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Lecture 11.1

Liquids and Intermolecular Forces The fundamental difference between states of matter is the distance between particles. In solids, particles are closely packed in an ordered array, their positions are :ixed. In gases, particles are far apart and possess complete freedom of motion. In liquids, particles are closely packed but randomly oriented, and retain freedom of motion. Table 11.1 lists properties of solids, liquids and gases. Condensed phase term given to solid and liquid states since particles are closer together. The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: (a) the kinetic energy of the particles, and (b) the strength of the attractions between the particles.

The :igure distinguishes intramolecular from intermolecular attraction. The attractions between molecules (intermolecular attractions) are not nearly as strong as the intramolecular attractions that hold compounds together. However, this intermolecular attractions are strong enough to control the physical properties of substances. These properties include boiling and melting points, vapor pressures and viscosities. Intermolecular forces are referred to as van der Waals forces; there are three types: (a) Dipoledipole interactions (b) Hydrogen bonding (c) London dispersion forces

London Dispersion Forces Simply called dispersion forces, are attractions between an instantaneous dipole and an induced dipole. Instantaneous dipole can develop temporarily when the electrons in an atom or molecule are distributed unsymmetrically about the nucleus because of the constant motion of the electrons.

The tendency of an electrons to be distorted or unsymmetrically distributed is called polarizability. Induced dipole forces result when an ion or a dipole induces a dipole in an atom or a molecule with no dipole. These are weak forces. Dispersion forces are present in all molecules, whether they are polar or nonpolar.

nucleus

electrons

+ - unsymmetrical distribution symmetrical distribution

Factors Affecting the Strength of London Dispersion Forces

The shape of the molecule: 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. Molecular weight or molar mass of the molecule: The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds that are easier to polarize. Boiling point is a property that measures strength of intermolecular forces; the graph below shows increasing boiling point with increasing atomic weight indicating increasing dispersion forces.

DipoleDipole Interactions Molecules that have permanent dipoles are attracted to each other. The positive end of one 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. The more polar the molecule, the higher its boiling point.

Which have a greater effect, dipole dipole interactions or dispersion forces? If two molecules are of comparable size and shape, dipoledipole interactions will likely be the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties. 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.

Hydrogen Bonding The dipoledipole interactions experienced when H is bonded to N, O, or F are unusually strong, these interactions are called hydrogen bonding. Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and :luorine. IonDipole Interactions Iondipole interactions are important in solutions of ions. The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

Use the :low-chart to determine the intermolecular forces in a molecule

Physical Properties Affected by Intermolecular Forces Intermolecular forces control how well molecules stick together. This affects many of the measurable physical properties of substances:

Melting and Boiling Points If molecules stick together more, they'll be harder to break apart Stronger intermolecular forces higher melting and boiling points Viscosity Viscosity is a measure of how well substances :low; resistance of a liquid to :low; It is related to the ease with which molecules can move past each other. Stronger intermolecular forces higher viscosity. Viscosity decreases with higher temperature.

Surface Tension Surface tension is a measure of the toughness of the surface of a liquid; results from the net inward force experienced by the molecules on the surface of a liquid Stronger intermolecular forces higher surface tension. Vapor Pressure This is a small amount of gas that is found above all liquids. Stronger intermolecular forces Lower vapor pressure. Sample Problem 1: In which of these substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH4), hydrazine (H2NNH2), methyl :luoride (CH3F), hydrogen sul:ide (H2S)? Why? Solution: The answer is hydrazine, H2NNH2S, since hydrogen bonding usually occurs only when the hydrogen is covalently bonded to N, O, or F. Sample Problem 2: List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling point. Solution: The boiling point depends in part on the attractive forces in each substance. We need to order these according to the relative strengths of the different kinds of intermolecular attractions. The attractive forces are stronger for ionic so BaCl2 should have the highest boiling point.

Sample Problem 2 continued: H2, CO and HF are covalent molecules, CO and HF are polar covalent, the polarity in H-F is stronger than CO, therefore HF has a higher boiling point than CO. In Ne and H2 are non-polar molecules, Ne has a higher mass than H2. The answer is: H2 < Ne < CO < HF < BaCl2 Energy Changes Associated with Phase Change

Heat of fusion is the amount of energy required to change a solid at its melting point to a liquid. Heat of vaporization is the amount of energy required to change a liquid at its boiling point to a gas. Heat of sublimation is the amount energy required to change a solid directly to a gas. The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during a phase change.

Sample Problem 3: Calculate the enthalpy change upon converting 1.00 mol of ice at -25 C to steam at 125 C under a constant pressure of 1 atm. The speci:ic heats of ice, liquid water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H2O, Hfus = 6.01 kJ/mol and Hvap = 40.67 kJ/mol. Solution: Ice at -25 C to ice at 0 C: Ice 0 C to water at 0 C: Water at 0 C to water at 100 C: Water at 100 C to water vapor at 100 C : Water vapor at 100 C to water vapor at 125 C: Total H = 0.914 + 6.01 + 7.52 + 40.7 + 0.828 = 56.0 kJ

! = 1.00!!"#!!! 18.0!!!"# !!2.03 !! ! ! 0 !25 !! = 914!!= 0.914!!" ! = 1.00!!"#!!! 6.01!!"!"# = 6.01!!"

! = 1.00!!"#!!! 18.0!!!"# !!4.18 !! ! ! 100 0 !! = 7520!! = 7.52!!" ! = 1.00!!"#!!! 40.67!!"!"# = 40.7!!" ! = 1.00!!"#!!! 18.0!!!"# !!1.84 !! ! ! 125 100 !! = 828!!= 0.828!!"

Vapor Pressure At any temperature some molecules in a liquid have enough energy to escape to the gaseous state. As the temperature rises, the fraction of molecules that have enough energy increases. As more molecules escape from the liquid to the gaseous state, the pressure (called the vapor pressure) exerted by these gases over the liquid increases. When the liquid molecules evaporate at the same rate as the vapor molecules condensing the liquid and vapor reach a state of dynamic equilibrium. The temperature at which the vapor pressure equals atmospheric pressure is called the boiling point of the liquid. The temperature at which the vapor pressure is at 760 torr is called the normal

boiling point of the liquid.

Phase Diagram Phase diagrams display the state of a substance at various pressures and temperatures, and the places where equilibria exist between phases. The liquidvapor interface starts at the triple point (T), at which all three states are in equilibrium, and ends at the critical point (C), above which the liquid and vapor are indistinguishable from each other. Each point along this line is the boiling point of the substance at that pressure. The interface between liquid and solid marks the melting point of a substance at each pressure. Below the triple point the substance cannot exist in the liquid state. Along the solidgas line those two phases are in equilibrium; the sublimation point at each pressure is along this line.

Phase Diagram of Water The slope of the solidliquid line is negative. This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid

Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

Sample Problem 4: Use the :igure in slide 14 to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm. Answer: Since 0.80 atm = 610 torr, the boiling point is at 27 oC Sample Problem 5: Use the phase diagram for methane, CH4, shown in the :igure: (a)What are the approximate temperature and pressure of the critical point? (b) What are the approximate temperature and pressure of the triple point? (c) Is methane a solid, liquid, or gas at 1 atm and 0 C? (d) If solid methane at 1 atm is heated while the pressure is held constant, will it melt or sublime? (e) If methane at 1 atm and 0 C is compressed until a phase change occurs, in which state is the methane when the compression is complete? Answers: (a) The critical point is at approximately -80 C and 50 atm. (b) The triple point is at approximately -180 C and 0.1 atm. (c) Gas (d) Melt (e) Gas

Liquid Crystals Some substances do not go directly from the solid state to the liquid state. In the intermediate state called liquid crystalline state, some traits of solids and some of liquids are exhibited by the molecule. Liquid crystals have properties between those of conventional liquid and those of solid crystal. For instance, a liquid crystal may :low like a liquid, but its molecules may be oriented in a crystal-like way. Liquid crystals can exhibit color changes with changes in temperature. Examples of liquid crystals can be found both in the natural world and in technological applications. Most contemporary electronic displays use liquid crystals.

Unlike liquids, molecules in liquid crystals have some degree of order. In nematic liquid crystals, molecules are only ordered in one dimension, along the long axis. In smectic liquid crystals, molecules are ordered in two dimensions, along the long axis and in layers.

In cholesteryl liquid crystals, nematic-like crystals are layered at angles to each other.

End of Chapter 11