Intermolecular Forces, Liquids & Solids

1. CHAPTER 11 AP Chemistry Intermolecular Forces, Liquids & Solids

2. δ- δ+ Bond Polarity, Geometry and Polar Molecules see Chapter 10 pages 409-414 Recall that if the EN between bonded atoms lies in the range of 0.30 to 1.7, we classify the bond as being polar covalent. The shared electrons tend to be shifted towards the atom with the greater EN. We represent this shift with anarrow, pointing towards the more electronegative atom, with a cross-bar at the more electropositive end of the bond. Consider HF: H F =

3. μ = δ x r Polar Molecules The magnitude of this shift is called the dipole moment(μ). The strength of the dipole moment depends primarily on the EN between the bonded atoms, which in turn determines the partial charge, +/- δ, induced in the molecule. The distance between the partial charges, r, is also a factor, so that: Since δ+ = δ- within the molecule, we just use the (+) value for δ. The units for μ are usually expressed in “debye” units (D), named after Peter Debeye, where 1 D = 3.36 x 10-30 C m

4. Polar Molecules For simple two-atom molecules, the greater the EN between the bonded atoms, the greater the magnitude of the dipole moment. Example: compare the dipole moments of HF, HCl, HBr and HI Increasing EN However, for polyatomic molecules, the geometry has a HUGE impact on the net dipole moment for that molecule, so we must be cautious about any generalizations.

5. Polar Molecules Polar molecules behave like very weak ions. They can form attractions between each other, or become aligned in an electric field like ions do:

6. Dipole moments are vector quantities – they have both magnitude and direction (direction of the arrow, ). In polyatomic molecules with more than one dipole moment, we must determine the net dipole moment, by summing up all the dipoles present. If the net dipole is 0 (Σμ = 0) then the molecule is not polar; if Σμ ≠ 0, then the molecule is polar. Determining if a molecule is polar or nonpolar Diatomic molecules bonded by polar covalent bonds are always polar molecules. Diatomic molecules made up of the same 2 elements, such as H2 or F2, are always nonpolar molecules – because there is no dipole moment in the molecule.

7. C O O Polar Molecules Consider CO2 as an example. The dipole moments of each C-O bond are equal in magnitude (identical C-O bonds means identical μ values) but opposite in direction. μ1 – μ2 = 0 Because of the geometry of the molecule, the dipoles cancel out,(Σμ = 0)and no net shift of electrons occurs – the molecule is nonpolar Electron density distribution diagram for carbon dioxide. We see that each C-O bond is polar, with electrons shifted towards the more EN oxygen – however, there is no net charge separation and the molecule is nonpolar.

8. Polar Molecules example: compare the dipoles in NH3 and NF3. Note that both molecules have a net or resultant dipole – the dipoles moments do not cancel. However, the net dipole in NH3 is much greater than in NF3 because the lone pair of electrons has an “effective” dipole that contributes to the net upward dipole in NH3, but partially cancels the downward components of the N-F dipole moments.

9. Cl C Cl Cl Cl Polar Molecules For many molecules, when we draw in the dipoles, it may not be clear if there is a net dipole, or if they all cancel – for example, if the molecule has tetrahedral geometry, some dipoles point out of the page, some into the page, and some in the plane of the page… Consider carbon tetrachloride: Is there a net downward dipole in the molecule CCl4, or do they all cancel? ?

10. Rule of Thumb to Determine if Dipoles Cancel The sum of the dipoles in a molecule are ZERO, (Σμ = 0), and the molecule is thus NONPOLAR if either of the following two conditions apply: • There are NO polar covalent bonds in the molecule. No dipoles mean Σμ = 0. *Note that μ ≈ 0 for C-H bonds. • If dipoles are present but all the elements attached to the central atom are identical AND the geometry of the molecule is ideal (no bond angle adjustments due to lone pairs, etc.), then Σμ = 0. Note that BOTH these conditions must be satisfied. If BOTH Rule 1 AND Rule 2 fail, the molecule is polar!

11. net dipole (Σμ = 1.88 D) polar molecule O O S net dipole (Σμ = 1.60 D) polar molecule H H H O C S S C Cl Cl Cl net dipole (Σμ = 1.01 D) polar molecule example: Which of the following molecules are polar? no net dipole (Σμ = 0) nonpolar molecule

12. Cl Cl C C H H Cl H C C H Cl example:C2H2Cl2 (dichloroethane) has two isomers: (b)trans-dichloroethane. (a)cis-dichloroethane and Determine if either isomer is polar. We see that the dipoles do NOT cancel out in the cis isomer …. cis Σμ = 1.88 D → molecule is polar However, the dipoles DO cancel out in the trans isomer. trans Σμ = 0 → molecule is nonpolar. Note that the trans isomer does not follow my rule of thumb…

13. EN 1.90 0.90 0.70 0.40 1.40 0.40 0.90 1.00 Note that IF the molecule is polar due to its geometry, then IN GENERAL the greater the EN between the bonded atoms, the greater the net dipole for the molecule. Note that there are many exceptions (not shown) to this generalization.

14. Dipole Moments for Polyatomic Molecules with Similar Formulas and Geometries For polyatomic molecules with identical geometries that differ by only one atom, the net dipole moment also tends to increase with increasing EN between bonded atoms, just as was the case with binary molecules – however, compounds involving atoms at the top of the family are often exceptions due to their small size (μ = δ x r )

15. Chapter 11 Intermolecular Attractive Forces

16. Intermolecular Attractive Forces We have been looking at the bonds that holds atoms and ions together within a specific compound. There are also attractive forces that can occur between two separate compounds or molecules. These are called intermolecular attractive forces, or van der Waals forces, after the Dutch chemist Johannes van der Waals, who first made use of them in explaining the boiling points of many compounds.

17. + 2+ F F F H H H H H H F F F The van der Waals forces listed in order of decreasing force are: • ion-dipole forces* • dipole-dipole forces • dipole-induced dipole forces • London (dispersion) forces *technically, ion-dipole is not a van der Waal force, but we include it here anyway.

18. Ion-Dipole Attractions A cation will be strongly attracted to the partial negative end of a polar molecule, and an anion will be strongly attracted to the partial positive end of the polar molecule. polar molecule Recall, this explained why water was able to attach to and “pull apart” ionic compounds in order to dissolve them. polar molecule

19. Dipole-Dipole Attractions The strength of the ion-dipole attraction depends on the charge and the size of the ion, and on the magnitude of the dipole moment of the polar molecule. In aqueous solutions, cations are typically surrounded by six water molecules in an octahedral arrangement. + 2+

20. Dipole-Dipole Attractions The partial positive end of one polar molecule will be attracted to the partial negative end of another polar molecule. This is called a dipole-dipole attraction. Because of these attractive forces, polar molecules tend to have higher melting and boiling points.

21. Hydrogen Bonding Molecules that contain H bonded to very small, highly electronegative atoms like N, O, or F form very strongly polar molecules. The dipole-dipole attractions between the H of one molecule and the N, O, or F of the other molecule are nearly as strong as ion-ion attractions. We call this strong dipole-dipole attraction a hydrogen bond.

22. ? ? ? Hydrogen Bonding We can see just how strong the H-bond is by comparing the boiling points of compounds that can form H-bonds with similar compounds that cannot. H2O can form 3 H-bonds per molecule, so even though its μ is smaller than that of HF, water’s boiling point is still higher Note that by the trends in the graph, we would expect the boiling points of H2O, NH3 and HF to be MUCH lower than what they actually are!

23. Hydrogen Bonding Hydrogen bonding is partly responsible for maintaining the “twist” in the alpha helix of DNA molecules. It is also the reason that ice forms a rigid crystalline structure, and helps explain why alcohols are soluble in water, among other things.

24. Induced Dipoles If a charged particle, such as an ion, or even a highly polar molecule, is brought close to another particle with no charge, the charged particle can induce charge separation (i.e., a dipole) on the neutral particle. ion –induced dipole: if an ion is used, the attractive force is stronger and is called an ion-induced dipole attraction. dipole –induced dipole: if a polar molecule is used, the attraction is weaker and is called a dipole-induced dipole attraction.

25.   Induced dipoles The repulsion of the excess e- on the δ- polar molecule (or an anion) repels the e- on the neutral atom. This induces a (+) charge on the front of the atom, and the two particles now attract each other. - + Alternately, a cation (or the δ+ end of a polar molecule) can attract the electrons of another atom, inducing a charge on the neutral atom, and the two now attract each other. + -

26. Polarizability The polarizability of an atom is a measure of how easily a dipole can be induced in that atom. The more electrons an atom has, the more diffuse the electron cloud is, and the more easily distorted the electron cloud is -- hence, the more polarizable the atom. This means larger atoms (more shells, less strongly held electrons) are more polarizable than smaller atoms.

27. Dispersion Forces There are very weak attractions, even between non-polar molecules. These attractions are called London forces (also known as dispersion forces). London forces can occur between any two atoms, and so they form whether the molecules are polar or nonpolar. London forces are, however, the ONLY type of intermolecular force that can occur between two nonpolar molecules.

28.   London (dispersion) Forces Due to the random changes in an atom’s electron density, for one very brief instant, most of the electrons may end up on one side of the atom. This makes the atom “polarized” for that one brief instant. We call it an instantaneous dipole. This instantaneous dipole can now induce a dipole on another atom, much like the dipole-induced dipole attractive force was formed. - + induced dipole instantaneous dipole

29.  London (dispersion) Forces The attraction between the two atoms lasts for only an instant (after which the electrons re-arrange themselves). However, the electrons in atoms move so fast that a new polar atom is created hundreds of times every microsecond, creating multiple opportunities to form an attraction to another atom. + If it weren’t for London forces, all nonpolar molecules would boil at temperatures near 0 K! +

30. H H H H H H C C H H H H H H H H H H H H H H H H C C C C C C H H H C C C C C H H C C C C C H H H H H H H H H H H H H H H H H C C H H H H London (dispersion) Forces Molecules with more “exposed sites” for London forces to act upon increases the likelihood of forming those London attractions. Thus, long, straight chain molecules form more London attractions than branched molecules: Multiple contact points where London attractions can form Only one contact point where London attractions can form

31. Boiling Point Trends and Intermolecular Forces • The stronger the net dipole moment, the stronger the attraction between polar molecules and the higher the boilint point. • The more polarizable the atom, the stronger the induced dipole and the stronger the attraction between two molecules, and the higher the boiling pt. • More massive molecules tend to have higher boiling pts as well, since they have greater inertia that must be overcome to move them apart. 4. Molecules with more “exposed sites” for London forces to act on increases the likelihood of forming London attractions, which increases the boiling point.

32. LIQUIDS

33. Properties of Liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. Strong intermolecular forces High surface tension

34. Cohesion is greater Adhesion is greater Properties of Liquids Capillary Action: A result of surface tension Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules a) adhesive forces are greater than cohesive forces: liquid rises. b) cohesive forces are greater than adhesive forces: liquid does not rise Hg & glass H2O & glass (a) (b)

35. Properties of Liquids Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity

36. SOLIDS

37. lattice point A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. An amorphoussolid does not possess a well-defined arrangement and long-range molecular order. A unit cell is the basic repeating structural unit of a crystalline solid. • At lattice points: • Atoms • Molecules • Ions Unit cells in 3 dimensions Unit Cell

38. The coordination numberis the number of atoms (or ions) surrounding an atom (ion) in a crystal lattice. The larger the coordination number, the more tightly packed the atoms are within the crystal. each Na+ ion in NaCl is surrounded by 6 Cl– ions; thus Na has a coordination number of 6.

41. Closest Packing Closest packing is the most efficient arrangement of spheres. There are several ways of arranging atoms to obtain closest packing:

42. Types of Crystals • Ionic Crystals • Lattice points occupied by cations and anions • Held together by electrostatic attraction • Hard, brittle, high melting point • Poor conductor of heat and electricity CsCl ZnS CaF2

43. Types of Crystals • Covalent Crystals • Lattice points occupied by atoms • Held together by covalent bonds • Hard, high melting point • Poor conductor of heat and electricity carbon atoms graphite diamond

44. Types of Crystals Molecular Crystals • Lattice points occupied by molecules • Held together by intermolecular forces • Soft, low melting point • Poor conductor of heat and • electricity

45. For most substances, the volume of the solid phase is LESS than the volume of the liquid phase. Thus, for most substances, the density of the solid is greater than the density of the liquid. When most liquids freeze, the solid form sinks to the bottom. solid benzene sinks in liquid benzene but ice floats in water

46. Water is an important exception. The density of ice is LESS than the density of liquid water, so it floats. This is because the structure of ice is an “open lattice” composed of 6-membered “rings” which increases its volume max density of ice is at 4°C

47. nucleus & inner shell e- mobile “sea” of delocalized e- Types of Crystals • Metallic Crystals • Lattice points occupied by metal atoms • Held together by metallic bonds • Soft to hard, low to high melting point • Good conductors of heat and electricity Cross Section of a Metallic Crystal

48. Crystal Structures of Metals

49. A Comparison of Crystal Types and Their Properties

50. Amorphous Solids An amorphoussolid does not possess a well-defined arrangement and long-range molecular order. A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Non-crystalline quartz glass Crystalline quartz (SiO2)