1 / 15

Liquids and Solids

Liquids and Solids. Phase Changes . Heating Curve: Heat of Fusion From solid to liquid or liquid to solid Heat of Vaporization From gas to liquid or liquid to gas Always larger than heat of fusion Heat of sublimation is the sum of fusion and vaporization

belva
Download Presentation

Liquids and 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. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Liquids and Solids

  2. Phase Changes • Heating Curve: • Heat of Fusion • From solid to liquid or liquid to solid • Heat of Vaporization • From gas to liquid or liquid to gas • Always larger than heat of fusion • Heat of sublimation is the sum of fusion and vaporization • Stronger IMF= larger heat of fusion and vaporization values (need more energy to overcome phases)

  3. Phase change problem • Calculate the energy change (enthalpy change) upon converting 1.00 moles of ice at -25 degrees Celsius to water vapor at 125 degrees Celsius under constant pressure. Specific heat of ice = 2.03 J/g-K, specific heat of water = 4.184 J/g-K, specific heat of water vapor = 1.84 J/g-K. For water, heat of fusion = 6.01 kJ/mol, and the heat of vaporization = 40.67 kJ/mol • Answer = 56 kJ

  4. Phases Change VOCABULARY • Critical Temperature = the highest temp at which a liquid can exist • Higher than the crit temp the kinetic energy is so great that the substance must be a gas • Critical Pressure = the pressure required to liquify at the critical temperature • Nonpolar (weak IMF)= low crit temps and pressures • Polar (strong IMF) = high crit temps and pressures

  5. Vapor Pressure • Vapor Pressure = The partial pressure exerted by a vapor in a closed system when it is at equilibrium at the solid or liquid phase. (When evaporation and condensation are at equilibrium) • Weak intermolecular forces and low molar masses have HIGH vapor pressures (because more will evaporate) • Increasing temp = increase vapor pressure (b/c more vapor!!)

  6. Boiling • The boiling point = the temp at which the vapor pressure equals the atmospheric pressure • The normal boiling point of a liquid is when the vapor pressure equals one atmosphere • (DEMO of water in vacuum) • Denver = low atmospheric pressure. • Low pressure = water will boil at a lower temp

  7. Phase Diagrams • Solid lines represent the temps and pressures where the phases of the substance are in equilibrium • Notice that ice melts under pressure (because of density) Phase diagram for Water (Line between liquid and solid is tilted toward the left- only substance to do this because solid is less dense than liquid)

  8. Phase diagram cont. • Where is critical point on phase diagram? • Represents both critical pressure and temp • After the critical point = supercritical fluid (where gas phase and liquid phase are indistinquishable) • TRIPLE POINT- where all phases exist in equilibrium

  9. Solids • Amorphous = no ordered structure • Don’t melt at specific temps because different intermolecular forces • Crystalline solid = ordered structure • Melt at specific temps because of consistent intermolecular forces

  10. Crystalline Structures • Unit Cell = the smallest unit in a crystal (The repeating pattern) • Crystal Lattice = an imaginary network of points on which the unit cells can be laid down so that the structure of the crystal is obtained.

  11. Common Unit Cells • Primitive Cubic = corners only • Body-centered cubic = corners and 1 in center • Face-centered cubic = corners and in the center of each side

  12. Bonding in Solids 1. Molecular Solids • Held together by intermolecular forces • Unit particle = atom or molecule • Generally the lowest melting point of all types of solids • Generally softer solids because of weaker IMF • Most molecular solids are gases or liquids at room temp • Examples: water, carbon dioxide, helium, sugar

  13. Bonding in solids cont 2. Covalent-Network Solids – • Held together with covalent bonds • HIGH melting points (highest of all types) • Very hard • Unit particle: Atoms connected in a network of covalent bonds • Examples: Diamond, quartz (SiO2)

  14. Bonding in solids cont 3. Ionic Solids – • Held together by ionic bonds (Lattice Energy) • Also high melting points, not quite as high as covalent-network solids • Hard and brittle • Unit particle: positive and negative ions • Examples: NaCl, Ca(NO3)2

  15. Bonding in solids cont. 4. Metallic Solids – (Metals) – • Held together by metallic bonds • delocalized valance electrons (a sea of electrons). • Unit particle: 1 atom • Soft to hard • Low to very high melting point • Excellent thermal and electrical conduction • Malleable • Ductile • Examples: Any metallic element (Cu, Fe, Al, Pt, Au, etc)

More Related