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Liquids and Solids

Liquids and Solids. Gas low density high compressibility completely fills its container Solid high density only slightly compressible rigid maintains its shape. Liquids and Solids. Liquids properties lie between those of solids and gases

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Liquids and Solids

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  1. Liquids and Solids • Gas • low density • high compressibility • completely fills its container • Solid • high density • only slightly compressible • rigid • maintains its shape

  2. Liquids and Solids • Liquids • properties lie between those of solids and gases • H2O(s) --> H2O(l) DHofus = 6.02 kJ/mol • H2O(l) --> H2O(g) DHovap = 40.7 kJ/mol • large value of DHvap suggests greater changes in structure in going from a liquid to a gas than from a solid to liquid • suggests attractive forces between the molecules in a liquid, though not as strong as between the molecules of a solid

  3. Liquids and Solids • Densities of the three states of water • H2O(g) D = 3.26 x 10-4g/cm3 (400oC) • H2O(l) D = 0.9971 g/cm3 (25oC) • H2O(s) D = 0.9168 g/cm3 (OoC) • Similarities in the densities of the liquid and solid state indicate similarities in the structure of liquids and solids

  4. Intermolecular Forces • Bonds are formed between atoms to form molecules • intramolecular bonding (within the molecule)

  5. Intermolecular Forces • The properties of liquids and solids are determined by the forces that hold the components of the liquid or solid together • may be covalent bonds • may be ionic bonds • may weaker intermolecular forces between molecules

  6. Intermolecular Forces • During a phase change for a substance like water • the components of the liquid or solid remain intact • the change of state is due to the changes in the forces between the components • e.g., H2O(s) --> H2O (l) …the molecules are still unchanged during the phase change

  7. Dipole-Dipole Forces • Polar molecules • line up in an electric field • positive end of molecule will line up with the negative pole of the electric field while the negative end of the molecule will line up with the positive pole • can attract each other • positive end of one molecule will attract the negative end of another molecule

  8. Dipole-Dipole Forces • Dipole-dipole forces • about 1% as strong as covalent or ionic bonds • become weaker with distance • unimportant in the gas phase

  9. Hydrogen Bonding • A particularly strong dipole-dipole force • When hydrogen is covalently bonded to a very electronegative atom such as N, O, or F • Very strong due to • great polarity of the bond between H and the N, O or F • close approach of the dipoles due to H’s small size

  10. Hydrogen Bonding • H-bonding has a very important effect on physical properties • For example, boiling points are greater when H-bonding is present

  11. London Dispersion Forces • aka Van der Waals forces • Nonpolar molecules must exert some kind of force or they would never solidify

  12. London Dispersion Forces • London dispersion forces (LDF) • due to an instantaneous dipole moment • created when electrons move about the nucleus • a temporary nonsymmetrical electron distribution can develop (I.e., all the electrons will shift to one side of the molecule)

  13. London Dispersion Forces • The instantaneous dipole moment can induce an instantaneous dipole moment in a neighboring molecule, which could induce another instantaneous dipole moment in a neighboring molecule, etc. (like a “wave” in the stands of a football game)

  14. London Dispersion Forces • The LDF is very weak and short-lived • To form a solid when only LDF exists requires very low temperatures • the molecules or atoms must be moving slowly enough for the LDF to hold the molecules or atoms together in a “solid” unit

  15. London Dispersion Forces • Element Freezing Point (oC) Helium -269.7 Neon -248.6 Argon -189.4 Krypton -157.3 Xenon -111.9

  16. London Dispersion Forces • Notice that as the MM of the noble gas increases, the freezing point increases • This implies that the LDF between the atoms is stronger as the MM increases • Large atoms with many electrons have an increased polarizability (the instantaneous dipole would be larger), resulting in a larger London Dispersion Force between the atoms than between smaller atoms

  17. The Liquid State • Properties of liquids • low compressibility • lack of rigidity • high density (compared to gases)

  18. The Liquid State • Surface Tension • results in droplets when a liquid is poured onto a surface • depends on IMF’s

  19. The Liquid State • Molecules at the surface experience an uneven pull, only from the sides and below. Molecules in the interior are surrounded by IMF’s • Uneven pull results in liquids assuming a shape with minimum surface area • Surface tension is a liquids resistance to an increase in surface area. • Liquids with high IMF’s have high surface tensions

  20. The Liquid State • Capillary Action • Exhibited by polar molecules • The spontaneous rising of a liquid in a narrow tube • due to two different forces involving the liquid

  21. The Liquid State • Cohesive forces - IMF between the liquid molecules • Adhesive forces - forces between the liquid molecules and the polar (glass) container • adhesive forces tend to increase the surface area • cohesive forces counteract this • Concave meniscus (water) - indicates adhesive forces of water towards the glass is greater than the cohesive forces between the water molecules. • Convex meniscus (nonpolar substances such as mercury) shows cohesive forces is greater than adhesive forces.

  22. The Liquid State • Viscosity • Measure of a liquid’s resistance to flow • Depends on strength of IMF’s between liquid molecules • molecules with large IMF’s are very viscous • Large molecules that can get tangled up with each other lead to high viscosity

  23. The Liquid State • So what does a liquid “look like?” • A liquid contains many regions where the arrangements of the components are similar to those of a solid • There is more disorder in a liquid than in a solid • There is a smaller number of regions in a liquid where there are holes present

  24. Types of Solids • Ways to classify solids • Crystalline vs. Amorphous Solid • Crystalline solids • regular arrangement of components • positions of components represented by a lattice • unit cell - smallest repeating unit of the lattice

  25. Types of Solids • three common unit cells exist • simple cubic • body centered cubic • face centered cubic

  26. Types of Solids • Amorphous Solids • noncrystalline • glass is an example • disorder abounds

  27. Types of Solids • X-ray diffraction • used to determine the structures of crystalline solids • diffraction occurs when beams of light are scattered from a regular array of points • obtain a diffraction pattern • Bragg equation: nl = 2d sinq

  28. Types of Solids • Where n is an integer • l is the wavelength of the x-rays • d is the distance between the atoms • q is the angle of incidence and reflection • Use x-ray diffraction to determine bond lengths, bond angles, determine complex structures, test predictions of molecular geometry

  29. Types of Solids • Example: • x-rays of wavelength 1.54 A were used to analyze an aluminum crystal. A reflection was produced at q = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing the reflection. • (D = 2.33 A)

  30. Types of Solids • Types of Crystalline Solids • Ionic Solids (e.g. NaCl) • Molecular Solids (e.g. C6H12O6) • Atomic Solids which include: • Metallic Solids • Covalent Network Solids

  31. Types of Solids • Classify solids according to what type of component is found at the lattice point (of a unit cell) • Atomic Solids have atoms at the lattice points • Molecular Solids have discrete, relatively small molecules at the lattice points • Ionic solids have ions at the lattice points

  32. Types of Solids • Different bonding present in these solids results in dramatically different properties • Element (atomic solid) M.P. (oC) Argon -189 C(diamond) 3500 Cu 1083

  33. Structure and Bonding inMetals • Properties of Metals • high thermal conductivity • high electrical conductivity • malleability (metals can be pounded thin) • ductility (metals can be drawn into a fine wire) • durable • high melting points

  34. Structure and Bonding inMetals • Properties are due to the nondirectional covalent bonding found in metallic crystals • Metallic crystal • contains spherical atoms packed together • atoms are bonded to each other equally in all directions

  35. Structure and Bonding inMetals • Closest Packing • most efficient arrangement of these uniform spheres • Two possible closest packing arrangements • Hexagonal Closest Packed Structure • Cubic Closest Packed Structure

  36. Structure and Bonding inMetals • Hexagonal Closest Packed Structure (hcp) • aba arrangement • First Layer • each sphere is surrounded by six other spheres

  37. Structure and Bonding inMetals • Second Layer • the spheres do not lie directly over the spheres in the first layer • the spheres lie in the indentations formed by three spheres • Third Layer • the spheres lie directly over the spheres in the first layer

  38. Structure and Bonding inMetals • Cubic Closest Packed Structure (ccp) • abc arrangement • First and Second Layers are the same as in hexagonal closest packed structure • Third Layer • the spheres occupy positions such that none of the spheres in the third layer lie over a sphere in the first layer

  39. Structure and Bonding inMetals • Finding the net number of spheres in a unit cell • important for many applications involving solids (when I figure it out, I’ll let you know…or when it shows up on the ACS or AP test…then I’ll figure it out!)

  40. Structure and Bonding inMetals • Examples of metals that are ccp • aluminum, iron, copper, cobalt, nickel • Examples of metals that are hcp • zinc, magnesium • Calcium and some other metals can go either way

  41. Structure and Bonding inMetals • Some metals, like the alkali metals are not closest packed at all • may be found in a body centered cubic (bcc) unit cell where there are only 8 nearest neighbors instead of the 12 in the closest packed structures

  42. Bonding Models for Metals • The model must account for the typical physical properties of metals • malleability • ductility • efficient and uniform conduction of heat and electricity in all directions • durability of metals • high melting points

  43. Bonding Models for Metals • To account for these physical properties, the bonding in metals must be • strong • nondirectional • It must be difficult to separate atoms, but easy to move them (as long as the atoms stay in contact with each other

  44. Bonding Models for Metals • Electron Sea Model (simplest picture) • Positive Metal ions (Metal cations) are surrounded by a sea of valence electrons • the valence electrons are mobile and loosely held • these electrons can conduct heat and electricity • meanwhile, the metal ions can move around easily

  45. Bonding Models for Metals • Band Model or Molecular Orbital (MO) model • related to the electron sea model • more detailed view of the electron energies and motions

  46. Bonding Models for Metals • MO model • electrons travel around the metal crystal in molecular orbitals formed from the atomic orbitals of the metal atoms • In atoms like Li2 or O2, the space between the energies of the molecular orbitals is relatively wide (big energy difference between the orbitals)

  47. Bonding Models for Metals • However, when many metal atoms interact, the molecular orbital energy levels are very close together • Instead of separate, discrete molecular orbitals with different energies, the molecular orbitals are so close together in energies, that they form a continuum of levels, called bands

  48. Bonding Models for Metals • Core electrons of metals are localized • the core electrons “belong” to a particular metal ion • The valence electrons of metals are delocalized • the valence electrons occupy partially filled, closely spaced molecular orbitals

  49. Bonding Models for Metals • Thermal and Electrical conductivity • metals conduct heat and electricity because of highly mobile electrons • electrons in filled molecular orbitals get excited (from added heat or electricity) • these electrons move into higher energy, empty molecular orbitals

  50. Bonding Models for Metals • Conduction electrons • the electrons that can be excited to empty MO’s • Conduction bands • the empty MO’s that can accept the conducting electrons

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