Chapter 11 Liquids and Intermolecular Forces MCQ 3 out of 60 FRQ occasionally

1. Lecture Presentation Chapter 11 Liquids and Intermolecular ForcesMCQ 3 out of 60 FRQ occasionally James F. Kirby Quinnipiac University Hamden, CT

2. 11.1 States of Matter • The fundamental difference between states of matter is the strength of the intermolecular forces of attraction. • Stronger forces bring molecules closer together. • Solids and liquids are referred to as the condensed phases(they are close together).

3. Differences in theStates of Matter • The state a substance in at a particular temperature and pressure depends on two antagonistic entities: • The kinetic energy of the particles; • The strength of the attractions between the particles

4. For a given substance, do you expect the density of the substance in its liquid state to be closer to the density in the gaseous state or in the solid state? • It is closer to the density of a gas. • b. It is closer to the density of a solid. • Both are somewhat close to one another in density. • d. They are significantly different in density.

5. Which State of Matter? • The answer to this question largely relies on the • balance between the kinetic energies of the particles. • interparticle energies of attraction.

6. 11.2 Intermolecular Forces • The attractions between molecules are not nearly as strong (15%) as the intramolecular attractions (bonds) that hold compounds together. • Many physical properties reflect intermolecular forces, like boiling points, melting points, viscosity, surface tension, and capillary action.

7. These intermolecular forces between neutral molecules are referred to as van der Waals forces

8. How would you expect the H—Cl distance represented by the red dotted line to compare with the H—Cl distance within the HClmolecule? • Smaller distance • b. Larger distance • c. Same distance

9. When water boils, what are the bubbles composed of? Oxygen Water Carbondioxide Otherimpurities

10. Types of Intermolecular Force • Weakest to strongest forces:

11. 1) London Dispersion Forces • The figure below shows how a nonpolar particle (in this case a helium atom) can be temporarily polarized to allow dispersion force to form. • The tendency of an electron cloud to distort is called its polarizability. (LMPEC)

15. Factors Affecting London Forces • number of electrons in an atom (more electrons, more dispersion force) • size of atom or molecule/molecular weight. Larger atoms have larger e- clouds which are easier to polarize (LMPEC). • shape of molecules : long, skinny molecules tend to have stronger dispersion forces than short, fat ones.

16. Polarizability & Boiling Point • If something is easier to polarize, it has a lower boiling point. • Remember: This means less intermolecular force (smaller molecule: lower molecular weight, fewer electrons).

17. Why is the boiling point of the halogen in each period greater than the noble gas? • The halogen gas molecules are closer together than the noble gas molecules. • b. The noble gas dispersion forces are greater. • The halogen molecule dispersion forces are greater. • d. The halogen molecule dipole forces are greater.

18. List the substances CCl4, CBr4, and CH4in order of increasing boiling point. CH4 < CBr4 < CCl4 CCl4< CH4 < CBr4 CH4< CCl4 < CBr4 CBr4< CCl4 < CH4

19. 2) Dipole–Dipole Interactions • Polar molecules have a permanent positive and a permanent negative end–a dipole (two poles, δ+ and δ−). • The oppositely charged ends attract each other. These forces are only important when the molecules are close to each other

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

21. Moving from left to right, do the dispersion forces get stronger, get weaker, or stay roughly the same in the molecules shown here? • Stronger • b. Weaker • c. About the same

22. 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 another, dispersion forces will likely determine its physical properties.

23. How Do We Explain This? • In a group, the period 3/4/5 elements have higher boiling points as the group member gets larger. • What happens with the period 2 elements? For group 4A, the trend is continued. What about for the other groups? Water, HF, and NH3 are quite an anomaly

24. Why is the boiling point of SnH4 higher than that of CH4? • SnH4 is more polar than CH4. • b. SnH4is smaller in size than CH4. • SnH4has greater internal dispersion forces than in CH4. • d. SnH4 is ionic in structure and CH4 is molecular.

25. 3) Hydrogen Bonding • The dipole–dipole interactions experienced when H is bonded to N, O, or F are unusually strong. - Is it strong H-NOF (enough) to be hydrogen bonding? • We call these interactions hydrogen bonds. • A hydrogen bond is an attraction between a hydrogen atom attached to a highly electronegative atom and a nearby small electronegative atom in another molecule or chemical group.

26. To form a hydrogen bond what must the non-hydrogen atom (N, O, or F) involved in the bond possess? • The non-hydrogen atom must have a nonbonding electron pair. • b. The non-hydrogen atom must have low electronegativity. • The non-hydrogen atom must have a large atomic size. • d. The non-hydrogen atom must have a small electron affinity.

27. What Forms Hydrogen Bonds? • Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. • These atoms interact with a nearly bare the H nucleus (which contains one proton).

29. Sample Exercise 11.1Identifying Substances That Can Form Hydrogen Bonds In which of these substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH4), hydrazine (H2NNH2), methyl fluoride (CH3F), hydrogen sulfide (H2S)? Solution Analyze We are given the chemical formulas of four compounds and asked to predict whether they can participate in hydrogen bonding. All the compounds contain H, but hydrogen bonding usually occurs only when the hydrogen is covalently bonded to N, O, or F. Plan We analyze each formula to see if it contains N, O, or F directly bonded to H. There also needs to be a nonbonding pair of electrons on an electronegative atom (usually N, O, or F) in a nearby molecule, which can be revealed by drawing the Lewis structure for the molecule. Solve The foregoing criteria eliminate CH4 and H2S, which do not contain H bonded to N, O, or F. They also eliminate CH3F, whose Lewis structure shows a central C atom surrounded by three H atoms and an F atom. (Carbon always forms four bonds, whereas hydrogen and fluorine form one each.) Because the molecule contains a C—F bond and not an H—F bond, it does not form hydrogen bonds. In H2NNH2, however, we find N—H bonds, and the Lewis structure shows a nonbonding pair of electrons on each N atom, telling us hydrogen bonds can exist between the molecules:

30. Sample Exercise 11.1Identifying Substances That Can Form Hydrogen Bonds Continued Check Although we can generally identify substances that participate in hydrogen bonding based on their containing N, O, or F covalently bonded to H, drawing the Lewis structure for the interaction provides a way to check the prediction. Practice Exercise 1 Which of the following substances is most likely to be a liquid at room temperature? (a) Formaldehyde, H2CO; (b) fluoromethane, CH3F; (c)hydrogen cyanide, HCN; (d) hydrogen peroxide, H2O2; (e) hydrogen sulfide, H2S. Practice Exercise 2 In which of these substances is significant hydrogen bonding possible: methylene chloride (CH2Cl2), phosphine (PH3), chloramine (NH2Cl), acetone (CH3COCH3)?

31. What is the approximate H—O••••H bond angle in ice, where H—O is the covalent bond and O••••H is the hydrogen bond? • Approximately 90° • b. Approximately 105° • Approximately 120° • d. Approximately 180°

32. What major type of attractive interaction must be overcome for water to evaporate? Hydrogenbonding Dispersion forces Covalent bonds Dipole-dipole forces

33. 4) Ion–Dipole Interactions • Ion–dipole interactions are found in solutions of ions. • The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

34. Why does the O side of H2O point toward the Na+ ion? • The H atoms repel the Na+ion. • b. The H atoms are attracted to other H2O molecules. • The angle of the HOH bond requires this orientation. • d. The negatively charged O is attracted to the positive Na+.

35. In which mixture do you expect to find ion dipole forces between solute and solvent? CH3OH in water or Ca(NO3)2in water. CH3OH in water, because CH3OH is a strongelectrolyte and forms ions. Ca(NO3)2 in water, because Ca(NO3)2is a strongelectrolyte and forms ions. CH3OH in water, because CH3OH is a weakelectrolyte and forms ions. Ca(NO3)2 in water, because Ca(NO3)2is a weakelectrolyte and forms ions.

36. Summarizing Intermolecular Forces

37. At which point in this flowchart would a distinction be made between SiH4 and SiH2Br2? • The presence of ions • b. The presence of polar molecules • H atoms bonded to N, O, and F • d. The flow chart would not differentiate between them.

38. Sample Exercise 11.2Predicting Types and Relative Strengths of Intermolecular Attractions List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling point. Solution Analyze We need to assess the intermolecular forces in these substances and use that information to determine the relative boiling points. Plan The boiling point depends in part on the attractive forces in each substance. We need to order these substances according to the relative strengths of the different kinds of intermolecular attractions. Solve The attractive forces are stronger for ionic substances than for molecular ones, so BaCl2should have the highest boiling point. The intermolecular forces of the remaining substances depend on molecular weight, polarity, and hydrogen bonding. The molecular weights are H2, 2; CO, 28; HF, 20; and Ne, 20. The boiling point of H2 should be the lowest because it is nonpolar and has the lowest molecular weight. The molecular weights of CO, HF, and Ne are similar. Because HF can hydrogenbond, it should have the highest boiling point of the three. Next is CO, which is slightly polar and has the highest molecular weight. Finally, Ne, which is nonpolar, should have the lowest boiling point of these three. The predicted order of boiling points is, therefore, H2< Ne < CO < HF < BaCl2 Check The boiling points reported in the literature are H2, 20 K; Ne, 27 K; CO, 83 K; HF, 293 K; and BaCl2, 1813 K—in agreement with our predictions.

39. Sample Exercise 11.2Predicting Types and Relative Strengths of Intermolecular Attractions Continued Practice Exercise 1 List the substances Ar, Cl2, CH4, and CH3COOH in order of increasing strength of intermolecular attractions. (a) CH4 < Ar < CH3COOH < Cl2 (b) Cl2< CH3COOH <Ar < CH4 (c) CH4< Ar < Cl2< CH3COOH (d) CH3COOH <Cl2 <Ar <CH4 (e) Ar < Cl2<CH4 <CH3COOH Practice Exercise 2 (a) Identify the intermolecular attractions present in the following substances and (b) select the substance with the highest boiling point: CH3CH3, CH3OH, and CH3CH2OH.

40. 11.3 Properties of Liquids

41. Intermolecular Forces • boiling point (previously discussed) and melting point • viscosity • surface tension • capillary action

42. Viscosity • Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

43. Surface Tension • Surface tension is the amount of energy required to increase the surface area of a liquid by a unit amount. • Therefore, surface molecules are packed more closely than bulk molecules. • This causes the liquid to behave as if it had a “skin”

44. Surface Tension • Surface tension results from the net inward force experienced by the molecules on the surface of a liquid - Therefore, surface molecules are packed more closely than bulk molecules. • This causes the liquid to behave as if it had a “skin” • Viscosity increase with the stronger intermolecular forces and decreases with the higher T.

45. How do viscosity and surface tension change? (a) as temperature increases (b) as intermolecular forces of attraction become stronger Viscosityincreasesas intermolecular forces increasewhile surface tension decreases. Bothviscosity and surface tension increasewithincreasingtemperature. Viscositydecreases as intermolecular forces increasewhile surface tension increases. Bothviscosityand surface tension increasewithdecreasingtemperature. Bothviscosity and surface tension increase as intermolecularforces increase and temperaturedecreases. Bothviscosity and surface tension decrease as intermolecularforces increase and temperatureincreases.

46. Types of Intermolecular Forces • Cohesive forces - Intermolecular forces that bind similar molecules to one another are called cohesive forces. • Adhesive forces - Intermolecular forces that bind a substance to a surface are called adhesive forces. These forces are important in capillary action.

48. Capillary Action • The rise of liquids up narrow tubes is called capillary action. • Adhesive forces attract the liquid to the wall of the tube. • Cohesive forces attract the liquid to itself. • Water has stronger adhesive forces with glass; mercury has stronger cohesive forces with itself.

49. If the inside surface of each tube were coated with wax, would the general shape of the water meniscus change? Would the general shape of the mercury meniscus change? • Shape of Water MeniscusShape of Hg Meniscus • Yes, inverted U Yes, downward U • b. No change Yes, downward U • No change No change • d. Inverted U No change

50. 11.4 Phase Changes • Conversion from one state of matter to another is called a phase change. • Energy is either added or released in a phase change. • Phase changes: melting/freezing, vaporizing/condensing, subliming/depositing.