Chapter 11: Liquids & Solids

1. Chapter 11: Liquids & Solids Renee Y. Becker Valencia Community College

2. Kinetic Molecular Theory Introduction • Gases • Gas particles act independent of one another • Attractive forces are very weak • Particles are free to move randomly • Occupy whatever space available Liquids and solids are different from gases in that they have strong attractive forces between particles

3. Polar Covalent Bonds • Polar Covalent Bonds • Form between a non-metal/non-metal of different electronegativities EN = ENA – ENB When EN  1.7 ionic bond When EN < 1.7 polar covalent bond When EN < .5 non-polar covalent bond

4. Polar Molecules • Polar Molecules • Just as bonds can be polar, molecules as a whole can be polar • Net sum of individual bond polarities and lone-pair contributions

5. Dipole Moment • Dipole moment, , (ionic and covalent) • Measure of net molecular polarity • The magnitude of the charge Q at either end of the molecular dipole times the distance, r, between the charges  = Q x r • Expressed in debyes, D, where 1 D = 3.336 x 10-30 coulomb meters • Q = 1.6 x 10-19 C (electron charge)

6. Dipole Moment % Ionic Character = experimental  (100%) calculated  a high % IC means that the bond is similar to or is an ionic bond a low % IC means that it is more like a covalent bond

7. Example 1 Chloromethane a) Calculate the dipole moment b) Calculate the % ionic character of the bond Experimentally measured dipole moment = 1.87 D C-Cl bond distance = 178 pm = 178 x 10-12 m If we assume that the contributions of the nonpolar C-H bonds are small, then most of the chloromethane dipole moment is due to the C-Cl bond

8. Example 2 Hydrochloric acid Calculate the % ionic charcater 1.      Distance between atoms is 127 pm 2.      Experimentally measured dipole moment = 1.03 D

9. Example 3 Tell which of the following compounds are likely to have a dipole moment and show the direction of each. a) SF6 b) CHCl3 c) CH2Cl2 d) CH2CH2

10. Intermolecular Forces • Van der Waals forces – intermolecular forces as a whole, all are electrical in origin and result from the mutual attraction of unlike charge or mutual repulsion of like charges. 4 main types • Dipole-dipole • Ion-dipole • Dispersion forces • Hydrogen bonding

11. Dipole-dipole • Neutral but polar molecules experience dipole-dipole forces as a result of electrical interactions among dipoles on neighboring molecules. • Forces can be attractive or repulsive, depending on the orientation of the molecules. • c) These forces are weak 3-4 kJ/mol and only significant if molecules are close

12. Ion-dipole Result of electrical interactions between an ion and the partial charges on a polar molecule b) Particularly important in aqueous solutions of ionic substances such as NaCl, in which polar water molecules surround the ions

13. London Dispersion Forces a) Result from the motion of electrons b) At any given time more electrons may be in a particular area of the molecule c)  This gives the molecule an instantaneous dipole d)  This short lived dipole can affect the electron distribution in neighboring molecules and induce temporary dipoles in them e) More electrons a molecule has the stronger the dispersion forces

14. Hydrogen Bonding a) Attractive interaction between a hydrogen atom bonded to a very electronegative atom (O, N, F) and an unshared electron pair on another electronegative atom • Hydrogen bonds arise because O-H, N-H, and F-H bonds are highly polar with partial positive charge on the hydrogen and partial negative on the electronegative atom. • Hydrogen has no core electrons to shield its nucleus and it is small so it can be approached closely by other molecules d) The dipole-dipole attraction between the hydrogen and an unshared electron pair on a nearby atom is usually strong

15. Hydrogen Bonding e) Water is able to form a vast 3D network of hydrogen bonds because each H2O molecule has two hydrogens and two electron pairs

16. Intermolecular Forces

17. Example 4 Identify the likely kinds of intermolecular forces in the following A)   HCl B) CH3CH3 C) CH3NH2 D) Kr

18. Example 5 Of the substances Ar, Cl2, CCl4 and HNO3 which has: a)  The largest dipole-dipole forces? b)   The largest hydrogen-bond forces? c) The smallest dispersion forces?

19. Properties of Liquids Viscosity 1. The measure of a liquids resistance to flow 2. Related to the ease with which individual molecules move around in the liquid and thus to the intermolecular forces present 3. Substances with small non-polar molecules have weak intermolecular forces and low viscosities (free flowing) 4. More polar substances have stronger intermolecular forces and have higher viscosities

20. Properties of Liquids Surface Tension 1. The resistance of a liquid to spread out and increase its surface area • Caused by differences in intermolecular forces experienced by molecules at the surface and the interior • Surface molecules feel attractive forces on only one side and are drawn in toward the liquid 4. Interior molecules are drawn equally in all directions 5. Higher in liquids that have stronger intermolecular forces

21. Phase Changes • Physical form changes but chemical identity does not change Fusion (melting) solid  liquid Freezing liquid  solid Vaporization liquid  gas Condensation gas  liquid Sublimation solid  gas Deposition gas  solid

22. Thermochemistry Free energy change, G 1. All naturally occurring processes, every phase change has a free-energy change 2.  G =  H - T S 3.   H enthalpy, heat flow, positive (from surrounding to system, bond breaking takes energy), negative (from system to surroundings, bond making) 4.   S entropy, disorder, positive (ordered to disorderd), negative (disordered to ordered)

23. Thermochemistry

24. Thermochemistry Calculating the temperature at a phase change • G > 0 non-spontaneous G < 0 spontaneous G = 0 equilibrium 2. Set G = 0 0 = H - TS and solve for T T = H/S

25. Heating Curve for H2O

26. Heating Curve for H2O Heating curve for H2O E = molar heat capacity (T) Energy to heat ice from -25C to 0C Molar heat capacity of ice= 36.57 J/molC E = Energy to heat H2O from 0C to 100C Molar heat capacity of liquid H2O = 75.4 J/ molC E =

27. Thermochemistry Heat of fusion, Hfusion The amount of energy required to overcome enough intermolecular forces to convert a solid into a liquid Heat of Vaporization, Hvap The amount of energy necessary to convert a liquid into a gas

28. Example 6 Chloroform has Hvap = 29.2 kJ/mol and Svap = 87.5 J/K mol. What is the boiling point of chloroform?

29. Evaporation, Vapor Pressure, and Boiling Point Vapor Pressure 1. The partial pressure of a gas in equilibrium and at constant temperature with liquid 2. Thepressure exerted by gaseous molecules above a liquid

30. Evaporation, Vapor Pressure, and Boiling Point

31. Evaporation, Vapor Pressure, and Boiling Point The higher the temperature and the lower the boiling point of the substance the greater the fraction of molecules in the sample that have sufficient kinetic energy to break free from the surface of the liquid and escape into the vapor.

32. Evaporation, Vapor Pressure, and Boiling Point Numerical value of Vapor Pressure depends on: a)      Magnitude of intermolecular forces The smaller the forces the higher the vapor pressure, loosely held molecules escape easily b)      Temperature The higher the temperature, the higher the vapor pressure, larger fraction of molecules have sufficient kinetic energy to escape

33. Evaporation, Vapor Pressure, and Boiling Point The Clausius-Clapeyron Equation ln Pvap = - Hvap 1 + C R T Y = m x + b m is the slope and b is the y-intercept Where R is the gas constant 8.314 J/K mol C is a constant characteristic of each specific substance T temperature in Kelvin

34. Evaporation, Vapor Pressure, and Boiling Point The Clausius-Clapeyron Equation ln P2 = Hvap 1 - 1 P1 R T1 T2

35. Evaporation, Vapor Pressure, and Boiling Point • This equation makes it possible to calculate the heat of vaporization of a liquid by measuring its vapor pressure at several temperatures and then plotting the results to obtain the slope • Once the heat of vaporization and the vapor pressure at one temperature are known, the vapor pressure of the liquid at any other temperature can be calculated.

36. Evaporation, Vapor Pressure, and Boiling Point Normal boiling point 1. The temperature at which boiling occurs when the pressure is exactly 1 atm. 2. Boiling point – when the vapor pressure of a liquid rises to the point where it becomes equal to the external pressure

37. Example 7 The vapor pressure of ethanol at 34.7C is 100.0 mm Hg, and the heat of vaporization of ethanol is 38.6 kJ/mol. What is the vapor pressure of ethanol in mm Hg at 65.0C?

38. Example 8 The normal boiling point of benzene is 80.1 C and the heat of vaporization is 30.8 kJ/mol. What is the boiling point of benzene (in C) on top of Mt. Everest where P = 260 mm Hg?

39. Phase Diagrams • Shows which phase is stable at different combinations of pressure and temperature.

40. Phase Diagrams Triple Point: The only condition under which all three phases can be in equilibrium with one another. Critical Temperature (Tc): The temperature above which the gas phase cannot be made to liquefy at any pressure. Critical Pressure (Pc) : The minimum pressure required to liquefy a gas at its critical temp. Supercritical Fluid: Neither true liquid nor true gas Normal boiling and melting point always at 1 atm

41. Example 9 Can you label the following? a) solid region b) Liquid region c)  Gas region d)  Normal boiling point e)  Normal melting point f)   Triple point g)  Supercritical fluid region h) Critical point, what is the critical pressure and temperature

42. Types of solids 1. Molecular solid -held together by intermolecular forces -H2O(s), CO2(s) 2. Metallic solid -positively charged atomic cores surrounded by delocalized electrons -Fe, Cu, Ag

43. Types of solids 3. Ionic solid -cations and anions held together by electrical attraction of opposite charges -NaCl 4. Covalent network solid -atoms held together in large networks by covalent bonds -diamonds