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

1. Liquids, Solids, and Van der waals (Intermolecular) Forces Ch 15 Lecture 16 3/27/13 HW: Ch 15: 1,4,13-17

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. Since thermal energy is required to overcome intermolecular forces, we can observe how the phase and temperature of a substance changes as heat is added (constant pressure). • A heating curve for water is shown below, going from -10o C to 125o C Heat of vaporization Heat of fusion Heat rate = 100

5. Energy is Required to Change Phase • The fusion(melting) of water can be represented by: • Therefore, the energy (heat) required to melt n moles of water would be: • The vaporizationof water can be represented by: • The energy (heat) required to vaporize n moles of water would be:

6. Example • Calculate the heat required to heat 28 g of H2O(s) at -10oC to H2O(L) at 50oC, given that the heat capacities of ice and liquid water are 37.7 and 75.3 J/mol K, respectively? Step 1: Raise to melting temp. Step 2: Fusion -10oC 0oC 0oC Step 3: Raise to 50oC 50oC

7. Example • Calculate the heat required to heat 28 g of H2O(s) at -10oC to H2O(L) at 50oC, given that the heat capacities of ice and liquid water are 37.7 and 75.3, respectively? melting ice heating ice heating water

8. Sublimation • Certain substances, like “dry ice” (CO2), convert straight from solid to gas without passing through a liquid phase. This is called sublimation.

9. 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 strong. • Therefore, all ionic compounds have very high melting/boiling points. Cl- Na+

10. Intermolecular Forces: Dipole-Dipole Forces • The values of ΔHvapand ΔHsub reflect how strongly the molecules attract one another in the liquid and solid phases. • The more strongly the molecules attract, the greater the values of ΔH. Hence, substances with stronger attractions have higher boiling/melting points • Recall polarity from chapter 8. Any molecule with a net dipole is polar. - δ + δ Cl H Partial positive character Partial negative character

11. 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.

12. 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

13. 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

14. 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

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

16. 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

17. 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. • Complete H-bonding creates gaps in the crystal structure. This causes the water to expand. • Therefore, we have the same mass of water, with a larger volume. Since ρ=(mass/volume), ρ decreases.

18. Hydrogen Bonding Causes Abnormalities in Boiling Point Trend