Chapters 11 and 12: Intermolecular Forces of Liquids and Solids

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1 1 Chapters 11 and 12: Intermolecular Forces of Liquids and Solids 11.1 A Molecular Comparison of Liquids and Solids The state of matter (Gas, liquid or solid) at a particular temperature and pressure depends on two factors: - The kinetic energy of the particles - The strength of the attraction between the particles See the figure below Gases-the average energy of the attraction between molecules is much smaller than their average kinetic energy. Thus, gas molecules are far apart and do not interact much with one another. They flow and diffuse readily. Liquid the average inter-atomic forces are strong enough relative to the average kinetic energy to hold particles together but not strong enough to keep the molecules in place. Thus, in liquids, molecules are free to move around. Solids inter-atomic forces strong enough relative to the average kinetic energy to hold particles in place. That is why solids are rigid. Thus, in solids molecules are locked into position 11.2 Intermolecular Forces These are forces that exist between molecules. Intermolecular forces are much weaker than ionic or covalent bonds (Intra-molecular attractions). See figure bellow. 1

2 2 When a substance boils and melts, intermolecular forces are broken. The boiling and melting points reflect the strength of the intermolecular forces. The higher the intermolecular force the higher the boiling point of a liquid or melting point of a solid. When a substance condenses, intermolecular forces are formed. van der Waals forces: are intermolecular forces that exists between neutral molecules. Types of van der Waals forces - Dipole-dipole - London dispersion - ydrogen bonding In addition to the intermolecular forces described above, two other types of interactions also exit between fully charged ions. - Ion ion forces: force of attraction that exists between fully charged particles (cation and anion). And they are very stronger electrostatic attractions, stronger than intermolecular forces. E.g., NaCl - Na + Cl - Na + Cl - - Cl - Na + Cl - Na+ - Na + Cl - Na + Cl - - Ion dipole forces: exist between an ion and a polar molecule. Positive ions are attracted to the negative end of a dipole and vice versa. This occurs commonly in polar solutions with dissolved salts (ionic compounds). The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents. E.g. NaCl dissolves readily in 2O. - O O Na + O O O O - Cl Dipole dipole forces: are forces that exist between polar molecules. A polar molecule has one end that is partially negatively charged and another that is partially positively charged. It is not as strong as those involve ions. Many properties of liquids and gases reflect the strength of the intermolecular forces between the molecules. For example the intermolecular forces between Cl molecules are so weak that Cl boils at -85 C. That is Cl(l) Cl(g). Cl Cl Generally, for molecules of approximately equal mass and size intermolecular attraction increases with increasing polarity (see table bellow or Table 11.2 text) 2

3 3 What force holds non-polar molecules together? London Dispersion Forces: these are the weakest of all forces and exist between nonpolar molecules. Generally, dispersion forces tend to increase in strength with increasing molecular weight (see table 11.3 text). London dispersion forces are the only forces that exist between non-polar molecules, but also exists all other molecules. ydrogen bonding unusually strong type of dipole-dipole attraction that occurs when ydrogen is bonded to a strongly electronegative atom: N, O, F (see examples below). E.g. F F ydrogen bonding is very important in biological systems. E.g., DNA strands held together by -bonding. Protein folding: - Bond. 3

4 4 ydrogen bonding influences boiling points and many other properties of molecules. The figure above shows the boiling point of simple hydrogen compounds. Generally, boiling point increases with increasing molecular weight, which is due to increase in the strength of dispersion forces. owever, water is the exception in this trend; whose boiling point is much higher than expected based on molecular weight. This indicates that the intermolecular forces between water molecules are extremely strong. This strong intermolecular force is the -bonding. See figure 11.7 text. Densities of liquid 2O and Ice The hydrogen bonding interaction that exists between liquid water molecules is random. owever, when water freezes the interaction assumes an ordered arrangement. Because the ordered structure has more hexagonal holes, ice is more open and less dense than liquid water (see figure 11.10). Summarizing Intermolecular Forces 4

5 Some Properties of Liquids Some of the properties of liquids include: - Viscosity - Surfaces tension - Capillary action - Boiling point and freezing point Viscosity a measure of liquid s resistance to flow. Example: oney, oil viscous flow slowly Water, gasoline-less viscous flow easily Viscosity depends on - The attractive force between molecules- intermolecular forces - Structural features that allow molecules to become entangled. Larger molecules have higher viscosities, because attractive forces between molecules are higher. Water with high attraction force between molecules has higher viscosity than other smaller non-polar molecules. Surface tension energy required to spread a liquid on a surface. Liquids with large surface tension tend to remain in sphere rather than spread out. Cause of surface tension: intermolecular forces in the center of the liquid that are experienced in all directions, the forces at the surface, however are only experienced in one direction inward towards the liquid center. Cohesive forces: intermolecular forces that bind similar molecules to one another. Adhesive forces Intermolecular forces that bind a substance to a surface. Meniscus of water: U-shaped because the force between water and glass are stronger than between water molecules. Meniscus of mercury: Curved downward because forces between mercury atoms are greater than force between mercury and glass. Capillary Action: is the rise of liquids up narrow tubes. 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 Phase Changes The transformations of the three state of matter from one state to another, is called phase change. This is summarized in the heating curve shown below using water as an example. 5

6 6 eating curve eating curves are graphs used to calculate energy changes that accompany phase changes. E.g., calculate the enthalpy change upon converting 1.00 mol of ice at -25 C to water vapor (steam) at 125 C under constant pressure of 1 atmosphere. Specific heat of ice = 2.09 J/g K Specific heat of water = 4.18 J/g K Specific heat of steam = 1.84 J/g K For water: fus = 6.01 kj/mol; vap = kj/mol Solution: Q (heat) = p = mc T AB: = (1.00 mol)(18.0g/mol)(2.09 J/g K)(25 K) = 940 J = 0.94 J BC: = (1.00 mol)(6.01 kj/mol) = 6.01 kj CD: = (1.00 mol)(18.0 g/mol)(4.18 J/g K)(100 K) = 7520 J = 7.52 kj DE: = (1.00 mol)(40.67 kj/mol) = 40.7 kj EF: = (1.00 mol)(18.0 g/mol)(1.84 J/g K)(25) = 830 J = 0.83 kj = 0.94 kj kj kj kj kj = 56.0 kj What happens when ice is heated? When heat is added the temperature rises and the ice starts to melt. At some point the temperature increase stops. Temperature remains constant during the phase change. The added energy is used to overcome the attractive forces holding between molecules. At melting point the following equilibrium exists 2 O(s) 2 O(l) fus = molar enthalpy or molar heat of fusion (melting) Melting or fusion: solid to liquid 6

7 7 Ones the ice is melted, the temperature of the liquid increases again, until another plateau is reached. The added heat is used to overcome the forces holding the molecules together, and vaporization occurs. At the boiling point the following liquid equilibrium exists. 2 O(l) 2 O(g) vap = molar enthalpy or molar heat of vaporization Vaporization: liquid to gas Additional heat added will increase the temperature of the gas. Sublimation: solid to gas Thus, some solids go directly from solid to gas phase: CO2 (s) (dry ices) CO2 (g) sub = molar enthalpy or heat of sublimation Critical Temperature and Pressure Gases may be liquefied by increasing the pressure at a particular temperature. Critical temperature: is the highest temperature at which a substance can exist as a liquid. Critical Pressure: is the pressure required for liquification at the critical temperature. The greater the intermolecular forces the, the higher the critical temperature, the easier to liquefy a substance Vapor pressure Consider a liquid in an evacuated closed container as shown in the figure below (see also figure text). A B When the ethanol is added, the liquid will start to evaporate (as in A) and will form some molecules of the liquid in the gas phase. Then the molecules in the gas phase will collide with each other (as in B) and with the walls of the vessel, thus the pressure above the 7