Chapter 13: Intermolecular Forces in Liquids & Solids

Chapter 5: Solids, Liquids & Phase Transitions

Overview

· organization chart for both condensed phases

· forces

· properties

· phase changes

5.1 Bulk Properties of Phases & Kinetic Molecular Theory

· recall kinetic molecular theory (for gases):

· gas atoms or molecules widely-separated

· no forces of attraction between them

· atoms or molecules in continual, random, rapid motion

· kinetic energy determined by temperature

· first two above are unique to gases, latter two are largely true for liquids and solids too

· comparison of bulk properties, in terms of:

· molar volume of gases, much larger than liquids & solids

· compressibility of gases (to form liquids) large; not dramatic for liquids (to form solids)

· thermal expansion of gases much larger than liquids & solids

· fluidity & rigidity

· diffusion

· surface tension (Fig. 5.4; not for gases)

5.2 Types of Intermolecular Forces

· in order of decreasing strength, Fig. 5.9; also important for solutions (chapter 6)

· ion - dipole interactions (ion - ion interactions treated separately)

· dipole - dipole interactions

· dipole (or ion-)- induced dipole

· induced dipole - induced dipole interactions

· similar repulsive forces for all

· term: van der Waals forces used for those interactions not involving ions; define van der Waals radius of atoms based on optimum between attractive and repulsive forces when ions not involved

Ion - Dipole Interactions (Fig. 5.6)

· dipole: charge separation, or (more likely) partial-charge separation, in a molecule i.e. - a polar molecule

· strength of ion - dipole interaction depends on:

· separation distance of ion and dipole

· charge on ion

· magnitude of dipole

· several examples based on the polar molecule, water:

· reaction is exothermic, heat given off = heat/enthalpy of hydration

(hydration here; more generally, solvation)

· distance effect seen with alkali cations

· compare with H+, DHhydr = -1090 kJ/mol (i.e.- hydronium ion, H3O+)

· summary: Li+, -515; Na+, -405; Cs+, -263 kJ/mol

· similarly, charge and distance effect seen

Dipole - Dipole Interactions (Fig. 5.5)

· energy generally released when molecules condensed, taken up when the condensed phase vaporized, i.e.- accompanying the equilibrium: gas D liquid

· for polar molecules, this is due to dipole - dipole interactions

· comparison of boiling points (a measure of the heat required for vaporization) allows categorization as polar or non-polar molecules (compare pairs of similar molar masses)

· solubility considerations are also due to a matching of polarities of solute and solvent (“like dissolves like”)

Hydrogen Bonding

· special class of dipole - dipole interactions due to small size and low electronegativity of H when bonded to small, electronegative atoms, especially N, O and F

· eg. Xd-- Hd+···Yd-

· the strongest are given in a Table (weaker ones with Cl, S, etc.)

· example, HF

· compare ethanol and dimethyl ether, both C2H6O:

· compare structures

· H-bonding in the alcohol, only dipole - dipole forces in the ether

· physical data:

dipole moment melting pt. boiling pt.

1.69D -114oC 78oC

1.30 -142 -25

· periodic trends in Fig. 5.10

Unusual Properties of Water

· H-bonding to extreme!!

· two H’s and two lone pairs (non-bonding pairs of electrons) on each O, Fig. 5.11

· form networks in the condensed phases

· perfect “diamond lattice” in ice, Fig. 5.12

· some disruption upon melting (still 85% of H-bonds), contraction of structure

· hence, ice has lower density than water; water max density at 4°C, Fig. 5.13

Dispersion Forces: Interactions with Induced Dipoles

· weakest of all intermolecular interactions

· two types:

· between polar and non-polar molecules, Fig. 5.7 (dipole (or ion-)- induced dipole)

· the larger the non-polar, the greater the interaction (eg. solubility of diatomic gases in water), due to polarizability

· between non-polar molecules, Fig. 5.8 (induced dipole - induced dipole)

· groups of examples in Table, trends in boiling points

Summary/Decision Tree for Intermolecular Interactions - Fig. Not in text

5.4 & 5.5 Phase Equilibria & Transitions

Physical Properties of Liquids

· distribution of kinetic energies of molecules in a liquid sample (similar to kinetic theory of gases)

· equilibrium with gas phase (part of Fig. 5.17):

· heat/enthalpy of vaporization, DHvap, is energy required (i.e.- endothermic) to escape intermolecular forces, Fig. 5.4

· eg. for water, DHvap = + 40.7 kJ/mol

· heat/enthalpy of condensation, DHcond, is energy liberated (i.e.- exothermic) on forming intermolecular interactions (same magnitude, opposite sign)

Vapor Pressure

· in a closed space above a liquid, Fig. 5.14, pressure in the gas phase stabilizes at a fixed value = equilibrium vapor pressure, dependent on temperature, Fig. 5.15 and Table 5.1; note:

· points on a line represent equilibrium pressure

· also a partial phase diagram - at a given T and P

· points to left of line represent the liquid phase region

· points to right of line represent the gas phase region

· boiling points at atmospheric pressure (see below)

· volatility, the tendency to escape into the gas phase ranked according to equilibrium vapor pressure

· practical application: “water pump” in lab better in winter than summer

Boiling Point

· Fig. 5.15, line at 1 atm

· equilibrium vapor pressure equals atmospheric pressure at the boiling point

· in an open vessel vaporized molecules can escape

· note dependence on pressure, applications:

· cooking in Salt Lake City

· vacuum distillations

Physical Properties of Solids

· disrupt the lattice to form a liquid (part of Fig. 5.17):

· fusion is endothermic, freezing/crystallization is exothermic

· also characterized by melting point, lowest temperature at which fusion occurs

· grouped according to forces; note ion-size dependence

· (note: sublimation also possible for solid D gas equilibrium; egs. H2O, I2, CO2, naphthalene; I2 on 1st page photo)

5.6 Phase Diagrams

· phase transitions, above

· influence of temperature and pressure on phases given in diagrams

· eg. H2O in Fig. 5.19: follow 2-phase lines (note negative slope to solid-liquid line); note triple point, freezing point, boiling point, critical point

· eg. CO2 in Fig. 5.21: note positive slope to solid-liquid line, critical point

Critical Temperature and Pressure

· keep increasing temperature in Fig. 5.19 and 5.21, two phases coalesce

· critical temperatures and pressures (both minima) shown

· new “phase”: supercritical gas/fluid has a density like that of the liquid but flow properties and ability of molecules to be separate from one another like a gas

· applications:

· liquefaction of gases, eg. air conditioning, fuels (must be below critical point)

· supercritical fluid extraction (must be above critical point), eg. CO2 (Tc = 31°C, Pc = 73 atm) used for decaffeinating coffee

Surface Tension, Capillary Action & Viscosity

· all are phenomena due to intermolecular interactions

· surface tension

· forces different in bulk liquid than at surface, Fig. 5.4; net, inward force at surface

· “skin” on surface, resists spreading as a film on another surface = surface tension

· surface layer interactions, in some cases, counteracted by interaction with another material

· eg. H2O with glass (H-O-H vs. Si-O-(H)), hence meniscus in a tube (Fig. 5.20), which is extreme in a very narrow tube (capillary)

· application: chromatography

· bulk liquid flow influenced by intermolecular interactions

· viscosity is the resistance to flow, which increases as the intermolecular interactions do

· eg. ethanol (two C’s) compared to longer chain alcohol, octanol (eight C’s) or to a “polyol” such as glycerol