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

1. Liquids, Solids, and Van der waals (Intermolecular) Forces Ch 15 HW: Ch 15: 1,5,13-17 SuggeSTED: 4, 9, 11

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 very strong. Coulombic attractions are the strongest of all intermolecular forces. • 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. • Recall polarity from chapter 8. Any molecule with a net dipole is polar. + δ Cl - H δ Partial negative character Partial positive 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