1. 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

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

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

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

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

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

8. * 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.

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

11. * WATER!! * density in the solid state (most dense at about 4 o C) * lake turnover * high heat capacity * coastal climates

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

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

14. * You mix the liquids water, CCl 4, and hexane (CH 3 CH 2 CH 2 CH 2 CH 2 CH 3 ). 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.

16. * Decide what type of intermolecular force is involved in (a) liquid O 2, (b) liquid CH 3 OH, (c) O 2 dissolved in H 2 O. Place the interactions in order of increasing strength.

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

18. * The heat energy required to vaporize a sample is given as the standard molar enthalpy of vaporization, ΔH o vap. (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!

19. * The molar enthalpy of vaporization of methanol, CH 3 OH, is 35.2 kJ/mol at 64.6 o C. How much energy is required to evaporate 1.00 kg of this alcohol?

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

21. * If 0.50 g of pure water is sealed in an evacuated 5.0L flask, and the whole assembly is heated to 60 o C, 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?

22. * Relates equilibrium vapor pressure to the molar enthalpy of vaporization at a specified temperature * You can calculate ΔH vap o for a liquid using the Clausius-Clapeyron equation ln(P 2 /P 1 ) = -ΔH vap o /R[1/T 2 – 1/T 1 ]

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?

24. * When the interface between the liquid and vapor disappears, this point is called the critical point and has a T c and P c. * 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 CO 2 used to decaffeinate coffee

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)

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

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)

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.

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!

30. * Using the ΔH f 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

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

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)

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

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

36. * the melting point is when the lattice collapses and the solid is converted to liquid * melting requires energy; enthalpy of fusion (ΔH fus ) * ionic compounds have higher lattice energies and therefore higher melting points

37. * 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)