Liquids, Solids, and Van der waals (Intermolecular) Forces ( Ch 15)

1. Liquids, Solids, and Van der waals (Intermolecular) Forces (Ch 15) Suggested HW: (Ch 15) 13-17

2. States of Matter Differ By Intermolecular Distance • The state of a substance at a given temperature and pressure is determined by two factors: • Thermal energy of the molecules • Intermolecular forces (called Van der walls forces) between molecules

3. States of Matter • Gases: • thermal energy is greater than the energy of attraction between the gas molecules, so molecules have enough energy to separate • have completely free motion (translational, rotational, and vibrational) • Liquids: • the thermal energy is somewhat less than the intermolecular attractive forces, so the molecules are slightly separated • the thermal energy available allows “tumbling” of molecules, which is why liquids can be poured • restricted translational, rotational, and vibrational movement • Solids: • the thermal energy is much less than the energy of attraction. • the molecules are completelyfixed in space • vibrational motion only

4. Intermolecular Forces: Coulombic Attractions • As you recall, ionic compounds are solids at room temperature. There are ion-ion attractions in ionic compounds. • The coulombic force that holds ions together is very strong. Coulombic attractions are the strongest of all intermolecular forces. • Therefore, all ionic compounds have very high melting/boiling points. Cl- Na+

5. Dipole-Dipole Forces • Polar molecules attract one another. This type of intermolecular force is called dipole-dipole attraction. + + δ δ - - δ δ Dipole-dipole interaction: Weaker than intra-molecular forces Covalent bond: Very Strong • Polar molecules will orient themselves in a way to maximize these attractions. The strength of these attractions increases with increasing polarity. Polar molecules have higher melting points than non polar ones.

6. Intermolecular Forces: Dipole-Dipole Forces • The magnitude of the melting/boiling temperatures of various substances reflect how strongly the molecules attract one another. • The more strongly the molecules attract, the harder it is to separate them. Hence, the higher the melting/boiling temperatures. • Recall polarity from chapter 8. Any molecule with a net dipole is polar. - δ + δ Cl H Partial negative character Partial positive character

7. Ion-Dipole Interactions • When salts are dissolved in water, the water dipoles are oriented around the ions. • This strong interaction explains why adding salt to water raises its boiling temperature, or why adding salt to snow causes it to melt.

8. Hydrogen Bonding • A special, and very strongtype of dipole-dipole interaction is hydrogen bonding. • Because hydrogen atoms are so small, the partial positive charge on H is highly concentrated. Therefore, it strongly attracts very electronegative elements. • Hydrogen bonds exist only between the H atom in an H—F, H—O, or H—N bond and an adjacent lone electron pair on another F, O, or N atom in another molecule

9. Hydrogen Bonding Causes Abnormalities in Boiling Point Trend

10. Structure and Density of Ice • Water is one of the few compounds that is less densein its solid phase than its liquid phase. • This is due to hydrogen bonding. • In liquid water, 80% of the atoms are H-bonded. In ice, 100% are H-bonded. • To maximize H-bonding, the water molecules in ice spread out. This causes expansion. • Therefore, we have the same mass of water, with a larger volume. Since ρ=(mass/volume), ρ decreases.

11. London Dispersion Forces • With nonpolar molecules, there are no dipoles, so we would not expect to see dipole-dipole interactions. Despite this, intermolecular interactions have still been observed. • For example, nonpolar gases like Helium can be liquified, but how can this happen? What force brings the He atoms together? • Fritz London, a physicist, proposed that the motion of electrons in a nonpolar molecule can create instantaneousdipoles

12. Lets take a Helium atom. At some moment in time, the electrons are spread out within the atom • However, because electrons are constantly moving, electrons can end up on the same side of the atom, creating a charge gradient (instantaneous dipole). This temporary dipole can induce a temporary dipole on another atom, yielding a weak dipole-dipole interaction called a London dispersion force. + + + δ δ δ e- e- e- e- e- e- e- e- e- e- 2+ 2+ 2+ 2+ 2+ - - - δ δ δ

13. London Dispersions • Because London dispersion forces depend on electron motion, the strength of these forces increases with the number of electrons. • The ease of the electron distortion is called polarizability. The more polarizable an atom/molecule, the more likely it is to induce instantaneous dipoles. • Hence, London dispersion forces increase with increasing molar mass because heavier atoms/molecules are more polarizable. All substances have dispersion forces. • In general, for covalently bonded molecules, boiling/melting point increases with molar mass. C5H12 C15H32 C18H38

14. Boiling Points Increase With Increasing Strength of London Dispersion Forces

15. Like Substances Mix • Polar substances are soluble in polar substances. • Nonpolar substances are soluble in nonpolar substances. • Polar and nonpolar substances DO NOT MIX(ex. oil and water) • Dipole-dipole interactions between water molecules are MUCH stronger than the London dispersion forces between oil and water • Water molecules would rather associate with other water molecules than to associate with oil. • Salts are highly soluble in water, but poorly soluble in other most other polar solvents.

16. Which of the following substances mix? • O2 and CH3CH2CH2CH3 (butane) • SO2 and CCl4 • H2O and CH3OH • K2S and H2O • H2O and CH3CH2CH2NH2

17. Summary • We can arrange the intermolecular forces by relative strength: • Coulombic attraction • Ion-dipole • Hydrogen bonding • Dipole-Dipole • London Dispersion * It is important to note that molecules with large nonpolar portions and polar end groups will not mix with polar solvents (ex. CH3CH2CH2CH2CH2OH) because the LD forces will dominate.

18. Example • Arrange the following in order of increasing boiling point: • H2O, Xe, CH4, NaCl (aq), NOCl, CH3OH, CH3COOH • To do this, we must consider the intermolecular forces that exist between the molecules. Those with the strongest forces between them will be the hardest to separate, and thus, have the highest boiling temperature. • DRAW THE LEWIS STRUCTURES (where applicable) • CH4< Xe < NOCl< CH3OH < H2O < CH3COOH, NaCl(aq)