chapter 11 liquids solids n.
Skip this Video
Loading SlideShow in 5 Seconds..
Chapter 11: Liquids & Solids PowerPoint Presentation
Download Presentation
Chapter 11: Liquids & Solids

Loading in 2 Seconds...

play fullscreen
1 / 59

Chapter 11: Liquids & Solids - PowerPoint PPT Presentation

  • Uploaded on

Chapter 11: Liquids & Solids. The molecular compounds like water, ammonia, and carbon dioxide have different physical properties because of the intermolecular forces. Comparison of all three phases:. Liquids & Solids. Liquids & Solids. Compressibility. Changes of State.

I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
Download Presentation

Chapter 11: Liquids & Solids

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
    Presentation Transcript
    1. Chapter 11: Liquids & Solids • The molecular compounds like water, ammonia, and carbon dioxide have different physical properties because of the intermolecular forces. • Comparison of all three phases:

    2. Liquids & Solids

    3. Liquids & Solids

    4. Compressibility

    5. Changes of State • Changes in state can be induced by a change in temperature or pressure.

    6. Intermolecular Forces • Forces between molecules. • Always LESS in energy than actual bond. • The attractive force between two HCl molecules is about 16 kJ/mol. • The bond dissociation energy of the HCl bond is about 431 kJ/mol.

    7. Intermolecular Forces • One method to compare the strength of intermolecular forces is to examine the substance’s boiling point. • When the forces are relatively weak, then the boiling point is small. • Ex) HCl, bp = -85oC. • There are three main types of intermolecular forces between neutral substances.

    8. Intermolecular Forces • In the LS packet, we identified molecules as being polar or non-polar based on shape, types of atoms, etc. • Polar molecules have a dipole – that is a positive and negative end.

    9. Intermolecular Forces • Thus, the first type of force is called Dipole-Dipole (or DD) forces and occurs for any polar molecule. • The larger the dipole moment, the more DD forces.

    10. Intermolecular Forces • If you cool any non-polar molecule or atom to a low enough temperature, then it will liquefy. • Yet, these have no reason to be attractive to each other. • Fritz London first proposed a theory in 1930. • On average, electrons in an atom like He are evenly distributed. • But, in one INSTANT, the two electrons may both be on the same side.

    11. Intermolecular Forces • Thus, in that one INSTANT, a He atom would have an instantaneous dipole. • This is called the London Dispersion (or LD) force.

    12. Intermolecular Forces • Since all molecules have electrons, they all have a LD force. • The polarizability of an atom or molecules electrons depends on two factors. • The number of electrons. More electrons = More LD forces. • The shape of molecule. More spread out = more LD forces.

    13. Intermolecular Forces

    14. Non-polar Alkanes

    15. Intermolecular Forces • The formula C5H12 has three structural isomers. CH3 – CH2 – CH2 – CH2 – CH3 CH3 CH3 CH3 – CH2 – CH – CH3 CH3 – C – CH3 CH3

    16. Non-polar Branched Alkanes

    17. Comparison of the group 4A, 5A, 6A, and 7A hydrides shows an interesting result. What type of forces do the group 4A have? Group 6A? What is the notable exception?

    18. Intermolecular Forces • The third type of force is a special case of DD force and is called Hydrogen Bonding (or HB). • The name “Hydrogen Bonding” is a misnomer! • HB can only occur when: • H is bonded to either N, O, or F. • The N, O, or F atom has at least one lone pair.

    19. Intermolecular Forces • The strength of Hydrogen Bonding varies from 5 kJ to 40 kJ, which is still much weaker than a covalent bond (200 – 1000 kJ). • However, it is MUCH stronger than DD or LD forces. • Thus, it can greatly increase the boiling point temperatures of molecules.

    20. Intermolecular Forces • HB forces are very important in biochemistry. • Proteins are made from the twenty amino acids. • The structure of the amino acid has both an –OH group and an –NH2 group that can HB. R O  H – N – C – C – O – H  H H

    21. Intermolecular Forces • Predicting relative boiling points. • Determine the molecular weight. • Determine the type(s) of intermolecular forces present. • If weights are similar, then LD < DD < HB • If weights are very dissimilar, then #2 probably does not matter. • However, HB can really distort the bp’s! • Ex) H2O, bp = 100oC, MW = 18 g/mol versus CCl4, bp = 76oC, MW = 154 g/mol.

    22. Intermolecular Forces • The strengths of attractions between the molecules may affect a liquids properties. • Viscosity • Surface Tension

    23. Intermolecular Forces • Viscosity is the resistance of a liquid to flow. • Liquids with low viscosity, like water, will produce a “splash” whereas liquids with high viscosity, like corn syrup or ketchup, will not.

    24. Intermolecular Forces • Viscosity tends to increase with more intermolecular forces and molecular weight. • Many liquids, like water, have a consistent viscosity over a wide range of temperatures. • Some liquids, like corn syrup, will decrease in viscosity as the temperature increases. • Multi-weight motor oil actually increases with an increase in temperature (ie. 5W – 30). • Non-Newtonian liquids (ie. “Slime”) have a variable viscosity at the same temperature.

    25. Surface Tension • Surface Tension is the “skin-like” appearance of the surface. • Results from surface molecules seeking six nearest neighbors like interior molecules.

    26. Surface Tension

    27. Phase Changes

    28. Phase Changes • Energy when changing between solid and liquid phase is called the Heat of Fusion and denoted as DHfus. • DHfusfor water is 6.01 kJ/mol or 334 J/g. • Energy when changing between liquid and gas is called the Heat of Vaporization and denoted as DHvap. • DHvapfor water is 40.67 kJ.mol or 2,260 J/g.

    29. Heating Curve

    30. Refrigeration • The basics of refrigeration. • First law of thermodynamics at work again! • Coolant is CF2Cl2 (old) or CF3CH2F (new).

    31. Vapor Pressure • Above the surface of any liquid, some liquid molecules will have enough energy to escape and become gas molecules. • In a closed system, an equilibrium will be achieved between the gas molecules and the liquid. • This is the vapor pressure.

    32. Vapor Pressure • As the temperature of the liquid increases, its vapor pressure will increase.

    33. Vapor Pressure • When the vapor pressure equals 1 atmosphere, then the liquid spontaneously becomes a gas. You would call this the boiling point. • Does pure water always boil at 100oC?

    34. Clausius-Clapeyron Equation • The graphs of vapor pressure versus temperature are approximately an exponential function. • Mathematically, if you take the natural logarithm (ln key on calculator) of the vapor press versus 1/T, then you get a linear relationship.

    35. Clausius-Clapeyron Equation • R is molar gas constant = 8.314 J/K mol and the T is the temperature in Kelvin • Heat of vaporization must be in J/mol. • Pressures can be in either atm or mmHg (must agree).

    36. Phase Diagrams • Display a single’s substances states of matter over a wide range of P and T.

    37. Carbon Dioxide • The phase diagram of CO2 shows that the liquid phase can only be found above a pressure of 5.11 atm. • As the temperature of solid CO2 increases, it undergoes sublimation.

    38. Water • The phase diagram of water has one very important difference. What is it?

    39. Solids • Solids can be either amorphous (random) or crystalline (repeating pattern). • Unit cell is the smallest repeating pattern for the crystalline structure. • Analogy: a hotel with many floors. • Structure of unit cell can have various lengths and angles.

    40. Solids • While many types of unit cells are possible, a few are seen many times in structures of metals, molecular, and ionic compounds. • Cubic unit cells – two main versions. • Body-centered cubic (BCC) – has atoms at each corner and an atom in the body-center. • Face-centered cubic (FCC) – has atoms at each corner and an atom on each face. • Important – just like a hotel room shares walls, floors, and ceilings with other rooms, so does a unit cell share atoms with other unit cells.

    41. Solids

    42. Solids

    43. Solids • Can also have atoms on edges in larger unit cells – namely for ionic compounds. • Thus, the following are the contributions for locations on or in a unit cell:

    44. Solids • Unit cell calculations will follow the formula: • Where Vcis the volume of the cubic unit cell, MW is the molar mass, C is the number of atoms per unit cell, D is the density (m/V), and Na is Avogadro’s Number.

    45. Solids • Another view – Closest Packing Model. • Assumes that atoms are hard spheres. • Maximize the density, minimize the empty spaces.

    46. Solids • First layer – what is the most efficient method of arrangement?

    47. Solids • Second layer is placed so that spheres sit in gaps from previous row. • Third layer can either repeat first layer yielding an ABABAB… pattern. • OR, the third layer is offset from the first two producing an ABCABCABC… pattern.

    48. Solids • The ABABAB… pattern produces a unit cell called hexagonal closest packing or HCP. • This is NOT a cubic unit cell!

    49. Solids • The ABCABCABC… pattern produces a unit cell called cubic closest packing or CCP. • However, CCP is the same as FCC!

    50. Solids