covalent Bonding: Occurs between two non-metals -network is infinite -includes diamond, graphite, quartz and most rocks. -covalent is stronger than other bonds due to a hard network. -solids are brittle -network solids have high melting (1200 – 3500°C) and boiling points.