UNIT 2

1. UNIT 2 Inorganic Nomenclature*, Intermolecular Forces, and Properties of Solutions *Students are responsible for reviewing nomenclature on their own.

2. Attention Chem 1215 Students: Part of your nomenclature requirement is to know the correct chemical name and formula of every chemical you use in the lab!

3. States of Matter • Gases • assume the volume AND shape of the container • are compressible • flow readily • Diffusion within a gas occurs rapidly.

4. States of Matter • Liquids • assume the shape of the part of the container they occupy but do not expand to fill the container • are virtually incompressible • flow readily • Diffusion within a liquid occurs slowly. The atoms or molecules of the liquid are in constant motion.

5. States of Matter Atoms in a solid NEVER sit still…they vibrate in place constantly. The higher the temperature, the more the atoms move about their position. • Solids • retain their own volume and shape • are virtually incompressible • do not flow • Diffusion within a solid occurs extremely slowly.

6. States of Matter • The state of a substance depends largely on the balance between • the kinetic energies of the particles (kinetic energy is also called thermal energy or energy of motion), and • the energies of attraction between particles (intermolecular attractions).

7. States of Matter • Gases: Kinetic energies are much larger than intermolecular attractions. • Liquids: Intermolecular attractions are significant, but only strong enough to impose short-range order. • Solids: Intermolecular attractions are strong enough to hold the molecules in place… • …but solids at T > 0 K still have energy of motion and will vibrate in place.

8. ATOMS DO NOT SIT STILL… EVER!

9. Intermolecular Forces • arise from the charged nature of the subatomic particles (electrons and protons). • are responsible for the cohesiveness of materials. • are what determine physical properties of pure substances such as melting point, boiling point, and vapor pressure.

10. Intermolecular Forces • are in general much weaker than intramolecular forces (aka: bonds). intermolecular forces are generally less than 15% as strong as ionic or covalent bonds

11. Intermolecular Forces • Substances that are gases at room temperature have weak intermolecular forces. • Substances that are condensed (liquids or solids) at room temperature have much stronger intermolecular forces. • If intermolecular forces did not exist, all substances would be gases, even at extremely low temperatures.

12. What are intermolecular forces? • Ion-dipole (strongest) • hydrogen bonding • dipole-dipole • dispersion (weakest)

13. Types of Intermolecular Forces • Clearly, to understand the various intermolecular forces, you need to be able to identify a dipole. • In chemistry, a dipole is a molecule containing a partial separation of charge (as opposed to an ionic compound, in which the charge separation is complete). • All polar molecules contain one or more dipoles.

14. What is a Polar Molecule? • For a molecule to be polar, it must have polar bonds the dipole moments of which do not completely cancel each other. • Review section 9.3 and 8.4 of your text, and study second half of “Review of Lewis Structures” on the Chem II web site.

15. Determining the Polarity of a Molecule • no polar bonds: molecule is nonpolar • 1 polar bond: molecule is polar • 2 or more polar bonds: polarity is a function of the geometry of the polar bonds in the molecule. • First we will see how to determine if a bond is polar.

16. Electronegativity and Bond Polarity Ionic and covalent bonds are the extremes: complete control of the valence electrons and complete sharing of the valence electrons. Most bonds are somewhere in between. Electronegativity is the ability of an atom in a bond to attract electrons. Atoms that are more electronegative will tend to have a partial negative charge. In a bond, a partial separation of charge means the covalent bond is polar.

17. Electronegativities can be found in your text. If the difference in electronegativities of the two atoms is between 0.4 and 2.0, the bond is polar (aka polar covalent). If the difference in electronegativities of the two atoms is ≤0.4, the bond is nonpolar.

18. Bond Polarity and Dipole Moments A bond is polar if it has a significant dipole moment. Dipole moment depends on both the electronegativities of the atoms in the bond, and on charge separation. Dipole moment (μ) μ = δ d δ is the amount of partial charged is the bond length

19. Identifying Polar Bonds To a first approximation, look at the periodic table. The most electronegative elements are in the upper right hand corner (disregarding the noble gases). So, C or H bound to these elements will form polar bonds. O—H C≡N C—O H—F C=O N—H

20. Determining the Polarity of a Molecule…Now That You Can Identify a Polar Bond • no polar bonds: molecule is nonpolar • All hydrocarbons are nonpolar. • 1 polar bond: molecule is polar • 2 or more polar bonds: polarity is a function of the geometry of the polar bonds in the molecule.

21. Water is Polar. CO2 is not. Water has 2 polar bonds and the shape of the molecule is nonlinear (bent). CO2 has 2 polar bonds and the shape of the molecule is linear, so the dipoles cancel.

22. Intermolecular forces: Ion-Dipole occur between an ion and a polar molecule

23. Intermolecular forces: Hydrogen bonding occurs between an electropositive (+) H and an electronegative (-) element (usually O, N, F).

24. Intermolecular forces: Dipole-Dipole Attractions occur between electropositive (+) elements and electronegative (-) elements involved in polar bonds.

25. Intermolecular forces: Dispersion (London) Forces occur in nonpolar compounds (actually, they occur in all compounds). They have electrons, too, and are capable of transient dipoles.

26. Intermolecular Forces: Dispersion • Dispersion forces or van der Waals forces arise from the fact that the electrons around an atom are constantly shifting and can - briefly - make it a dipole. This can then induce a dipole on a nearby atom. • This attraction is also called an instantaneous dipole - induced dipole attraction.

27. Types of Intermolecular Forces: Dispersion • Dispersion forces are present between all moleculesand account for why nonpolar substances can be liquefied. • These forces increase with increasing surface area and/or molecular weight. • These forces can be significant even between polar molecules such as HCl.

28. Dispersion Forces Dispersion forces increase with increasing molecular weight and surface area.

29. Hydrogen Bonding Hydrogen bonding is the interaction between a H atom in a polar bond (particularly H-F, H-O, or H-N) and an unshared pair of electrons on a nearby electronegative ion or atom (usually F, O, or N).

30. Hydrogen Bonding Hydrogen bonding is the reason water has an abnormally high boiling point.

31. Hydrogen Bonding Hydrogen bonding is the reason ice is less dense than water.

32. Hydrogen Bonding Hydrogen bonding is what holds the two strands of a DNA or RNA molecule together.

33. Summary of Intermolecular Forces Examples Ar(l), I2(s) H2S, CH3Cl H2O(l), H2O(s), HF, NH3 NaCl in water weakest Dispersion forces (van der Waals forces) Dipole - dipole interactions Hydrogen bonding (a special instance of the dipole-dipole interaction involving an H in a polar bond and O, F, or N in another polar bond) Ion-dipole strongest