Ch. 11: Liquids, Solids, and Intermolecular Forces Dr. Namphol Sinkaset Chem 200: General Chemistry I.

Ch. 11: Liquids, Solids, and Ch. 11: Liquids, Solids, and Intermolecular ForcesIntermolecular Forces

Dr. Namphol Sinkaset

Chem 200: General Chemistry I

I. Chapter OutlineI. Chapter Outline

I. Introduction

II. Intermolecular Forces

III. Vaporization and Vapor Pressure

IV. Energies of Phase Changes

V. Phase Diagrams

I. Electrostatic ForcesI. 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.

I. Solid, Liquid, or Gas?I. Solid, Liquid, or Gas?

• Whether a substance exists as a solid, liquid, or gas depends on the relationship between the intermolecular attractions and the kinetic energy of the molecules. It’s a battle – which dominates? The KE or

the IM attractions? Recall that the average KE of a sample is

related to its temperature.

I. KE vs. IM ForcesI. KE vs. IM Forces

• Gas: the kinetic energy of the molecules is much greater than the intermolecular attractions.

• Liquid: the kinetic energy of the molecules is moderately greater than the intermolecular attractions.

• Solid: the kinetic energy of the molecules is less than the intermolecular attractions.

II. Intermolecular ForcesII. Intermolecular Forces

• IM forces originate from interactions between charges, partial charges, and temporary charges on molecules.

• IM forces are relatively weak because of smaller charges and the distance between molecules.

II. Types of IM ForcesII. Types of IM Forces

• There are different kinds of IM forces, each with a different level of strength. Dispersion force Dipole-dipole force *Hydrogen “bonding” Ion-dipole force

II. Dispersion ForceII. Dispersion Force

• Dispersion force (London force) is present in all molecules and atoms and results from changes in e- locations.

II. Instantaneous DipolesII. Instantaneous Dipoles

• Charge separation in one creates charge separation in the neighbors.

II. Dispersion Force StrengthII. 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 PointsII. Noble Gas Boiling Points

II. Dispersion Force and SizeII. Dispersion Force and Size

• Molecular size is not the only factor…

II. Dispersion Force and ShapeII. Dispersion Force and Shape

• Shape influences how the molecules interact with one another…

II. Dipole-Dipole ForceII. Dipole-Dipole Force

• Occurs in polar molecules which have permanent dipoles, so attraction is always present.

II. Effect of Dipole-Dipole ForceII. Effect of Dipole-Dipole Force

• Polar molecules have dispersion forces and dipole-dipole forces.

• Effects can be seen in boiling and melting points.

II. “Like Dissolves Like”II. “Like Dissolves Like”• Polar liquids are miscible with other polar

liquids, but not with nonpolar liquids.• Can be explained with intermolecular forces.

II. Hydrogen “Bonding”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 force.

II. H “Bonding” in Ethanol & WaterII. H “Bonding” in Ethanol & Water

II. Effect of H “Bonding”II. Effect of H “Bonding”

• Hydrogen “bonding” is a very strong intermolecular force.

• Without hydrogen “bonding” life as we know it could not exist!

II. Ion-Dipole ForceII. Ion-Dipole Force

• This type has already been discussed.

• Example: NaCl(s) dissolved in water.

II. Summary of IM ForcesII. Summary of IM Forces

III. Vaporization and IM ForcesIII. Vaporization and IM Forces

• From experience, we know that water evaporates in an open container.

• What factors influence rate of vaporization?

III. Vaporization VariablesIII. Vaporization Variables

• Temperature• Surface area• IM forces

III. Heat of VaporizationIII. Heat of Vaporization• The energy needed to vaporize 1 mole of a

liquid to gas is the heat of vaporization.• Can be thought of the energy needed to

overcome IM forces of the liquid.

III. Dynamic EquilibriumIII. Dynamic Equilibrium

• In an open flask, a liquid will eventually evaporate away.

• What about a closed flask?

III. Dynamic EquilibriumIII. Dynamic Equilibrium• As evaporation occurs,

headspace fills with gas molecules.

• Gas molecules condense back to liquid phase.

• Eventually, rates become equal.

• Pressure of gas at dynamic equilibrium is called the vapor pressure.

III. Dynamic EquilibriumIII. Dynamic Equilibrium

• Systems at dynamic equilibrium will seek to return to dynamic equilibrium when disturbed.

III. Vapor Pressure and Temp.III. Vapor Pressure and Temp.

• Vapor pressure depends on temperature and IM forces.

• Why?

III. Clausius-Clapeyron III. Clausius-Clapeyron EquationEquation

• The nonlinear relationship between vapor pressure and temperature can be written in a linear form:

III. Clausius-Clapeyron III. Clausius-Clapeyron Equation, 2-point FormEquation, 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. Boiling PointIII. Boiling Point

• When temperature is increased, the vapor pressure increases due to the higher number of molecules that can break away and enter the gas phase.

• What if all molecules have the necessary thermal energy?

• At this point, vapor pressure equals the external pressure, and the boiling point is reached.

III. Boiling PointIII. Boiling Point

• At the boiling point, those aren’t air bubbles!

IV. Energies of Phase ChangesIV. 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 1.00 moles of H2O at 1.00 atm pressure.

• Note that there are sloping regions and flat regions in the curve. (Why?)

IV. Heating Curve for HIV. Heating Curve for H22OO

IV. Heating Curve, Segment 1IV. Heating Curve, Segment 1

• At this stage, we are heating ice from -25 °C to 0 °C, increasing KE (vibrational motions).

• The heat required depends on the specific heat capacity of ice.

IV. Heating Curve, Segment 2IV. Heating Curve, Segment 2

• 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 heat of

fusion, the heat needed to melt 1 mole of solid.

IV. Heating Curve, Segment 3IV. Heating Curve, Segment 3

• During this stage, water is being heated from 0 °C to 100 °C; again, KE is increasing.

• The heat gained depends on the specific heat capacity of water.

IV. Heating Curve, Segment 4IV. Heating Curve, Segment 4

• Again, the temperature stays the same, so the average KE stays the same.

• PE must be increasing.• The heat gained is a factor of the heat of

vaporization.

IV. Heating Curve, Segment 5IV. Heating Curve, Segment 5

• During this stage, steam is heated from 100 °C to 125 °C; average KE is increasing.

• The heat gained depends on the specific heat capacity of steam.

V. Phase DiagramsV. Phase Diagrams

• The relationship between pressure, temperature, and the three phases can be summarized in a phase diagram.

• A phase diagram allows the prediction of how a substance will respond to changes in pressure and/or temperature.

V. Phase Diagram of WaterV. Phase Diagram of Water

V. IV. I22 and CO and CO22 Phase Diagrams Phase Diagrams