Intermolecular Forces and Liquids and the Chemistry of Solids

1. Intermolecular Forces and Liquids and the Chemistry of Solids Chapters 12 and 13

2. More complicated than gases… • particles are close together due to attractive forces • these attractive forces are mostly ignored when dealing with gases 12.1 State of Matter & IMF

3. directly related to properties such as melting point and boiling point (energy needed to overcome attractions) • solubility of gases, liquids, and solids in various solvents • determines structures of biomolecules such as DNA and proteins Intermolecular Forces

4. Forces between particles • in ionic compounds, the ions are held together by electrostatic attraction • in molecular compounds, the intermolecular forces (forces between particles) are based on electrostatic attractions that are weaker than ionic forces Kinetic-Molecular Theory

5. Polar Molecules mixed with Ionic Compounds • Ions will be attracted to polar ends of molecules. • Can be used to determine the enthalpy of solvation (or hydration) • Ion dipole forces 12.2 IMF Involving Polar Molecules

6. Polar Molecules mixed with Ionic Compounds (cont.) • The force depends on: • distance between ion and dipole. • charge on ion. • magnitude of dipole. Polar Molecules

7. Molecules with Permanent Dipoles • molecules with dipoles interact by dipole-dipole attraction • these attractions influence endothermic evaporation (ΔHvap) and exothermic condensation • polar bonds are stronger and require more energy to break bonds Polar Molecules

8. Intermolecular Forces • Solubility • “like dissolves like” means that polar molecules are more likely to dissolve in polar solvents, and nonpolar molecules are more likely to dissolve in nonpolar solvents

9. A hydrogen bond is an attraction between the hydrogen atom of an X-H bond and Y, where X and Y are atoms of highly electronegative elements and Y has a lone pair of electrons. • These bonds are an extreme form of dipole-dipole in which one atom is always H and the other is often O, N, or F. Hydrogen Bonding

10. due to large electronegativity differences, these hydrogen bonds are very polar; partial charges are formed • hydrogen atom becomes a bridge between electronegative elements Hydrogen Bonding

11. WATER!! • density in the solid state (most dense at about 4oC) • lake turnover • high heat capacity • coastal climates Hydrogen Bonding

12. Nonpolar Molecules • polar molecules (like water) can induce a dipole in molecules without a permanent dipole (such as oxygen gas and water) • the force of attraction is called a dipole/induced dipole interaction • The process of inducing a dipole is called polarization (a molecule/atom has a certain polarizability) • The higher the molar mass, the larger the cloud and greater the polarizability of the molecule 12.3 IMF involving Nonpolar Molecules

13. London Dispersion Forces • when two nonpolar atoms approach each other, attractions or repulsions between their electrons and lead to distortions in their clouds; leading to intermolecular attraction • Arise between all molecules • this force of attraction in nonpolar molecules is an induced dipole/induced dipole force, or London dispersion forces Intermolecular Forces

14. You mix the liquids water, CCl4, and hexane (CH3CH2CH2CH2CH2CH3). For each pair of compounds, what type of intermolecular forces can exist between the compounds? If you mix these three liquids, describe what observations you might make. Practice Problem

16. Decide what type of intermolecular force is involved in (a) liquid O2, (b) liquid CH3OH, (c) O2 dissolved in H2O. Place the interactions in order of increasing strength. Practice Problem

17. particles interact with each other like a solid, but there is little order in their arrangement • vaporization or evaporation is the process in which a liquid becomes a gas; molecules escape the liquid surface and enter the gaseous state Properties of Liquids - Vaporization

18. The heat energy required to vaporize a sample is given as the standard molar enthalpy of vaporization, ΔHovap. (kJ/mol) • The opposite process is condensation, in which a molecule may reenter the liquid phase. This releases energy, which is why a steam burn is much worse than one from boiling water! Vaporization

19. The molar enthalpy of vaporization of methanol, CH3OH, is 35.2 kJ/mol at 64.6oC. How much energy is required to evaporate 1.00 kg of this alcohol? Practice Problem

20. liquid in a sealed flask will form a dynamic equilibrium • when this equilibrium has been established, the pressure exerted by the vapor is the equilibrium vapor pressure (a measure of the tendency of molecules to escape to the vapor phase) • volatility Vapor Pressure

21. If 0.50 g of pure water is sealed in an evacuated 5.0L flask, and the whole assembly is heated to 60oC, will the pressure be equal to or less than the equilibrium vapor pressure of water at this temperature? What if you use 2.0 g of water? Under either set of conditions is any liquid water left in the flask, or does it all evaporate? Practice Problem

22. Relates equilibrium vapor pressure to the molar enthalpy of vaporization at a specified temperature • You can calculate ΔHvapofor a liquid using the Clausius-Clapeyron equation ln(P2/P1) = -ΔHvapo/R[1/T2 – 1/T1] Clausius-Clapeyron

23. boiling point is the temperature at which a liquid’s vapor pressure is equal to the external pressure; at standard pressure this point is called the normal boiling point • Water boils at a lower temperature at higher altitudes…why? Boiling Point

24. When the interface between the liquid and vapor disappears, this point is called the critical point and has a Tc and Pc. • At this point, the substance is called a supercritical fluid, meaning that it is like a gas under high pressure so that its density is like a liquid but its viscosity is like a gas. • supercritical CO2 used to decaffeinate coffee Critical Temperature and Pressure

25. surface molecules are attracted to molecules below them, makes the liquid behave as if it had a skin – toughness of that “skin” is surface tension (water striders) • molecules may be attracted to adhesive forces between two different substances in such a way that overcomes the cohesive forces between the molecules themselves – capillary action (paper) • viscosity – resistance to flow (honey vs. water) Surface Tension, Capillary Action, and Viscosity

26. After reading sections 12.1-12.4, you should be able to do the following… • P. 581 (2-24 even) Homework

27. molecules, atoms, or ions cannot move (although they vibrate or rotate some) • solids have regular, repeating patterns of atoms or molecules within the structure • attractive forces are maximized and repulsive forces are minimized • the unit cell within a crystalline solid is the smallest repeating unit (such as a “repeat” in a wallpaper pattern) 13.1 Crystal Lattice and Unit Cells

28. a crystal lattice refers to a bunch of unit cells all put together • The lattice points defining each unit cell in solids represent identical environments for the ions, atoms, or molecules. Solid State

29. Ionic compounds have high melting points due to the strength of bonding in the lattice. • Lattice energy is a measure of the strength of ionic bonding • Measured as lattice enthalpy! 13.3 Lattice Energy

30. Using the ΔHf of compounds and enthalpy for ion formation in gas phase, you can calculate the lattice enthalpy of a compound. • Add up the steps: • Formation of solid sodium chloride = formation of each element as a gas plus the formation of ions plus the lattice enthalpy Born-Haber Cycle

31. 13.4 Other Kinds of Solid Materials

32. low vapor pressure due to strong interactions of positive and negative ions • brittle due to repulsion of like charges caused when one layer slides across another layer • do not conduct electricity unless melted or in solution • do not dissolve in nonpolar solvents Ionic Solids

33. positive kernels consisting of nucleus and inner electrons surrounded by a sea of mobile valence electrons • good conductors • malleable and ductile • pure substances or mixtures (alloys) • interstitial alloys (steel) • substitutional alloys (brass) Metallic Solids

34. Only formed from nonmetals: elemental (diamond, graphite) or two nonmetals (silicon dioxide, silicon carbide) • Covalent network solids have high melting points • Generally for in the carbon group due to their ability to form four covalent bonds • Graphite is an allotrope of carbon that forms sheets; high melting point due to covalent bonds, but soft due to LDF layers Covalent Solids

35. Nonmetals, diatomic elements, two or more nonmetals • Nonconductors • Low melting points due to weak IMF • Sometimes very large molecules or polymers Molecular Solids

36. the melting point is when the lattice collapses and the solid is converted to liquid • melting requires energy; enthalpy of fusion (ΔHfus) • ionic compounds have higher lattice energies and therefore higher melting points 13.5 Phase Changes Involving Solids

37. Solids – Physical Properties • Molecules can escape directly from the solid to the gas phase by sublimation, which is endothermic. • frost

38. After reading sections 13.1-13.6, you should be able to do the following… • P. 611 (14-18 even, 22) Homework