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CHEM PAC 13

CHEM PAC 13. Liquids, Solids, and Water. Kinetic Molecular Theory Description - Liquids. Molecules of a liquid are always moving because they possess kinetic energy .

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CHEM PAC 13

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  1. CHEM PAC 13 Liquids, Solids, and Water

  2. Kinetic Molecular Theory Description - Liquids • Molecules of a liquid are always moving because they possess kinetic energy. • Molecules of a liquid are almost as close together as they can possibly be so liquids are denser than gases and practically non-compressible.

  3. Kinetic Molecular Theory Description - Liquids • Molecules of a liquid are held together by attractive forces (more so than gases) so liquids have definite volume and can have one free surface. • Molecules of a liquid move with respect to one another so liquids are fluid and take the shape of their container.

  4. Properties of liquids • definite volume (KMT 3) • fluidity(takes the shape of its container, can flow) (KMT 4) • relatively high density(particles close together) (KMT 2) • relative noncompressibility (particles close together) (KMT 2) • dissolving ability(coming in Pac 14!!) (KMT 1+ IMF) • ability to diffuse(slower than gases) (KMT 1 + 2 + IMF)

  5. Properties of liquids • surface tension(particles on the surface pull together, can support some solids, droplets are spherical, explains capillary action) (KMT 3) • tendency to evaporate and to boil (KMT 1) ( liquid + energy gas) • tendency to solidify ( liquid solid + energy)

  6. Predicting phases of substances

  7. Kinetic Molecular Theory Description - solids • Molecules of a solid are always moving but much less than gases because they possess less kinetic energy. Their movement is essentially vibrational. • Molecules of a solid are as close together as they can possibly be so solids have relatively high density and are said to be noncompressible.

  8. Kinetic Molecular Theory Description - solids • Molecules of a solid are held together tightly by attractive forces so solids have definite shape and volume and can have all free surfaces.

  9. Amorphous solids - lack an orderly arrangement of particles ex. glass, paraffin • definite shape (temporary, changes over time) • definite volume • fluid (resemble liquids which flow very slowly) • lack definite melting point • high density • noncompressible • very slow diffusion rate

  10. Crystalline solids - particles arranged in an orderly, geometric, repeating pattern ex. most solids • definite shape (geometric) • definite volume • nonfluid • definite melting point • high density • noncompressible • very slow diffusion rate (Want proof?)

  11. Proof that solids diffuse The evidence that solids diffuse came in the form of a very simple but conclusive experiment. After two different metal bars are fixed in place for an extended period of time, traces of each metal are found in the other bar.

  12. Some terminology • CRYSTAL: a homogeneous portion of a substance bounded by plane surfaces making definite angles with each other, giving a regular geometric form. • CRYSTAL LATTICE: the pattern of particle arrangement. • UNIT CELL: the smallest portion of the crystal lattice which exhibits the pattern of the lattice structure.

  13. Crystals form when... • Solutions of chemical compounds evaporate. • Hot, saturated solutions cool. • Change from liquid to solid phase (freezing). • Change from gas to solid phase (opposite of subliming). • Ions, atoms or molecules arrange themselves in positions of least energy.

  14. A Virtual Field Trip to aNatural History Museum Images from The Smithsonian National Mining Museum Western Museum of Mining

  15. Gold nugget & Hope Diamond

  16. Molecular solids • arrangement of particles (unit cell) individually distinct molecules • attractive forces nonpolar molecules: d-i forces polar molecules: both d-i & dipole-dipole forces • melting point: low • volatile(easily vaporized) • hardness: soft • flexibility: brittle • conductor: no EXAMPLES: I2, CO2, H2O

  17. Covalent network solids • arrangement of particles (unit cell) array of atoms sharing electrons • attractive forces strong covalent bonds in fixed directions • melting point: high • nonvolatile(not easily vaporized) • hardness: hard • flexibility: brittle • conductor: no EXAMPLES: diamond, -SiO4, oxides of transition metals

  18. ionic solids • arrangement of particles (unit cell) positive and negative ions • attractive forces strong binding forces from attraction between ions • melting point: high • nonvolatile(not easily vaporized) • hardness: hard • flexibility: brittle • conductor: no EXAMPLES: NaCl, MgF2, CaCO3

  19. metallic solids • arrangement of particles (unit cell) electron sea model • attractive forces strong binding forces from metallic bonds • melting point: very high • nonvolatile(not easily vaporized) • hardness: hard • flexibility: flexible • conductor:yes EXAMPLES: Na, Fe, Cu, Ag

  20. How do metals conduct electricity? The electron sea model shows the positive nuclei of metal atoms surrounded by electrons which are free to migrate through the crystal lattice.

  21. Equilibrium & Le Chatelier’s Principle Equilibrium is a dynamic state in which two opposingprocesses occur at equal rates within a closed system.

  22. Equilibrium & le chatelier’s principle In time within the closed beaker, the number of molecules moving between the liquid and gas phases will be equal and the processes of evaporation and condensation will be at equilibrium.

  23. Equilibrium & le chatelier’s principle Remember at equilibrium: • There is no net change in the number of liquid or vapor molecules. • Instead, molecules will continue to move, exchanging places liquid vapor vapor liquid

  24. Equilibrium & le chatelier’s principle Evaporation can be represented as: liquid + energy vapor Condensation can be represented as: vapor liquid + energy So the system at equilibrium can be represented as: liquid + energy vapor

  25. So how does evaporation occur? • Some liquid molecules acquire enough kinetic energy (through elastic collisions) to escape from the surface to the vapor phase. • Evaporation occurs without heating. • Evaporation takes place only at the surface.

  26. Equilibrium & le chatelier’s principle 1888 - French Chemist Le Chatelier recognized that changes which alter either the rate of a forward or reverse reaction in an equilibrium system disrupt that system. However, the system responds to this stress upon it by shifting in the direction which relieves this stress. For example...

  27. Equilibrium & le chatelier’s principle If the beaker on the left represents the system in equilibrium: liquid + energy vapor what would happen to the system if the temperature is increased by 10oC? The rate of evaporation would exceed condensation and we say equilibrium shifts to the right. In time, equilibrium will be re-established at this new higher temperature.

  28. Equilibrium & le chatelier’s principle If the beaker on the left represents the system in equilibrium: liquid + energy vapor what would happen to the system if the temperature is decreased by 10oC? The rate of condensation would exceed evaporation and we say equilibrium shifts to the left. In time, equilibrium will be re-established at this new lower temperature.

  29. Equilibrium & le chatelier’s principle We can now improve on a definition from earlier chapters… vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid and it depends on the nature of the liquid (attractive forces between molecules) and on the temperature of the liquid (kinetic energy of the molecules).

  30. Phase changes - boiling Heating liquid molecules causes their kinetic energy to increase. Bubbles will form on the bottom. • What is in these bubbles? Vapor of the liquid • Where/Why do they form? They form where heating is strongest. The vapor bubble can only exist if it has enough pressure to overcome the pressure exerted on it from all sides.

  31. Phase changes - boiling As the vapor pressure gets closer to atmospheric pressure, the bubble is able to move closer to the surface of the liquid, ready to exit to the atmosphere. (This is the slow bubbling seen prior to boiling.) When the vapor pressure increases with the water temperature, a temperature is reached at which the equilibrium vapor pressure is equal to the pressure of the atmosphere acting on the surface of the liquid. This temper-ature is called the boiling point.

  32. How is boiling different from evaporation? • Boiling occurs with heating. • Boiling does not only occur from the surface. • Boiling depends on the atmospheric pressure.

  33. Phase changes - boiling T 100oC E M boiling point of water P 0oC TIME At 1.00atm water boils at 100oC (373K). DURING BOILING THE TEMPERATURE REMAINS CONSTANT heat of vaporization: the heat energy required to vaporize one mole of a liquid at its standard boiling point. liquid + energy vapor (heat of vaporization) (gas)

  34. Phase changes - freezing and melting • freezing point: the temperature at which a liquid changes phase to solid liquid solid + energy • melting point: the temperature at which a solid changes phase to liquid solid + energy liquid • for pure crystalline solids, the temperatures at which these two processes occur coincide and this can be shown as an equilibrium solid + energy liquid (heat of fusion) • heat of fusion: heat energy required to melt one mole of a solid at its melting point

  35. Phase changes __A__ solid _ABC_ melting __B__ melting point __C__ liquid _CBA_ freezing __B__freezing point __E__gas _CDE_ boiling __D__boiling point _CDE_ evaporating _EDC_ condensing __D__heat of vaporization __B__heat of fusion

  36. Phase changes - subliming sublimation: the solid “skips” the liquid phase and passes directly to the gas phase. solid + energy vapor (gas) The phase change diagram for such a substance would look like: T gas E deposition M sublimation P solid TIME

  37. water Intramolecular forces within water molecules: water molecules _ are O polar (dipoles) H + H water molecules contain polar bond angle: 105ocovalent bonds

  38. water Intermolecular forces between water molecules: _ O * 1. London dispersion _ H + H forces O 2. Dipole-dipole forces H + H 3. Hydrogen bonding * All of these attractive forces result in an unusually high boiling point for such a low molecular mass polar molecule.

  39. water Hydrogen bonds in liquid water are longer and more flexible than those in ice. Therefore, water (l) is denser than water (s) and ice can float on water. 0 oC 4 oC 100 oC H-bonds begin MAXIMUM kinetic energy increases, to break, water DENSITY molecules overcome molecules (1.00 g/mL) attractive forces and move closer move apart SOLID | LIQUID | GAS

  40. Hydrates - water trapped in a crystal lattice • hydrate: a crystalline substance with its water(s) of crystallization. ex. CuSO4 . 5H2O copper (II) sulfate pentahydrate • efflorescence: loss of water(s) of crystal-lization when a hydrate is heated or exposed to air • deliquescence: uptake of water molecules from the air into the crystal lattice.

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