Liquids and Intermolecular Forces (1)

1. Liquids and Intermolecular Forces (1) Lecture 15

3. States of Matter The fundamental difference between states of matter is the distance between particles. The distance between particles decreases when going from gas to liquid to solid.

4. Solids, Liquids, and Gases

5. Condensed Phases Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

7. The States of Matter and Intermolecular Forces • The state of a substance at a particular temperature and pressure depends on two antagonistic factors: • The kinetic energy of the particles. • The strength of the attractions between the particles.

8. Electronegativity and Bond Polarity

10. Electronegativity • The electronegativity value, as pertinent to scope discussed here: • indicates the attraction of an atom for shared electrons. • increases from left to right going across a period on the periodic table. • is low for the metals (metals are usually electropositive and form ionic bonds). • Through a period, bond polarity increases towards right nonmetal as the distance between nonmetals increases. • Through a group, bond polarity increases towards upper nonmetal as distance between the two nonmetals increases .

11. Nonpolar Covalent Bonds • A nonpolar covalent bond: • occurs between nonmetals. • is an equal or almost equal sharing of electrons. • has almost no electronegativity difference (0.0 to 0.4). • Examples: • Electronegativity • Atoms Difference Type of Bond • N-N3.0 - 3.0 = 0.0 Nonpolar covalent • Cl-Br3.0 - 2.8 = 0.2 Nonpolar covalent • C-H 2.1 - 2.0 = 0.1 Nonpolar covalent

12. Polar Covalent Bonds • A polar covalent bond: • occurs between nonmetal atoms. • is an unequal sharing of electrons. • has a moderate electronegativity difference (0.5 to 1.7). • Examples: • Electronegativity • Atoms Difference Type of Bond • P-Cl3.0 - 2.1 = 0.9 Polar covalent • C-F4.0 - 2.5 = 1.5 Polar covalent • O-S 3.5 - 2.5 = 1.0 Polar covalent • X-H >1 Polar covalent

13. Comparing Nonpolar and Polar Covalent Bonds

14. Predicting Bond Types

15. Intermolecular Forces The properties of liquids are intermediate between those of gases and solids but are more similar to solids. In contrast to intramolecular forces, such as the covalent bonds that hold atoms together in molecules and polyatomic ions, intermolecular forces hold molecules together in a liquid or solid. Intermolecular forces are generally much weaker than covalent bonds

16. Intermolecular Vs Intramolecular (Chemical Bonds) forces • Intermolecular forces are much weaker than intramolecular forces. . For example, it requires 927 kJ to overcome the intramolecular forces and break both O–H bonds in 1 mol of water, but it takes only about 41 kJ to overcome the intermolecular attractions and convert 1 mol of liquid water to water vapor at 100°C.

17. Intermolecular Forces The attractions between molecules are much weaker than the intramolecular attractions that hold compounds together.

18. Intermolecular Forces These intermolecular attractions are, however, strong enough to control physical properties, such as boiling and melting points, vapor pressures, and viscosities.

19. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces.

20. van der Waals Forces Three types of intermolecular forces exist between neutral molecules, these are collectively known as van der Waals forces: • Dipole–dipole interactions • London dispersion forces • Hydrogen bonding But not ion-dipole forces or ionic bonding

21. London Dispersion Forces Why are many nonpolar molecules, such as bromine, benzene, and hexane, liquids at room temperature, while others, such as iodine and naphthalene, are solids?. What kind of attractive forces can exist between nonpolar molecules or atoms? This question was answered by Fritz London in 1930, where London proposed that temporary fluctuations in the electron distributions within atoms in nonpolar molecules could result in the formation of short-lived instantaneous dipole moments, which produce attractive forces called London dispersion forces between otherwise nonpolar substances.

22. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

23. Dispersion forces, charge distribution for a pair of helium atoms at three instants.

24. London Dispersion Forces • These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called polarizability. • Polarizability increases as the number of electrons in atom increases (or increasing molecular weight)

25. Factors Affecting London Forces • The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane) tend to have stronger dispersion forces than short, fat ones (like neopentane). This is due to the increased surface area in n-pentane.

26. London Forces and Molecular Size

27. London Forces and Atomic Size • Larger atoms have electron clouds that are easier to polarize.

28. London Forces and Molecular mass • The strength of dispersion forces tends to increase with increased molecular weight. Boiling points increase as the formula weight is increased, due to stronger dispersion forces

29. Boiling points increase as the molecular mass is increased, due to stronger dispersion forces

30. Dipole–Dipole Interactions • As a result of differences of electronegativities of atoms connected by covalent bonds, a dipole is formed. • Molecules that have permanent dipoles are attracted to each other. • The positive end of one dipole is attracted to the negative end of the other, and vice versa. • These forces are only important when the molecules are close to each other.

33. In the solid, CH3CN molecules are arranged with the negatively charged nitrogen end of each molecule close to the positively charged -CH3 ends of its neighbors. In the liquid, the molecules are free to move with respect to one another, and their arrangement becomes more disordered. At any moment both repulsive and attraction forces exist.

34. Dipole–Dipole Interactions For molecules of approximately equal mass and size, the more polar the molecule, the higher its boiling point.

35. Which has a higher boiling point: ethyl methyl ether (CH3OCH2CH3), or 2-methylpropane [isobutane, (CH3)2CHCH3],. The two compounds have essentially the same molar mass (58–60 g/mol), therefore, they should have similar London forces, and thus we must look at differences in polarity to predict the strength of the intermolecular dipole–dipole interactions and thus the boiling points of the compounds. The first compound, 2-methylpropane, contains only C–H bonds, which are not very polar because C and H have similar electronegativities, and thus will have a lower boiling point than ethyl methyl ether.

36. Relationships between the Dipole Moment and the Boiling Point for Organic Compounds of Similar Molar Mass

37. Which Have a Greater Effect, Dipole–Dipole Interactions or Dispersion Forces? • If two molecules are of comparable size and shape, dipole–dipole interactions will likely be the dominating force. • If one molecule is much larger than the other, dispersion forces will likely determine its physical properties.

38. How Do We Explain This? • The nonpolar series (SnH4 to CH4) follow the expected trend. • The polar series follow the trend until you get to the smallest molecules in each group, where these have boiling points much higher than expected!!!

39. Hydrogen Bonding • Hydrogen bonding typically occurs when a hydrogen atom bonded to O, N, or F, is electrostatically attracted to a lone pair of electrons on an O, N, or F atom in another molecule.

40. Hydrogen Bonding • The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong. • We call these interactions hydrogen bonds.

41. Hydrogen Bonding • Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

42. Hydrogen bonding inacetic acid and 1-propanol. The greaterthe number of hydrogen bonds possible, themore tightly the molecules are held togetherand, therefore, the higher the boiling point.

43. Ion–Dipole Interactions • Ion–dipole interactions (a fourth type of intermolecular forces) are important in solutions of ions. • The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

45. Examples: Which of the following pairs has a higher boiling point? CH3OH or CH3CH3OH H2O or H2S Br2 or ICl CHCl3 or CHBr3 CH3CH2OH or HOCH2CH2OH

46. Which is greater: dipole dipole or dispersion forces? CH3F has a dipole moment of 1.8 Debye, while CCl4 has no dipole moment. The boiling point of the two compounds are -78.4 and 76.5 oC. Although CH3F has both dipole – dipole forces and dispersion forces, it has a much lower BP (smaller intramolecular forces) than CCl4 which has dispersion forces only. However this is not always true.

47. What types of intermolecular forces exists between the following pairs: a. HBr and Na+ b. Cl2 and CHBr3 c. H2O and NO3- d. NH3 and HCl e. KBr and H2O • HF and H2O

48. Which one of the following should have the lowest boiling point? A) CH3OH B) H2S C) NH3 D) HCl E) CH4

49. Of the following substances, __________ has the highest boiling point. A) CH3CH2OH B) C2H6 C) N2 D) F2 E) HOCH2CH2OH

50. Of the following, __________ has the highest boiling point. A) N2 B) Br2 C) H2 D) Cl2 E) O2