CHEM 200/202. Professor Gregory P. Holland Office: GMCS-213C. All s are to be sent to:

Transcription

1 CHEM 200/202 Professor Gregory P. Holland Office: GMCS-213C All s are to be sent to: My office hours will be held in GMCS-212 on Monday from 12:00 pm to 2:00 pm or by appointment.

2 SUPPLEMENTAL INSTRUCTION Study sessions lead by former CHEM 200/202 students that excelled in the previous semesters class (18 sessions/ wk). Free to access, no reporting to faculty.

3 LECTURES, HOMEWORK & EXAM 4 (CHAPTERS 1-11) Homework 9 due December 7th Homework 10 & 11 due December 14th Quiz 4 is due December 13th Exam 4, Final is December 15th at 7:30 PM Note: OWL HW and Quizzes are due by 11:55 PM on due date

4 LECTURE OBJECTIVES Chapter 10.4 Phase Diagrams Explain the construction and use of a typical phase diagram. Use phase diagrams to identify stable phases at given temperatures and pressures, and to describe phase transitions resulting from changes in these properties. Describe the supercritical fluid phase of matter. Chapter 10.5 & 10.6 The Solid State of Matter and Lattice Structures Explain the main types of crystalline solids and their bonding. Describe crystal defects. Describe the arrangement of atoms and ions in crystal structures. Compute ionic radii using unit cell dimensions.

5 PHASE DIAGRAMS Chapter 10.4

6 PHASE DIAGRAMS Phase diagrams depict how a substance changes its physical state as a function of temperature and pressure. Phase diagrams can be used to predict at what temperature and pressure certain transitions will occur, or if they can occur.

7 PHASE DIAGRAM OF WATER The phase diagram for water reveals a unique property of water, that its solid is less dense then the liquid. Seen in a negative slope from B-D. Liquid water is favored at higher pressures because it can compact more (more mass per volume).

8 PHASE DIAGRAM OF CO2 Supercritical fluids (SCF) occur when the gas and liquid phase have the same density. SCF have unique properties, similar to a gas, but with solubility like a liquid.

9 PROBLEM Examine the phase diagram for the substance X and select the correct statement. 1.X(s) has a lower density than X(l). 2.The triple point for X is at a higher temperature than the melting point for X. 3.X changes from a solid to a liquid as one follows the line from C to D. 4.Point B represents the critical temperature for X.

10 THE SOLID STATE OF MATTER Chapter 10.5

11 CHEMICAL SOLIDS When chemicals are cooled sufficiently, the molecules will ultimately form a solid. Crystalline solids arise when the molecules/ ions/atoms are arranged in a defined repeating pattern. Amorphous solids (noncrystalline) can arise when liquids freeze before the molecules/ atoms/ions can be arranged in a defined patter. Metals & ionic compounds are good at forming crystalline structures. Large molecules & mixtures tend to form amorphous solids.

12 of the metals varymolecules widely. Mercury is a such liquid as indicated by the melting points points of the crystals. Small symmetrical (nonpolar molecules), as Hat 2, room temperatur N2, O2, and F2, have weak attractive forces and form molecular solids with very lowhave meltinglow points (below 200 Several post-transition metals also melting points, whereas the C). Substances consisting of larger, nonpolar molecules have larger attractive forces and melt at higher temperatures C. These differences differences strengths of metallic bo Molecular solids composed of molecules with permanent dipole momentsreflect (polar molecules) melt atin still higher Figure 10.40Examples Copper is a metallic solid. temperatures. include ice (melting point, 0 C) and table sugar (melting point, 185 C). free to move. Many simple compounds formed by the reaction of a metallic onic. TYPES OF CRYSTALLINE SOLIDS Covalent Network Solid Covalent network solids include crystals of diamond, silicon, some other nonmetals, and some covalent compound such as silicon dioxide (sand) and silicon carbide (carborundum, the abrasive on sandpaper). Many minerals hav networks of covalent bonds. The atoms in these solids are held together by a network of covalent bonds, as shown i Figure To break or to melt a covalent network solid, covalent bonds must be broken. Because covalent bond are relatively strong, covalent network solids are typically characterized by hardness, strength, and high meltin points. For example, diamond is one of the hardest substances known and melts above 3500 C. solid. Ionic Molecular Solid Metallic Figure Carbon dioxide (CO2) consists of small, nonpolar molecules and forms a molecular solid with a melting point of 78 C. Iodine (I2) consists Figure of larger, nonpolar and is forms a molecular solid. solid that melts at 114 C molecules Copper a metallic Properties of Solids Covalent Network Solid A crystalline solid, like those listed in Table 10.4, has a precise melting temperature because each atom or molecule of the same type is held in place with the same forces or energy. Thus, the attractions between the units that make up the crystal all have the same strength and all require the same amount of energy to be broken. The gradual softening Covalent network solids include crystals of diamond, silicon, some othe of an amorphous material differs dramatically from the distinct melting of a crystalline solid. This results from the such as silicon dioxide (sand) silicon carbide (carborundum, the ab structural nonequivalence of the molecules in the amorphous solid. Some forcesand are weaker than others, and when an amorphous material is heated, the weakest intermolecular attractions breakthe first. atoms As the temperature increasedare held together b networks of covalent bonds. in theseis solids further, the stronger attractions are broken. Thus amorphous materials soften over a range of temperatures. per, aluminum, and iron are formed by metal atoms Figure The ribed as a uniform distribution of atomic nuclei within a sea of delocalized ic solid are held together by a unique force known as metallic bonding that Network Figure To break or to melt a covalent network solid, covalent bo k properties. All exhibit high thermaltypes and ofelectrical conductivity, metallic Crystalline Solids and Their Properties Covalent org/content/col11760/1.9 are relatively strong, covalent network solids are typically characteriz Type of Properties Examples points. For example, diamond is one of the hardest substances known an Attractions Type of Solid Type of Particles ionic ions ionic bonds hard, brittle, conducts electricity as a liquid but not as a solid, high to very high melting points NaCl, Al2O3 metallic atoms of electropositive elements metallic bonds shiny, malleable, ductile, conducts heat and electricity well, variable hardness and melting temperature Cu, Fe, Ti, Pb, U covalent bonds very hard, not conductive, very high melting points covalent network atoms of electronegative C (diamond),

13 TYPES OF CRYSTALLINE SOLIDS & THEIR PROPERTIES Type of Solid Type of Particles Type of Attractions Properties Examples ionic ions ionic bonds hard, brittle, conducts electricity as a liquid but not as a solid, high to very high melting points NaCl, Al 2 O 3 metallic atoms of electropositive elements metallic bonds shiny, malleable, ductile, conducts heat and electricity well, variable hardness and melting temperature Cu, Fe, Ti, Pb, U covalent network atoms of electronegative elements covalent bonds very hard, not conductive, very high melting points C (diamond), SiO 2, SiC molecular molecules (or atoms) IMFs variable hardness, variable brittleness, not conductive, low melting points H 2 O, CO 2, I 2, C 12 H 22 O 11

14 Vacancies - gaps in the network, lacking an atom/molecule. Interstitial - atoms within the spaces of the regular crystalline structure. CRYSTAL DEFECTS Substitution - different atoms/ molecules replacing the regular ones in the crystalline structure. Some defects are made intentionally to give rise to unique properties in the materials (i.e. brass, carbon steel, semiconductors)

15 LATTICE STRUCTURES OF CRYSTALLINE SOLIDS Lecture Objectives Chapter 10.6 Describe the arrangement of atoms and ions in crystalline structures. Compute ionic radii using unit cell dimensions.

16 THE CRYSTAL LATTICE AND THE UNIT CELL A crystal lattice is a regular repeating pattern in a structure. Lattice points all have identical surroundings. Unit cell is the smallest portion of the crystal, that can be repeated on all faces to grow the crystal.

17

18 SIMPLE CUBIC UNIT CELL Coordination number = 6

19 BODY-CENTERED CUBIC UNIT CELL Coordination number = 8

20 FACE-CENTERED CUBIC UNIT CELL Coordination number = 12

21 QUESTION The crystal structure of iron contains two atoms per unit cell. What type of structure does it have? Simple cubic Face-centered cubic Body-centered cubic None of the above

22 PACKING OF IDENTICAL SPHERES Simple cubic (52% packing efficiency) Body-centered cubic (68% packing efficiency)

23 Hexagonal closest packing Layer a Layer b Cubic closest packing abab...(74%) abcabc...(74%) Layer c Hexagonal unit cell Expanded side views Face-centered unit cell

24 CRYSTAL STRUCTURES OF MOLECULES Cubic closest packing for frozen argon. Cubic closest packing for frozen methane.

25 QUESTION Lead crystallizes in the face-centered cubic lattice. What is the coordination number for Pb? Coordination number = 12

26 PROBLEM Determine the atomic radius and density of polonium. It crystallizes in a simple cubic lattice with a unit cell edge length of 336 pm. Calculate the atomic radius of polonium. Calculate the density of polonium in g/cm 3.

27 PROBLEM Iridium has the highest density of any element. It crystallizes in the facecentered cubic lattice with a unit cell edge length of pm. Calculate the atomic radius of iridium. Calculate the density of an iridium crystal in g/cm 3

28 CRYSTALS OF IONIC SOLIDS To maximize the interactions between ions of opposing charges, ionic solids attempt to surround each ion type with as many opposing charges as possible. This typically results in the smaller cations fitting into the spaces between the larger anions. The unit cell for an ionic solid consists of the smallest portion of the crystal that maintains the composition of the ionic compound (ratio equivalent to the empirical formula). Ionic crystals take on various forms depending on the ratios of the ions (charges) involved in the solid.

29 THE SODIUM CHLORIDE STRUCTURE The blending of two face-centered cubic lattices. Unit cell contains 4 Cl - ions and 4 Na + ions.

30 THE FLUORITE STRUCTURE Characteristic of calcium fluoride. Results in a face-centered cubic lattice. Unit cell contains 8 F - ions and 4 Ca 2+ ions.

31 THE ZINC BLEND Characteristic of zinc sulfide. Results in a face-centered cubic lattice. Unit cell contains 4 S 2- ions and 4 Zn 2+ ions.