Intermolecular Forces, Liquids, and Solids

Intermolecular Forces, Liquids, and Solids Some Characteristic Properties of the States of Matter Assumes both the volume and shape of the container Is Compressible Diffusion within a gas occurs rapidly Flows readily Gas: Liquid Assumes the shape of the portion of the container it occupies Does not expand to fill the container Is virtually incompressible Diffusion within a gas occurs slowly Flows readily Retains its own shape and volume Is virtually incompressible Diffusion within a solid occurs extremely slowly Does not flow Solid

The strength of intermolecular forces vary over a wide range, but are generally weaker than covalent or ionic bonds 16 kJ will overcome the intermolecular attraction leading to vaporization 431 kJ of energy is required to break the covalent bond between HCl

http://phet.colorado.edu/en/simulation/states-of-matter-basics

the motions of electrons on its near neighbors London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment The instantaneous distribution at any given moment can be different from the average distribution producing an instantaneous dipole moment Because electrons repel, the motions of electrons in one atom influence the motions of electrons on its near neighbors

London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment The ease with which the charge distribution in a molecule can be distorted by an external force is called its polarizability. The greater the polarizability of a molecule, the more easily its electron cloud can be distorted to give a momentary dipole. In general, larger molecules tend to have greater polarizability because their electrons are farther from the nucleus Because molecular size and molecular mass tend to parallel one another, dispersion forces tend to increase in strength with increasing molecular weight. Halogen Boiling pt (K) Noble gas Boiling point (K) F2 85.1 He 4.6 Cl2 238.6 Ne 27.3 Br2 332.0 Ar 87.5 I2 457.6 Kr 120.9

London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment The shapes of molecules can play a role in the magnitude of dispersion forces.

London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment London dispersion forces may also operate within polar molecules and contribute to the overall attractive forces between the molecules; some times more than the dipole-dipole forces HCl: bp = 189.5 K Dipole Moment = 1.03 D HBr: bp = 206.2 K Dipole Moment = 0.79 D

Dipole Dipole forces exist between neutral polar molecules, are effective only when polar molecules are very close together, and are generally weaker then ion-dipole interactions Substance Dipole moment Bp For two particles of equal mass and size, the strengths of the intermolecular attractions increase with increasing polarity Propane 0.1 231 Dimethyl ether 1.3 249 Methyl chloride 2.0 249 Acetaldehyde 2.7 293 Molecules that are attracting one another spend more time near each other than those molecules which repel. Thus the overall effect is a net attraction

Hydrogen bonding is a special type of intermoleculer attraction that exists between the hydrogen atom in a polar bond (e.g., H F, H O, or H N) and an unshared electron pair on a nearby electronegative atom (usually an F, O, or N atom on another molecule) Because F, N, and O are so electronegative, a bond between hydrogen and any of these elements is quite polar, with hydrogen at the positive end.

Ion-Dipole Forces exist between an ion and the partial charge on the end of a polar molecule Cl- The magnitude of the interaction depends upon the charge of the ion, the dipole moment of the polar molecule and the distance from the center of the ion to the midpoint of the dipole

Metallic solids consist of metal atoms Metallic solids consist of metal atoms. Bonding is due to valence electrons that are delocalized throughout the entire solid. We can visualize these electrons as an array of positive ions immersed in a sea of delocalized electrons.

Intermolecular Forces, Liquids, and Solids Metallic bond examples: Cu, Fe, or alloy

ionic solids

network solids

Intermolecular Forces, Liquids, and Solids Properties of Liquids: Viscosity and Surface Tension The resistance of liquids to flow is called their viscosity. The greater the viscosity, the more slowly the liquid flows. Viscosity decreases with increasing temperature. Viscosity is related to the ease with which individual molecules of a liquid can move with respect to one another. It thus depends on the attractive forces between molecules and and their structural character.

Intermolecular Forces, Liquids, and Solids Properties of Liquids: Viscosity and Surface Tension Molecules in the interior are attracted equally in all directions, whereas those at the surface experience a net inward force This net inward force pulls molecules from the surface into interior, thereby reducing the surface area Surface tension is therefore defined as the energy required to increase the surface of a liquid by a unit amount Forces that bind like molecules to one another are called cohesive forces. Forces that bind molecules to a surface are called adhesive forces

http://apod.nasa.gov/apod/ap130424.html

Changes of State Vapor Pressure: pressure exerted by evaporating liquid in the space above that liquid once dynamic equilibrium has been established Substances with a high vapor pressure evaporate more rapidly than thosewith low vapor pressures. Theses substances are more volatile The vapor pressure of a substance increases as the temperature increases

Changes of State Whenever a change of state involves going to a less ordered state, energy must be supplied in order to overcome intermolecular forces The heat needed for vaporization is called the heat of vaporization (e.g. 40.67 kJ/mol for water) ENDOTHERMIC EXOTHERMIC Melting is also called fusion. The enthalpy of change associated with melting is called heat of fusion. (e.g. 6.01 kJ/mol for water)

Changes of State Enthalpy and Temperature Changes Accompanying Heating

Ht = .940 kJ + 6.01 kJ + 7.52 kJ + 40.67 kJ + .830 kJ = 55.97 kJ Calculate the enthalpy change associated with converting 1.00 mol of ice at -25°C to water vapor at 125 °C and 1 atm. The heat capacities per gram (specific heats) of ice, water, and steam are 2.09 J/g-°C , 4.18 J/g-°C , and 1.84 J/g-°C, respectively. The heat of fusion of ice is 6.01 kJ/mol, and the heat of vaporization of water is 40.67 kJ/mol Heat required to bring the ice from -25 °C to 0 °C : 18.0 g. 2.09 J 1.00 mol H2O 25 °C = 940 J 1.00 mol H2O g -°C Heat of fusion: 6.01 kJ 1.00 mol H2O = 6.01 kJ mol H2O Heat required to bring liquid water from 0 °C to 100 °C : 18.0 g. 4.18 J 1.00 mol H2O 100 °C =7520 J 1.00 mol H2O g -°C Heat of vaporization 40.67 kJ 1.00 mol H2O = 40.67 kJ mol H2O Heat required to raise the temperature of the vapor to 125 °C 18.0 g. 1.84 J 1.00 mol H2O 25°C =830 J 1.00 mol H2O g -°C Ht = .940 kJ + 6.01 kJ + 7.52 kJ + 40.67 kJ + .830 kJ = 55.97 kJ

Vapor Pressure and Boiling Point: A liquid boils when its vapor Changes of State Vapor Pressure and Boiling Point: A liquid boils when its vapor pressure equals the external pressure acting on the surface of the liquid The higher the outside pressure, the higher the boiling point

A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the difference states of matter. Line A-D represents the change in melting point of the solid under increased pressure Line A-B is the vapor curve of the liquid The line A-C represents the variation in vapor pressure during sublimation.

A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the difference states of matter. Note that the solid phase is the stable phase under condition of high pressure and low temperature Note that the gas phase is the stable phase under condition of low pressure and high temperature

Why are the fusion-freezing lines different for H2O and CO2? Why does CO2 sublime under normal conditions? How is freeze-dry food prepared?

AS THE TEMPERATURE IS INCREASED : At point 6-A, the H2O exists entirely as a solid. When the temperature reaches point 4-B, the solid begins to melt, and an equilibrium condition occurs between the solid and the liquid. At a yet high temperature, point 7-C, the solid has been converted entirely to a liquid. When point 8-D is encountered, vapor forms, and a liquid-vapor equilibrium is achieved. Upon further heating, to point 9-E, the H2O is converted entirely to the vapor phase. 6 4 7 9 8 7 5 3 2 1

A crystalline solid is a solid whose atoms, ions, or molecules are ordered in well-defined arrangements. Amorphous solids have no structure. A unit cell is the repeating unit that makes up the crystalline solid. The array of repeating points is called the crystal lattice

Covalent Network Solids are Held together in large networks or chains by covalent bonds. These solids are much harder and have higher melting points than molecular solids Molecular solids consist of atoms or molecules held together by dipole-dipole or London dispersion forces or hydrogen bonds

of the bonds depend upon the charges of the ions. Melting point Ionic solids consist of ions held together by ionic bonds. The strength of the bonds depend upon the charges of the ions. Melting point increases as bond strength increases NaCl The structure of the solid is dependent upon the charges and relative sizes of the ions. The coordination number increases as the ratio of the cation radius to the anion radius increases.

Atomic solids – (london) eg. He, Ar etc. atoms exist at lattice pts molecular solids (London, dipole, H-Bond) CH4, NH3 molecules exist at lattice pts Ionic solids (electrostatic attractions) NaCl cations & anions exist at lattice pts Metallic solids- (Cations held by a sea of electrons) Cu, Fe cations exist at lattice pts Network solids – (3 dimensional covalent bonds) diamond (C), quartz (SiO2)