Intermolecular Forces

Intermolecular Forces
Forces between (rather than within) molecules. dipole-dipole attraction: molecules with dipoles orient themselves so that “+” and “” ends of the dipoles are close to each other. hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10. 2: (a) The electrostatic interaction of two polar molecules
Figure 10.2: (a) The electrostatic interaction of two polar molecules. (b) The interaction of many dipoles in a condensed state.

Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Figure 10.3: (a) The polar water molecule. (b) Hydrogen bonding among water molecules. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.4: The boiling points of the covalent hydrides of the elements in Groups 4A, 5A, 6A, and 7A. General trend to increase down the group Small size Large electronegativity Copyright©2000 by Houghton Mifflin Company. All rights reserved.

London Dispersion Forces
relatively weak forces that exist among noble gas atoms and nonpolar molecules. (Ar, C8H18) caused by instantaneous dipole, in which electron distribution becomes asymmetrical. the ease with which electron “cloud” of an atom can be distorted is called polarizability. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.5: (a) An instantaneous polarization can occur on atom A, creating an instantaneous dipole. This dipole creates an induced dipole on neighboring atom B. (b) Nonpolar molecules such as H2 also can develop instantaneous and induced dipoles.

Some Properties of a Liquid
Surface Tension: The resistance to an increase in its surface area (polar molecules). Capillary Action: Spontaneous rising of a liquid in a narrow tube. Adhesion- stick to walls Cohesion- stick to each other Viscosity: Resistance to flow (molecules with large intermolecular forces). Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.6: A molecule in the interior of a liquid is attracted by the molecules surrounding it, whereas a molecule at the surface of a liquid is attracted only by molecules below it and on each side. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.7: Nonpolar liquid mercury forms a convex meniscus in a glass tube, whereas polar water forms a concave meniscus. Which has stronger adhesive forces? water Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.8: Two crystalline solids: pyrite (left), amethyst (right).
Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Types of Solids Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)]. Amorphous solids: considerable disorder in their structures (glass). Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Representation of Components in a Crystalline Solid
Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Crystal lattice A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance. TiO2 NaCl CsCl

Coordination number of an ion is
the number of ions of opposite charge That surround the ion in a crystal Sodium Chloride: Each Na+ is surrounded by 6 Cl- , therefore the Coordination # = 6 Each Cl- is surrounded by 6 Na+ Cubic centered

Representation of Components in a Crystalline Solid
Unit Cell: The smallest repeating unit of the lattice. simple cubic body-centered cubic face-centered cubic Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.9: Three cubic unit cells and the corresponding lattices.

Types of Crystalline Solids
Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl). Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice). Atomic Solid: have atoms at the lattice points. metallic, network, and group 8A solids Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Metallic Crystals Metals are made up of closely packed cations rather than neutral atoms. Chemical bonding results from the attraction between metal ions and the surrounding sea of electrons. Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move about the metal.

Because of this arrangement
Metals are good conductors of heat and electricity Metals are malleable Metals are ductile Metals have high tensile strength Metals have luster

Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors.

Every atom has 8 neighbors Every atom has 12 neighbors Every atom has 12 neighbors

Metal Alloys Substitutional Alloy: some metal atoms replaced by others of similar size.

Metal Alloys Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.

Network Solids Composed of strong directional covalent bonds that are best viewed as a “giant molecule”. brittle do not conduct heat or electricity carbon, silicon-based graphite, diamond, ceramics, glass Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Diamond (sp3) tetrahedral arrangement hard, colorless, an insulator Graphite (sp2) layers of six member rings slippery, black, conductor (delocalization of the π bonds) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10. 22: The structures of diamond and graphite
Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

Figure 10.24: The p orbitals (a) perpendicular to the plane of the carbon ring system in graphite can combine to form (b) an extensive π-bonding network. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.25: Graphite consists of layers of carbon atoms.

Silicon Silicon-oxygen compound SiO4 silica (tetrahedral) Silicates- found in rocks, soil, and clay. Silica is heated and cooled rapidly- glass Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Computer-generated model of silica.

ceramics From clay (contain silicates, like glass) and firing at high temp. non metallic, strong, brittle, and resistant to heat and chemicals Glass- homogeneous Ceramics- heterogeneous ( crystal of silicates suspended in a glassy cement) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Semiconductors A substance in which some electrons can cross the band gap. Conductivity increases w/increase in temp. Conductivity is enhanced by doping with group 3a or group 5a elements. N-type (As; has 1 more e-) P-type (B; has 1 less e-) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.29: (a) A silicon crystal doped with arsenic, which has one more valence electron than silicon. (b) A silicon crystal doped with boron, which has one less electron than silicon.

Figure 10.31: The p-n junction involves the contact of a p-type and an n-type semiconductor.

Figure 10.33: The steps for forming a transistor in a crystal of initially pure silicon.

Molecular Solids Ice- lattice is occupied by H2O molecules Dry ice- solid CO2 Sulfur- contain S8 Phosphorus- P4 (white phosphorus) CO2, I2, S8, P4 (London dispersion) (gas) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

A “steaming” piece of dry ice.

Figure 10. 34: (a) Sulfur crystals (yellow) contain S8 molecules
Figure 10.34: (a) Sulfur crystals (yellow) contain S8 molecules. (b) White phosphorus (containing P4 molecules) is so reactive with the oxygen in air that it must be stored under water.

Ionic Solids Stable, high melting point substances, held together by strong electrostatic forces between oppositely charged ions. Closest packing model, where the smaller cation fits in the holes. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Vapor Pressure and Changes of State
Molecules of a liquid escape the liquid’s surface and form a gas- Vaporization or Evaporation. Endothermic (overcome intermolecular forces) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Heat of vaporization, or enthalpy of vaporization (ΔHvap )
Energy required to vaporize 1mol of liquid at 1atm. H2O has a large heat of vaporization(40.7kJ) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.38: Behavior of a liquid in a closed container.
Initial transfer of liquid to the vapor phase. Initially the # of vapor molecules increase and so does the rate of return. (condensation). Eventually equilibrium occurs Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Vapor Pressure . . . is the pressure of the vapor present at equilibrium. . . . is determined principally by the size of the intermolecular forces in the liquid. . . . increases significantly with temperature. Volatile liquids have high vapor pressures. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.40: (a) The vapor pressure of a liquid can be measured easily using a simple barometer of the type shown here. (b) The three liquids, water, ethanol (C2H5OH), and diethyl ether [(C2H5)2O], have quite different vapor pressures. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

In general substances with large molar masses have low vapor pressures (Large dispersion forces) ( the more e- the more polarizable) H2O smaller molar mass but has H-bonding (stronger!!!) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

(a) The vapor pressure of water, ethanol, and diethyl ether as a function of temperature. Nonlinear (b) Plots of In(Pvap) versus 1/T (Kelvin temperature) for water, ethanol, and diethyl ether. linear Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Straight line y=mx + b ln(Pvap) = -Hvap/R (1/T) + C
Hvap = enthalpy of vap R = universal gas const J/Kmol C = const for each liquid

Why do we need to know this?
Can find Hvap by measuring Pvap at several temps and evaluating slope or..

If we know Hvap and Pvap at one temp, we can find them at another
Can combine the eqn b/c C does not depend on temp in order to solve for Pvap at another temp

Equation ln(PvapT1) - ln(PvapT2) = Hvap/R (1/T2 - 1/T1) OR
ln(PvapT1) + Hvap/RT1) = C = ln(PvapT2) + Hvap/RT2) ln(PvapT1) - ln(PvapT2) = Hvap/R (1/T2 - 1/T1) OR Ln(PT1/PT2) = Hvap/R (1/T2 - 1/T1)

Solve The vp of water at 25oC is 23.8 torr and the Hvap at 25oC is 43.9 kJ/mol. What is the vp at 50. oC? 93.7 torr

Melting Point Heat of fusion, enthalpy of fusion ΔHfus
Vibration of water molecules increases as T increases. Molecules break loose from lattice points and solid changes to liquid. (Temperature is constant as melting occurs.) vapor pressure of solid = vapor pressure of liquid Plateau at 0°C. Heat of fusion, enthalpy of fusion ΔHfus

If the vapor pressure of the solid is greater than that of the liquid? What state does the substance exist in? Liquid Will this temperature be above or below the Melting Point temperature? Above Copyright©2000 by Houghton Mifflin Company. All rights reserved.

If the vapor pressure of the solid is less than that of the liquid? What state does the substance exist in? Solid Will this temperature be above or below the Melting Point temperature? Below Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.46: Solid and liquid phases in equilibrium with the vapor pressure.

Boiling Point At 100°C the liquid water reaches BP.
Temperature remains constant: added energy is used to vaporize the liquid. vapor pressure of liquid = pressure of surrounding atmosphere ***All intermolecular changes (physical) no intramolecular bonds broken.

Normal melting point- T at which the solid and liquid states have the same vapor pressure under conditions were the total pressure is 1atm. Normal boiling point- T at which the vapor pressure of the liquid is exactly 1atm Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Super cooling Cooled below O°C at 1atm and remain in the liquid state. As it is cooled it may not reach the degree of organization necessary to form ice at O°C. Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Boiling chip releasing air bubbles acts as a nucleating agent for the bubbles that form when water boils. Superheating: Rapid heating of a liquid may allow it to exist as a liquid at temperatures above the normal boiling point . VP of liquid > atmosphere The gas bubbles burst before rising to the surface, blowing liquid out of the container (bumping) Copyright©2000 by Houghton Mifflin Company. All rights reserved.

Figure 10.49: The phase diagram for water.

Figure 10.47 Water in a Closed System with a Pressure of 1 atm Exerted on the Piston
To form a bubble here, the piston would have to move upwards! This won’t happen unless Pbubble ≥ Pexternal Pbubble = VP of liq! So VP ≥ Pexternal for boiling Copyright © Houghton Mifflin Company. All rights reserved.

Filled w/ice (no air space) T=-20°C Piston= 1atm pressure
Experiment 1 Pressure at 1atm Filled w/ice (no air space) T=-20°C Piston= 1atm pressure Since T<0, VP of ice is <1atm, no vapor present As it is heated, only ice until: T=0 *VPsolid=VPliquid melting Still no vapor ; PL <1atm T increases, solidliquid Now only water (still no vapor) T=100C ; VPliquid = 1atm; boiling occurs; change to vapor T increases, Liquidvapor

Experiment 2 P=2 torr (.0026 atm) T= -20 Just ice
T increases to -10 , sublimation occurs VPsolid=VPatm No liquid because the VP of liquid is always greater than 2 torr

Experiment 3 Pressure 4.58 torr (0.006 atm) T= -20, only ice
When T=0.01°C; Triple Point solid and liquid water have identical VP of 4.58 torr All three phases exist. Closed system!! Can not happen outside.

Experiment 4 P=225atm T=300°C ; liquid ??
As T increases the liquid changes to vapor, goes through intermediate “fluid” region. (neither liquid nor vapor: beyond the critical point) critical temperature: temperature above which the vapor can not be liquefied no matter how much pressure. critical pressure: pressure required to liquefy AT the critical temperature. critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

How about “Expt 5”? Expt 5 Why?? Density of ice is less than
What variable is changing, which variable is held constant, and what phase change occurs (and at what point)? ANSWERS: P increases T is held constant Melting occurs (at crossing point) Expt 5 Why?? Density of ice is less than that of liquid. Volume Increases. Copyright © Houghton Mifflin Company. All rights reserved.

Figure 10. 51: The phase diagram for water
Figure 10.51: The phase diagram for water. At point X on the phase diagram, water is a solid. What happens as you increase the pressure? MP decreases Why?? Density of ice is less than that of liquid. Volume Increases.

Intermolecular Forces

Similar presentations

Presentation on theme: "Intermolecular Forces"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Intermolecular Forces

Similar presentations

Presentation on theme: "Intermolecular Forces"— Presentation transcript:

Similar presentations

About project

Feedback