1 / 39

Liquids and Solids

Liquids and Solids. “CONDENSED STATES OF MATTER”. intramolecular forces inside molecules hold atoms together into molecule intermolecular forces between molecules get weaker as phase changes from S – L – G Generally, inter molecular forces are much weaker than intra molecular forces.

oro
Télécharger la présentation

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 “CONDENSED STATES OF MATTER”

  2. intramolecular forces inside molecules hold atoms together into molecule intermolecular forces between molecules get weaker as phase changes from S – L – G Generally, intermolecular forces are much weaker than intramolecular forces. Intra- vs. Inter-

  3. Dipole-Dipole Forces Orientation of Polar Molecules in a Solid Attractive forces between polar molecules

  4. happens between H and N, O, or F very strong type of dipole-dipole attraction because bond is so polar because atoms are so small Hydrogen Bonding

  5. Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

  6. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. • increase with the size of the molecules. • in every molecular compound • only important for nonpolar molecules and noble gas atoms • weak, short-lived • caused by formation of temporary • dipole moments Fritz London 1900-1954

  7. London Dispersion Forces

  8. Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Strongest Weakest Covalent bonds (400 kcal/mol) (intramolecular) Hydrogen bonding (12-16 kcal/mol) (intermolecular) Dipole-dipole interactions (2-0.5 kcal/mol) (intermolecular) Londonforces (less than 1 kcal/mol) (intermolecular)

  9. London Forces in Hydrocarbons Dispersion forces usually increase with molar mass.

  10. Boiling point as a measure of intermolecular attractive forces

  11. Practice • which has highest boiling pt? • HF, HCl, or HBr? • Identify the most important intermolecular forces : • BaSO4 • H2S • Xe • C2H6 • P4 • H2O • CsI HF ionic dipole-dipole London Dispersion London Dispersion London Dispersion H-bonding ionic

  12. CO2 or OCS CO2: nonpolar so only LD OCS: polar so dipole-dipole PF3 or PF5 PF3: polar so dipole-dipole PF5: nonpolar so only LD Which has stronger IMF’s? • SF2 or SF6 • SF2: polar so dipole-dipole • SF6: nonpolar so only LD • SO3 or SO2 • SO3: nonpolar so LD only • SO2: polar so dipole-dipole ? ? ? ?

  13. DO WE GET IT? Please begin the following problems. You may work cooperatively on this assignment. Problems: pg. 475 #12, 20, 30, 32, 34 and 36 (a,b and d only!) To be collected!

  14. Some Properties of a Liquid • Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals). Strong intermolecular forces High surface tension =

  15. EMOO BUG Surface of water behaving like it had an “elastic skin”

  16. Adhesion • Capillary Action: Spontaneous rising of a liquid in a narrow tube. Cohesion is the intermolecular attraction between like molecules Cohesion Adhesion is an attraction between unlike molecules attracted to glass attracted to each other

  17. Properties of a Liquid • Viscosity: Resistance to flow • High viscosity is an indication of strong intermolecular forces

  18. Which of the following consistently have the highest melting points? • Metals • Salts • Molecular crystals • Alkanes • Hydrogen-bonded compounds. Answer: B

  19. How sharp are you? Gases can be compressed more easily than liquids because: • Gas molecules are smaller than liquid molecules B. The kinetic energy of gas molecules is higher than that found in liquids • The average intermolecular distances are greater in gases than those found in liquids. • Intermolecular forces increase as gas moleculues are brought closer together. E. None of the above. Answer: C In a gas, the molecules are separated by a large distance and are able to be compressed by increasing the pressure. They move independently of one another because there is no appreciable intermolecular interactions among them.

  20. In which of the following are the intermolecular forces listed from the weakest to the strongest? • Dipole-dipole>London>hydrogen bonds • London<dipole-dipole<hydrogen bonds • Hydrogen bonds<dipole-dipole<London • London>hydrogen bonds>dipole-dipole • London>Javier bonds>dipole-dipole>Ali forces Answer: B

  21. Which of the following compounds will NOT hydrogen- bond>? • CF4 • CH3OH • H2NCH2CH2CH3 • HOCH2CH2OH • HClO Answer: A

  22. Water has a higher capillary action than mercury due to: • Higher dipole-dipole forces between the water molecules • Strong cohesive forces within water. • Very significant induced intermolecular attractions. • Weak adhesive forces in water • Strong cohesive forces in water which work with strong adhesive forces. Answer: E. The strong adhesive forces leads to a creeping effect as water moves up the narrow tubing and the strong cohesive forces attempt to minimize the surface area.

  23. Small drops of water tend to bead up because of: • High capillary action • the shape of the meniscus • The resistance to increased surface area. • Low London dispersion forces • Weak covalent bonds. Answer: C. This is a description of surface tension, which is a result of high dipole-dipole forces between water molecules. These intermolecular forces are also called…..hydrogen bonds!

  24. Several liquids are compared by adding them to a series of 50 mL graduated cylinders, then dropping a steel ball of uniform size and mass into each. The time required for the ball to reach the bottom of the cylinder is noted. This is a method used to compare the differences in a property of liquids known as: • Surface tension • Buoyancy • Capillary action • Viscosity • Surface contraction Answer: D. The resistance to flow of any fluid is called viscosity. As You would predict, liquids with high viscosity (ex: maple syrup) have large intermolecular forces.

  25. General Classification of Solids Crystalline Solids: Well-ordered, definite arrangement of atoms. (Examples- metals, H2O, diamond) • Amorphous: No pattern to the arrangement of particles. (Examples- glass, plastic, wax)

  26. Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

  27. Types of Crystalline Solids There are four types of crystalline solid: - Molecular (formed from atoms or molecules) - usually soft with low melting points and poor conductivity. - Covalent network (formed from atoms) - very hard with very high melting points and poor conductivity. - Ionic (formed form ions) - hard, brittle, high melting points and poor conductivity. - Metallic (formed from metal atoms) - soft or hard, high melting points, good conductivity, malleable and ductile.

  28. Know these!! Bonding in Crystalline Solids • Metallic bonds are formed from metal nuclei floating in a sea of electrons.

  29. Crystalline Solids Molecular Covalent Network Ionic Metallic

  30. Metal Alloys • Substitutional Alloy: some metal atoms replaced by others of similar size. • brass = Cu/Zn

  31. Metal Alloys(continued) • Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. • steel = iron + carbon

  32. Phase Changes & Energy Endothermic: melting, evaporating/boiling & sublimation Exothermic: freezing, condensation, & deposition

  33. Phase Changes & Energy • heat of vaporization: the heat energy required to evaporate a given mass of liquid at a constant temperature • heat of fusion: the heat energy required to melt a given mass of solid at a constant temperature • The temperature, (average KE), during a phase change (such as boiling) does not change! • Any heat added during boiling gives more molecules enough energy to escape the liquid. Heating Curve

  34. Phase Changes & Energy • Generally, it will take more heat to vaporize a liquid than to melt a solid… (∆H(vap) > ∆H(fusion) ) Why? • Every intermolecular bond is broken when vaporizing, but only some of the intermolecular forces break when melting solids. ?

  35. Practice Time: Please attempt the following problems (at your workstations!): pg. 478-80: 72, and 88.

  36. Liquefying Gases Gases can be liquefied by: • increasing pressure at some temperature. • decreasing the temperature at some pressure. - Critical temperature: the highest temperature at which a substance can remain a liquid regardless of the pressure applied. - Critical pressure: the pressure needed at the critical temperature.

  37. Phase Diagrams • Shows the relationship between the 3 phases of matter at various temperatures and pressures. Triple Point: All 3 phases of matter at equilibrium. Critical Point: The highest temperature at which the liquid phase can exist.

  38. Phase Diagrams of H2O and CO2 • Notice the slope of the solid–liquid equilibrium line.

  39. Practice Time: Please attempt the following problems. You may scatter and work with anyone you please. pg. 480: #91.

More Related