Chapter 12 IM forces

Chapter 12 IM forces - • Why is water usually a liquid...

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Unformatted text preview: • Why is water usually a liquid and not a gas? • Why does liquid water boil at such a high temperature for such a small molecule? • Why does ice float on water? • Why do snowflakes have 6 sides? • Why is I2 a solid whereas Cl2 is a gas? • Why are NaCl crystals little cubes? C hapte 12:I M r Force The s. WHY? WHY? Chapter 12: Intermolecular Forces and Liquids • Gas: The Ideal Gas Law – why does it work? Answers: gas molecules act independently; they move about freely; the attractive forces between gas molecules are very weak near STP. For liquids & solids the molecules do not act independently as they take up less space; the attractive forces are stronger due to intermolecular forces. We have studied We INTRAMOLECULAR forces INTRAMOLECULAR —the forces holding atoms the together to form molecules. together Now turn to forces between Now molecules — molecules INTERMOLECULAR forces. INTERMOLECULAR Forces between molecules, between ions, or between molecules and ions. molecules I nte rm cular ole Force s Force I on-I on Force s for com parison of m agnitude Na+—Cl- in salt These are the These strongest forces. strongest Lead to solids with high melting temperatures. temperatures. NaCl, mp = 800 oC NaCl, MgO, mp = 2800 oC MgO, Attraction Be e I ons and twe n Pe ane Dipole rm nt s Pe •• •• water -δ dipole O H H +δ Water is highly polar Water and can interact with positive ions to give hydrate ions in d water. water. Attraction Be e I ons and Attraction twe n Pe ane Dipole rm nt s Pe Attraction between ions and dipole depends on ion charge and ion-dipole distance. ion ion-dipole Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+ δ- H O H δ+ ••• Mg2+ -1922 kJ/mol δ- H O H δ+ δ- H O H δ+ ••• Na + ••• Cs+ -405 kJ/mol -263 kJ/mol Dipole -DipoleForce s Dipole Such forces bind molecules having Such permanent dipoles to one another. permanent Dipole -DipoleForce s Dipole Influence of dipole-dipole (permanent or Influence induced) forces is seen in the boiling points of simple molecules of same mass. points Compd Mol. Wt. Boil Point 2 N (nonpolar) o 28 -196 C -196 o CO 28 -192 C -192 Hydroge Bonding n Hydroge A special form of dipole-dipole special attraction, which enhances dipoleattraction, dipole attractions. H-bonding is strongest when X and Y are N, O, or H-bonding F:highly electronegative F:highly H-Bonding Be e Two twe n Me thanol Mole s cule Me -δ +δ -δ H-bond 2 Hydroge Bonding in H O n Hydroge H-bonding is especially H-bonding strong in water because because • the O—H bond is very the polar(diff electronegativity) electronegativity) • there are 2 lone pairs there on the O atom on Accounts for many of Accounts water’s unique water’s properties. properties. 2 Hydroge Bonding in H O n Hydroge Ice has open lattice-like structure. Ice density is < liquid and so solid floats on Ice water. water. One of the VERY few One substances where solid is LESS DENSE than the liquid. than H bonds ---> abnormally high specific heat bonds capacity of water (4.184 J/g•K) capacity This is the reason water is used to put out fires, it This is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy. energy. Hydroge Bonding in H O n Hydroge Boiling Points of S ple im Hydroge ontaining n-C C pounds om Active Figure 13.8 Double helix Double of DNA of Portion of a DNA chain DNA FORC I NVOLVI NG ES I NDUCED DI POLES How can non-polar molecules such as O2 and I2 dissolve in water? The water dipole INDUCES a dipole INDUCES The in the O2 electric cloud. in Dipole-induced dipole dipole FORC I NVOLVI NG ES I NDUC DI POLES ED • Process of inducing a Process dipole is polarization polarization • Degree to which electron Degree cloud of an atom or molecule can be distorted Increases with mass of the gas. in its in polarizability. polarizability London Forces • Fritz London suggested it in 1928. He said that molecules set up instantaneous dipoles because of shifting electron density. These instantaneous dipoles affect neighbor molecules, distorting their electron clouds and inducing dipoles in them. They actually occur in all molecules – nonpolar as well as polar. • They are the only forces for nonpolar molecules • They are sometimes more important than dipole-dipole forces • Large, high MW molecules have high polarizability or are more polarizable, which means more distortion and, thus strong (or stronger) London forces. These have higher enthalpies of vaporization and higher melting and boiling points FORC I NVOLVI NG ES I NDUC DI POLES ED Formation of a dipole in two nonpolar I2 molecules. molecules. Induced dipoleinduced dipole induced (London) (London) FORC I NVOLVI NG ES I NDUC DI POLES ED The induced forces between I2 molecules are very weak, so solid I2 sublim s (goes from e very sublim a solid to gaseous molecules). solid FORC I NVOLVI NG ES I NDUCED DI POLES The magnitude of the induced dipole depends The on the tendency to be distorted. Higher molec. weight ---> larger induced dipoles. dipoles. Boiling Points of Non-polar Molecules Halogen Boiling Point (oC) Noble Gas Boiling Point (oC) Fluorine (F 2 ) Chlorine (Cl2) Bromine (Br2) Iodine (I2 ) – 188.1 Helium – 268.6 – 34.6 Neon – 245.9 + 58.8 Argon – 185.7 +184.4 Krypton – 152.3 Xenon – 107.1 Radon – 61.8 for molecules with similar molecular masses, the more compact the molecule, the lower its boiling point I nte ole rm cular Force S m s um ary Liquids, Changes of State (Phase Changes) and Dynamic Equilibrium • A phase change is the process that occurs when a substance is transformed from one physical state to another • Solid to Liquid is called fusion. The reverse is called freezing. • Liquid to Vapor is called vaporization (or evaporation). The reverse is called condensation. The energy required to vaporize a sample is the standard molar enthalpy of vaporization, ∆ Hvap • Solid to Vapor is called sublimation. The reverse is called deposition. Liquids Liquids In a liquid • molecules are in molecules constant motion constant • there are appreciable there intermolec. forces intermolec. • molecules close molecules together together • Liquids are almost Liquids incompressible incompressible • Liquids do not fill the Liquids container container Liquids Liquids The two key properties we need to The describe are EVAPORATION and its EVAPORATION opposite—CONDENSATION opposite— LIQUID evaporation---> Add energy break IM force VAPOR make IM force Remove energy <---condensation A liquid in a beaker • Phase changes tend toward a condition called “dynamic equilibrium.” • Open beaker – the liquid evaporates until it is all gone • Covered beaker – only some of the liquid evaporates. In the covered beaker both evaporation and condensation occur. Initially, the rate of evaporation is greater than the rate of condensation. After a time, the rate of condensation becomes equal to the rate of evaporation. This defines the equilibrium. At this point the pressure of the vapor is constant and it is called the vapor pressure of the liquid. Liquids—Evaporation Liquids—Evaporation To evaporate, To molecules must have sufficient energy to break IM forces. IM Breaking IM forces Breaking requires energy. The process of evaporation is endothermic. endothermic. Distribution of Ene in a Liquid rgy Distribution N u m b e r o f m o le c u le s lo w e r T h ig h e r T At higher T a much larger number of At molecules has high enough energy to break IM forces and move from liquid to vapor state. vapor 0 M o le c u la r e n e r g y m in im u m e n e r g y n e e d e d to b r e a k IM fo r c e s a n d e v a p o r a te High E molecules carry away E. You High cool down when sweating or after swimming swimming When molecules of liquid When are in the vapor state, they exert a VAPOR PRESSURE PRESSURE EQUI LIBRIUM VAPOR EQUI PRES URE is the S PRES pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation. the Liquids Liquids VAPOR PRESSURE VAPOR • Vapor Pressure depends upon (a) the temperature and (b) the chemical composition. (a) temperature – VP rises as the temperature of a liquid increases. Compare plots of P vs. T for a gas (what do you think this looks like) with plots of P (vapor pressure of a liquid) vs. T as shown in figure 12.18. • Chemical composition – compare different compounds. See the data below and figure 12.18. • Vapor pressure is independent of how much liquid is present or the surface area of the liquid as long as some liquid remains when equilibrium is reached Me asuring Equilibrium Vapor Pre Vapor ssure Liquid in flask evaporates and exerts pressure on manometer. Active Fig. 13.17 Compound (liquid) Vapor Pressure (mm Hg) at 20 C Water (H2O) 17 Methanol (CH3OH) 65 Carbon Tetrachloride (CCl4) 77 Acetone (CH3COCH3) 185 Diethylether (CH3CH2OCH2CH3) 442 EquilibriumVapor Pre ssure Equilibrium ActiveFigure12.18 VAPOR PRESSURE CHANGES WITH TEMPERATURE Boiling Point • Boiling Point – the temperature where the vapor pressure is equal to the external atmospheric pressure. If this external pressure is 1 atm, the temperature is called the normal boiling point. The boiling point (or normal boiling point) depends on the strengths of intermolecular attractions. Some examples can clearly be seen in Table 12.4. Note the relation between the enthalpy of vaporization and the normal boiling point of the liquid Liquids Liquids HEAT OF VAPORI ZATION is the heat req’d (at ’d constant P) to vaporize 1 mol of the liquid. constant LIQ + heat ---> VAP Compd. ∆Hvap (kJ/mol) IM Force IM H2O 40.7 (100 oC) H-bonds 40.7 SO2 26.8 (-47 oC) 26.8 dipole Xe 12.6 (-107 oC) induced dipole 12.6 induced Clausius-Clapeyron equation • A plot of lnVP (the natural log of the vapor pressure) versus 1/T (T in kelvin) yields a straight line. The slope of this line given the enthalpy of vaporization . The enthalpy of vaporization can be estimated from only two data points (two VP values at two temperatures) using an equation called the Clausius-Clapeyron equation. This is on page 612 in the textbook with an example. I will propose a different equation and give another example. EquilibriumVapor Pre ssure& theC lausiusEquilibrium C yron Equation lape • Clausius-Clapeyron Clausius-Clapeyron equation — used to find ∆H˚vap. ∆H˚ • The logarithm of the vapor The pressure P is proportional to ∆Hvaporiation and to 1/T. ∆H • ln P = –(∆H˚vap/RT) + C ∆Hvap 1 P2 1 ln = ­ P1 R T1 T2 Phase Diagrams (Section 12.10). • Phase Diagrams: pressure (P) versus Temperature (T) graphs showing the equilibria between the solid, liquid and gas phases of a substance. • a) along each curve or line there are 2 phases in equilibrium. b) at the triple point there are 3 phases in equilibrium. c) there are 3 regions – solid, liquid, and vapor. d) the normal melting point and boiling point are the temperatures of the s-l line and l-v curve where they cross the pressure axis at 760 mm (1 atm). e) water has an abnormal phase diagram. The s-l line slopes to the left. f) the triple point for carbon dioxide is above one atmosphere. g) the l-v curve ends at a point where the interface between liquid and vapor disappears. This point is called the critical point. Here the temperature is the critical temperature and the pressure is the critical pressure. PhaseDiagram (wate s r) Phase SOLID LIQUID VAPOR C ritical T and P As P and T increase, you finally reach As the CRITICAL T and P CRITICAL Pcritical Note that line goes straight up High Pressure . LIQUID Tcritical GAS High Temperature Above critical T Above no liquid exists no matter how high the pressure. pressure. C ritical T and P COMPD Tc(oC) Pc(atm) H2O CO2 374 31 218 73 ...
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