Presentation on theme: "Ch. 11: Liquids, Solids, and Intermolecular Forces Dr. Namphol Sinkaset Chem 200: General Chemistry I."— Presentation transcript:
Ch. 11: Liquids, Solids, and Intermolecular Forces Dr. Namphol Sinkaset Chem 200: General Chemistry I
I. Chapter Outline I.Introduction II.Intermolecular Forces III.Vaporization and Vapor Pressure IV.Energies of Phase Changes V.Phase Diagrams
I. 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? 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 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 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 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 Force Dispersion force (London force) is 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. Dispersion Force and Size Molecular size is not the only factor…
II. Dispersion Force and Shape Shape influences how the molecules interact with one another…
II. Dipole-Dipole Force Occurs in polar molecules which have permanent dipoles, so attraction is always present.
II. 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” Polar liquids are miscible with other polar liquids, but not with nonpolar liquids. Can be explained with intermolecular forces.
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.
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. 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 Equilibrium In an open flask, a liquid will eventually evaporate away. What about a closed flask?
III. 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 Equilibrium Systems at dynamic equilibrium will seek to return to dynamic equilibrium when disturbed.
III. Vapor Pressure and Temp. Vapor pressure depends on temperature and IM forces. Why?
III. Clausius-Clapeyron Equation The nonlinear relationship between vapor pressure and temperature can be written in a linear form:
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. 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 Point At the boiling point, those aren’t air bubbles!
IV. 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 H 2 O at 1.00 atm pressure. Note that there are sloping regions and flat regions in the curve. (Why?)
IV. 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 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 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 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 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 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.