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Liquids and Solids

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  1. Liquids and Solids AP Chem Unit 10

  2. Sections • Intermolecular Forces • Liquid state • Solid Structures • Metal Structures • Carbon and Silicon Networks

  3. Sections • Molecular Solids • Ionic Solids • Vapor pressure and State Change • Phase Diagrams

  4. States of Matter When considering the three states of matter, properties of gases are strikingly different than solids and liquids. Liquids and solids share many similar characteristics • compressibility • density • intermolecular forces

  5. States of Matter • H2O(s)H2O(l) ΔH°fus = 6.02 kj/mol • H2O(l)H2O(g) ΔH°vap = 40.7 kj/mol • Water densities: • 25°C and 1atm .99707g/cm3 • 25°C and 1065 atm 1.046g/cm3 • 400°C and 1atm 3.26x10-4 g/cm3 • 400°C and 242 atm .157g/cm3

  6. Intermolecular Forces 10.1

  7. Intermolecular Forces • Electrons shared within the molecule are called intramolecular bonding. • In the condensed states of matter the attraction between molecules are called intermolecular forces.

  8. Intermolecular Forces It is important to realize that when a molecule changes state, the molecule stays intact. The changes in state are due to the change in forces surrounding the molecule not from changes within the molecule. • 40.7kj needed to vaporize water • 934kj to break the O-H bond

  9. Dipole–Dipole Forces Dipole-dipole forces occur when polar molecule (molecules with dipole moments) electrostatically attract each other by lining up the positive and negative ends of the dipoles. • Dipole-dipole forces are about 1% as strong as a covalent or ionic bond and rapidly become weaker when distances between the dipoles increases. The distances in a gas make these attractions relatively unimportant

  10. Dipole-Dipole Forces • In a condensed state, molecules line up dipoles to minimize repulsions and maximize attractions.

  11. Dipole-Dipole Forces Some dipole-dipole forces are unusually strong. These usually form between H and another very electronegative atom. • These are stronger due to the high polarity of the bond and the closeness of the dipoles between the atoms. • These strong attractions have a strong impact on melting points and boiling points.

  12. Boiling Points of Covalent Hydrides

  13. Hydrogen bonds Hydrogen bonds are the strongest in the smallest and lightest of the covalent molecules. This is primarily due to two factors: • large difference in electronegativities • small size of the atoms allows for close dipole interactions.

  14. Hydrogen bonds

  15. Hydrogen Bonds and Organics • Methanol (CH3OH) and ethanol (CH3CH2OH) have much higher boiling points than would be expected from their molar masses because of the O-H bonds that produce hydrogen bonding.

  16. London Dispersion Forces Even without dipoles, molecules exert forces on each other. • The forces that exist among noble gas atoms and nonpolar molecules are called London dispersion forces.

  17. London Dispersion Forces Usually it is assumed that electron dispersion is uniform throughout the molecule, but this is not always the case. • Since the movements of the electrons around the nucleus are somewhat random, a momentary nonsymmetrical electron distribution can develop that creates a temporary dipolar arrangement of charge.

  18. London Dispersion Forces • This temporary change in polarity can, in turn, temporarily change the distribution of the neighboring molecule. • This phenomenon leads to an inter-atomic attraction that is relatively weak and short-lived, but can be significant in larger atoms at lower temperatures. • larger atoms have more electrons and increases the probability of a temporary dipole.

  19. London Dispersion Forces

  20. London Dispersion Forces Polarizability is the ease at which an electron cloud can be distorted into a temporary dipole. • large atoms have a larger polarizability than smaller atoms • This also applies to molecules like H2, CH4, CCl4 and CO2;smaller molecules, but nonpolar.

  21. The Liquid State 10.2

  22. Liquid Characteristics • lack of rigidity • low compressibility • high density • rounded droplets • capillary action • viscosity

  23. Rounded Droplets • Occur due to the intermolecular forces of the liquid. The liquid molecules are subject to attraction from the side and from below, so liquid tends to form a shape with the minimum surface area – sphere. • The resistance of a liquid to increase surface area is from the energy that it takes to overcome intermolecular forces. This resistance is called surface tension.

  24. Rounded Droplets • Molecules that are polar and have stronger intermolecular forces have stronger surface tensions.

  25. Surface Tension

  26. Capillary Action Capillary action is the spontaneous rising of a liquid in a narrow tube. This action is due to two forces • cohesive forces- the intermolecular forces among the molecules. • adhesive forces – the attractive forces between the liquid and the container.

  27. Adhesive forces Adhesive forces happen when bonds within the container have polar bonds • For example: glass has O atoms that carry a partial negative charge that attracts the partial positive charge of the hydrogen in water. This balance between the strong cohesive forces and the strong adhesive forces produce a meniscus.

  28. Adhesive forces • A nonpolar substance, such as mercury, has a convex meniscus because the cohesive forces are stronger than the adhesive forces.

  29. Meniscus: Water vs. Mercury

  30. Viscosity Viscosity is a fluids resistance to flow. • liquids with strong cohesive forces tend to be highly viscous. • Example: glycerol is highly viscous because of its ability to create hydrogen bonds.

  31. Viscosity • Molecular complexity also can affect viscosity because they can become entangled in each other. • Example: Gasoline has carbon chains from 3-8C long and is nonviscous. Grease is 20-25C long and is very viscous.

  32. Introduction to Structures and Types of Solids 10.3

  33. Types of Solids • Crystalline solids • Amorphous solids

  34. Crystalline Solids Crystalline solids have a regular arrangement of components at a microscopic level and produce beautiful, characteristic shapes of crystals:

  35. Crystalline Solids The positions of components are usually represented by a lattice. • lattice is a three dimensional system of units repeating in a pattern. The smallest repeating unit of the lattice is called the unit cell.

  36. Three types of Crystalline Solids

  37. Amorphous Solids Amorphous solids have considerable disorder in their structures. • Example: Common glass looks like a solution frozen in place. It has a rigid shape but a great deal of disorder within its structure.

  38. X-ray Analysis of Solids The structures of crystalline solids are commonly determine by X-ray diffraction. • This type of diffraction occurs when beams of light are scattered as they go through spaces between substances. Light scatters when the size of the spaces are similar to the wavelength of light.

  39. X-ray Analysis of Solids

  40. X-ray Analysis of Solids • A single wavelength is directed at the crystal and a diffraction pattern is obtained. The diffraction pattern is a series of light and dark areas on a photographic plate from constructive and destructive interference from waves of light. • The diffraction pattern can then be used to determine the interatomic spacings.

  41. X-ray Analysis of Solids • A diffractometer is a computer-controlled instrument used for carrying out the X-ray analysis of crystals • It rotates the crystal with respect to the X-ray beam and collects the data produced by the scattering. The techniques have been refined to the point that very complex structures can be determined, such as large biological enzymes.

  42. X-ray Analysis of Solids The Bragg equation combines trigonometry and physics to determine the atomic spaces between crystals: • nλ = 2d sin θ • d is the distance between atoms and θ is the angle of incidence and reflection of the light. n is an integer, most commonly 1. (n is usually given)

  43. X-ray Analysis of Solids

  44. Example Problem X-rays of wavelength 1.54 Â were used to analyze an aluminum crystal. A reflection was produced at θ = 19.3°. Assuming n=1, calculate the distance d between the planes of atoms producing this reflection 2.33 Á

  45. Types of Solids • Ionic solids • ionic solids are made of ions • Molecular solids • Molecular solids have small units of covalently bonded molecules. • Atomic solids • Atomic solids are made of elements such as carbon (graphite, diamond and the fullerenes), boron, silicon, and all metals.

  46. Fullerenes

  47. Types of Solids

  48. Atomic Solids Atomic solids are broken down into subgroups depending on the bond that exists in the solid: • Metallic solids • Has delocalized nondirectional covalent bonding. • Network solids • atoms bond with strong directional covalent bonding that lead to giant molecules and networks

  49. Atomic Solids • Group 8A solids • noble gases are attracted to each other with London dispersion forces.

  50. Classification of Solids