Ch 10 Liquids and Solids Dr Namphol Sinkaset

Ch. 10: Liquids and Solids Dr. Namphol Sinkaset Chem 200: General Chemistry I

I. Chapter Outline I. III. IV. Introduction Intermolecular Forces Phase Transitions Phase Diagrams

I. Condensed States • Liquids and solids are the condensed states because of the close proximity of atoms/molecules to one another. • This proximity leads to much more frequent interactions than in gases. • Interactions depend on chemical identity of the substance and determine many physical properties.

II. States of Matter • The state of a sample if matter is the result of a battle between attractive forces between particles and kinetic energy.

II. Electrostatic Forces • Every molecule in a sample of matter experiences two types of electrostatic forces. § Intramolecular forces: the forces that exist within the molecule (bonding). These forces determine chemical reactivity. § Intermolecular forces: the forces that exist between molecules. These forces determine physical properties.

II. Intermolecular Forces • Intermolecular forces are attractive forces that originate from interactions between charges, partial charges, and temporary charges on molecules.

II. Types of IM Forces • There are different kinds of IM forces, each with a different level of strength. § Dispersion forces § Dipole-dipole attractions § *Hydrogen “bonding”

II. Dispersion Forces • Dispersion forces (London forces) are present in all molecules and atoms and results from changes in e- locations.

II. Instantaneous Dipoles • Charge separation in one creates charge separation in the neighbors.

II. Dispersion Force Strength • The ease with which e-’s can move in response to an external charge is known as polarizability. • Large atoms with large electron clouds tend to have stronger dispersion forces. • Larger molecules tend to have stronger dispersion forces.

II. Noble Gas Boiling Points

II. Dispersion Forces and Shape • Molecular size is not the only factor…

II. Dispersion Forces in a Family

II. Dispersion Forces in Action

II. Dipole-Dipole Attractions • Occur in polar molecules which have permanent dipoles, so attraction is always present.

II. Hydrogen “Bonding” • This IM force is a misnomer since it’s not an actual bond. • Occurs between molecules in which H is bonded to a highly electronegative element (N, O, F), leading to high partial positive and partial negative charges. • It’s a “super” dipole-dipole attraction.

II. H “Bonding” Water

II. Boiling Point Trend

II. Effect of H “Bonding”

II. Sample Problem • Which substance has the highest boiling point and why? a) CH 3 OH b) CO c) N 2

III. Vaporization and IM Forces • From experience, we know that water evaporates in an open container. • What factors influence rate of vaporization?

III. Vaporization Variables • Temperature • Surface area • IM forces

III. Energetics of Vaporization • As molecules evaporate, what happens to the temperature of the samples left in the beaker? • Vaporization is an endothermic process – it’s the reason why we sweat when we get too hot. • Condensation is an exothermic process.

III. Dynamic Equilibrium • In an open flask, a liquid will eventually evaporate away. • What about a closed flask?

III. Vapor Pressure

III. Vapor Pressure and Temp. • Vapor pressure depends on temperature and IM forces. • Why?

III. Boiling Point • When T is increased, the vapor pressure increases due to the higher # of molecules that can break away and enter gas phase. • What if all molecules have necessary thermal energy? • At this point, vapor pressure = external pressure, and boiling point is reached. • The temperature at which vapor pressure equals 1 atm is the normal boiling point.

III. Boiling Point vs. Altitude

III. Pvap – T Relationship • The Clausius-Clapeyron equation describes the relationship between vapor pressure and temperature.

III. Linear Form • This equation is in linear form, y = mx + b. • The heat of vaporization can be found using graphical analysis. • Use R = 8. 314 J/mole·K.

III. Graphical Analysis

III. Clausius-Clapeyron Equation, 2 -point Form • If you have two sets of pressure, temperature data for a liquid, the more convenient 2 -point form of the Clausius-Clapeyron equation can be used.

III. Sample Problem • Propane has a normal boiling point of -4. 20 °C and a heat of vaporization of 19. 04 k. J/mole. What is the vapor pressure of propane at 25. 0 °C?

III. Other Phase Changes • Sublimation is the direct conversion of particles from the solid phase to the gas phase. § Average KE is low, but always some that have enough KE to break away. • Fusion is the conversion of solid to liquid. • Also have deposition and freezing.

III. Energetics of Fusion • Different compounds have different heats of fusion. • Notice they are much lower than heats of vaporization – why?

III. Energies of Phase Changes • The enthalpies involved in a phase change depends on the amount of substance and the substance itself. • We look at a heating curve for H 2 O at 1. 00 atm pressure. • Note that there are sloping regions and flat regions in the curve. (Why? )

III. Heating Curve for H 2 O

III. Heating Curve, Sloped Regions • In these regions, heat is being used to increasing KE – hence changes in T. • The heat required depends on the specific heat capacity of the phase.

III. Heating Curve, Flat Regions • Here, the temperature stays the same, so the average KE stays the same. • Thus, the PE must be increasing. • The heat gained is a factor of the ΔH of the phase change.

IV. Phase Diagrams • Measurements of phase transitions over a variety of different temperatures and pressures are used to construct phase diagrams. • Phase diagrams allow predictions of the phase in which a substance will exist under specific conditions.

IV. Generic Phase Diagram

IV. Phase Diagram for H 2 O

IV. The Critical Point • In a sealed container, as T of liquid is heated, more and more vapor is formed, and P increases. • At the critical temperature, a supercritical fluid forms; liquid can’t exist above this temperature.