1 / 73

Syllabus Chemistry 102 Spring 2011 Sec. 501, 503 (MWF 8-8:50, 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone O

Syllabus Chemistry 102 Spring 2011 Sec. 501, 503 (MWF 8-8:50, 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone Office: Room 123E Heldenfels (HELD) Telephone: 845-3010 (no voice mail) or leave message at 845-2356 email: stone@mail.chem.tamu.edu

liam
Télécharger la présentation

Syllabus Chemistry 102 Spring 2011 Sec. 501, 503 (MWF 8-8:50, 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone O

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. Syllabus Chemistry 102 Spring 2011 Sec. 501, 503 (MWF 8-8:50, 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone Office: Room 123E Heldenfels (HELD) Telephone: 845-3010 (no voice mail) or leave message at 845-2356 email: stone@mail.chem.tamu.edu (put CHEM 101-Sec. # + subject in subject line of your email) Office Hours: HELD 408: Tue. And Thurs. 8:00-10:50 AM I.A. Esther Ocola S.I. Leader: Mary Hesse

  2. Kotz and Treichel 7th ed. TEXTBOOKS Averill and Eldridge • Hardbound ~$200 • Solution Manual ~$40 • Online Tutor ~$45 • Total ~$285 • Ebook $45 per semester • Includes • Text • Solution manual • Online tutorial Helpful Online Dictionary of Chemistry Useful As A Second Language General Chemistry I and Organic Chemistry I (There are O-chem II and Physics books in this series if you find these useful and will have to take those classes. Yvette Freeman Publisher's Representative Pearson Education Yvette.Freeman@Pearson.com Chang’s Essentials

  3. http://slc.tamu.edu/ • Tutoring • Supplemental Instruction • Courses • Texas Success Initiative • About UsContact Us • 118 Hotard North of Sbisa, between • Neeley Hall and the Northside Post Office • (979) 845-2724 • The Student Learning Center has won the 2008 National College Learning Center Association • Frank L. Christ Outstanding Learning Center Award! • The award recognizes the center's commitment to supporting and strengthening the • Academic experience of students at Texas A&M University by providing a variety of programs and • services that promote retention and success. Read more... •   The Student Learning Center provides Supplemental Instruction and tutoring free of charge to all • Texas A&M University students. The SLC oversees the STLC courses (formerly CAEN), which teach • students how to improve their study skills and prepare for the job market. The SLC manages • Developmental programs for students who have not yet passed the assessment tests required by the state. • Study Tips • General • Time Management • Reading Textbooks • Setting Goals • Preparing for Exams • Success Tips from Fellow Aggies • Tutoring • During the Fall 2008 semester, drop-in tutoring will be offered Sunday nights 5-8pm and Monday through Thursday nights from 5-10pm. Tutoring will begin on Monday, September 1st. Tutor Zones are currently planned for Studio 12 of The Commons. Look for our table and tutors and just ask for help! See the schedule... Drop-in tutoring is available for most lower level math and science courses on the first floor of Hotard Hall. Tutors are also available to help out with many other courses. If you need help in a particular course and would like to check to see if a tutor is available for that course, you can contact our tutor coordinator, Linda Callen, at 845-2724.  

  4. How grades are determined The way the real world works Individual Mastery compared to a large population What you are used to and I will report 1) Raw scores are determined. Sum of points assigned to correct responses 2) Individual scores are normalized. A context-free evaluation of relative performance 3) Normalized scores are transformed. An absolute score is assigned to a defined scale 4) Letter grades are assigned>89.501 A >79.501 B >69.501 C >59.501 D <59.501 F

  5. Grading: • Your grade will be based on • Four one-hour examinations (200 raw points – T score 100+) • A final examination (400 raw points – 2 x T score 100+)

  6. The mere formulation of a problem is far more often essential than its solution, which may be merely a matter of mathematical or experimental skill. To raise new questions, new possibilities, to regard old problems from a new angle requires creative imagination and marks real advances in science. • ~Albert Einstein Problem - A situation that presents difficulty, uncertainty, or perplexity: Question - A request for data: inquiry, interrogation, query. Answer - A spoken or written reply, as to a question. Solution - Something worked out to explain, resolve, or provide a method for dealing with and settling a problem.

  7. c (ms-1) l (m) • Numbers – Significant Figures, Rounding Rules, Accuracy, Precision, Statistical Treatment of the Data • Units – 5 of the 7 • Time – seconds • Length – Meters Density? • Mass – grams Molecular Weight (Mass) • Amount – Moles Mole Ratio, Molarity, molality • Temperature – Kelvins • Vocabulary – Approximately 100 new terms or words and applying new or more rigid definitions to words you may already own. • Principles (Theories and Laws) – Stoichiometry, Quantum Theory, Bonding, Chemical Periodicity, Solutions, Thermodynamics, Intermolecular Forces, Gas Laws, Collogative Properties, Kinetics, Equilibrium, Electrochemistry • cp = q/mDT rate = k[A]m[B]n ∆E = q + w • DG = DH – TDS Eocell = Ecathode = Eanode • PV = nRT %yield = actual/theoretical * 100% K = • DT = Kmi [C]c[D]d [A]a[D]b E = n =

  8. Inter-molecular Forces Last Semester you studied INTRAmolecular forces—the forces holding atoms together to form molecules. Now turn to forces between molecules —INTERmolecular forces. Forces between molecules, between ions, or between molecules and ions.

  9. Ion-Ion Forcesfor comparison of magnitude Na+—Cl- in salt These are the strongest forces. Lead to solids with high melting temperatures. NaCl, mp = 800 oC MgO, mp = 2800 oC

  10. •• water dipole d - O •• H d + H Ion – Permanent DipoleAttractions Water is highly polar and can interact with positive ions to give hydrated ions in water.

  11. Attraction Between Ions and Permanent Dipoles Many metal ions are hydrated. This is the reason metal salts dissolve in water.

  12. Attraction Between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion charge and ion-dipole distance. Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+ -1922 kJ/mol -405 kJ/mol -263 kJ/mol

  13. Dipole-Dipole Forces Such forces bind molecules having permanent dipoles to one another. • Influence of dipole-dipole forces is seen in the boiling points of simple molecules. • Compd Mol. Wt. Boil Point • N2 28 -196 oC • CO 28 -192 oC • Br2 160 59 oC • ICl 162 97 oC

  14. Interactions in Liquid Solutions • Hydrophilic and hydrophobic solutes – A solute can be classified as hydrophilic, meaning that there is an electrostatic attraction to water, or hydrophobic, meaning that it repels water. 1. Hydrophilic substance is polar and contains O–H or N–H groups that can form hydrogen bonds to water; tend to be very soluble in water and other strongly polar solvents 2. Hydrophobic substance may be polar but usually contains C–H bonds that do not interact favorably with water; essentially insoluble in water and soluble in nonpolarsolvents – The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems. – Vitamins can be classified as either fat soluble or water soluble. • Fat-soluble vitamins (Vitamin A) are nonpolar, hydrophobic molecules and tend to be absorbed into fatty tissues and stored there. • Water-soluble vitamins (Vitamin C) are polar, hydrophilic molecules that circulate in the blood and intracellular fluids and are excreted from the body and must be replenished in the daily diet. – Ionic substances are most stable in polar solvents. – Water is the most common solvent for ionic compounds because of its high polarity. – A more useful measure of the ability of a solvent to dissolve ionic compounds is its dielectric constant (), which is the ability of a bulk substance to decrease the electrostatic forces between two charged particles.

  15. Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. • H-bonding is especially strong in water because • the O—H bond is very polar • there are 2 lone pairs on the O atom • Accounts for many of water’s unique properties. H-bonding is strongest when X and Y are N, O, or F

  16. Hydrogen Bonding in H2O Ice has open lattice-like structure. Ice density is < liquid and so solid floats on water. One of the VERY few substances where solid is LESS DENSE than the liquid. H bonds  abnormally high specific heat capacity of water (4.184 J/g•K) This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

  17. Boiling Points of Simple Hydrogen-Containing Compounds Active Figure 13.8

  18. Double helix of DNA Portion of a DNA chain H-bonding is especially strong in biological systems — such as DNA. DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O. Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases. —adenine with thymine —guanine with cytosine

  19. Base-Pairing through H-Bonds

  20. Base-Pairing through H-Bonds

  21. FORCES INVOLVING INDUCED DIPOLES How can non-polar molecules such as O2 and I2 dissolve in water? The water dipole INDUCES a dipole in the O2 electric cloud. Dipole-induced dipole

  22. - d I-I I-I + d The alcohol temporarily creates or INDUCES a dipole in I2. - - d d O O R H R H + + d d FORCES INVOLVING INDUCED DIPOLES Consider I2 dissolving in ethanol, CH3CH2OH. Solubility increases with mass the gas • Process of inducing a dipole is polarization • Degree to which electron cloud of an atom or molecule can be distorted in its polarizability. • Non-polar gases are most soluble in non-polar solvents.

  23. FORCES INVOLVING INDUCED DIPOLES Formation of a dipole in two nonpolar I2 molecules. The induced forces between I2 molecules are very weak, so solid I2 sublimes (goes from a solid to gaseous molecules). Induced dipole-induced dipole

  24. FORCES INVOLVING INDUCED DIPOLES The magnitude of the induced dipole depends on the tendency to be distorted. Higher molec. weight → larger induced dipoles. Molecule Boiling Point (oC) CH4 (methane) - 161.5 C2H6 (ethane) - 88.6 C3H8 (propane) - 42.1 C4H10 (butane) - 0.5

  25. Intermolecular Forces Summary – London dispersion forces, dipole-dipole interactions, and hydrogen bonds that hold molecules to other molecules are weak. – Energy is required to disrupt these interactions, and unless some of that energy is recovered in the formation of new, favorable solute-solvent interactions, the increase in entropy on solution formation is not enough for a solution to form.

  26. Vapor Pressure • When the liquid is heated, its molecules obtain sufficient kinetic energy to overcome the forces holding them in the liquid and they escape into the gaseous phase. • The result of this phenomenon is that the molecules from the liquid phase generate a population of molecules in the vapor phase above the liquid that produces a pressure called the vapor pressure of the liquid. • Plotting the fraction of molecules with a given KE against their KE gives the KE distribution of the molecules in the liquid. • Increasing the temperature increases both the average KE of the particles in a liquid and the range of KE for the molecules. • Molecules with KE greater than Eo can escape from the liquid to enter the vapor phase • Molecules must also be at the surface where it is physically possible for the molecule to leave the liquid surface.

  27. Equilibrium Vapor Phase • Volatile liquids have high vapor pressures and tend to evaporate readily; nonvolatile liquids have low vapor pressures and evaporate more slowly. • Substances with vapor pressures higher than that of water are volatile, whereas those with vapor pressures lower than water are nonvolatile. • Equilibrium vapor pressure of a substance at a particular temperature is a characteristic of the material. – Does not depend on the amount of liquid – Depends strongly on the temperature and intermolecular forces

  28. ether alcohol water dipole- dipole H-bonds H-bonds extensive increasing strength of IM interactions Liquids Equilibrium Vapor Pressure VP as a function of T. 1. The curves show all conditions of P and T where LIQ and VAP are in EQUILIBRIUM 2. The VP rises with T. 3. When VP = external P, the liquid boils. This means that BP’s of liquids change with altitude. 4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT 5. VP of a given molecule at a given T depends on IM forces. Here the VP’s are in the order O O O H5C2 C2H5 H5C2 H H H

  29. Liquids HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid. LIQ + heat VAP Compd. ∆Hvap (kJ/mol) IM Force H2O 40.7 (100 oC) H-bonds SO2 26.8 (-47 oC) dipole Xe 12.6 (-107 oC) induced dipole

  30. Phase Diagrams Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line. (At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

  31. Phase Equilibria — Water Solid-liquid Gas-Liquid Gas-Solid

  32. Triple Point — Water At the TRIPLE POINTall three phases are in equilibrium.

  33. Phases Diagrams—Important Points for Water T(˚C) P(mmHg) Normal boil point 100 760 Normal freeze point 0 760 Triple point 0.0098 4.58

  34. Critical T and P As P and T increase, you finally reach the CRITICAL T and P Above critical T no liquid exists no matter how high the pressure. The gas and liquid phases are intermixed and indistinguishable. At P < 4.58 mmHg and T < 0.0098 ˚C solid H2O can go directly to vapor. This process is called SUBLIMATION This is how a frost-free refrigerator works.

  35. Terminology and Supercritical Fluids • Critical temperature (Tc) – Temperature above which the gas can no longer be liquefied, regardless of pressure – The highest temperature at which a substance can exist as a liquid – Above the critical temperature, the molecules have too much kinetic energy for the intermolecular attractive forces to hold them together in a separate liquid phase • Critical pressure (Pc) – Minimum pressure needed to liquefy a substance at the critical temperature • Critical point — combination of critical temperature and critical pressure • As the temperature of a liquid increases, its density decreases. • As the pressure of a gas increases, its density increases. • At the critical point, the liquid and gas phases have exactly the same density, and only a single phase exists, called a supercritical fluid, which exhibits many of the properties of a gas but has a density typical of a liquid.

  36. P Liquid H O 2 Solid Normal H O 2 freezing point 760 mmHg 0 °C T Solid-Liquid Equilibria Raising the pressure at constant T causes water to melt. The NEGATIVE SLOPE of the S/L line is unique to H2O. Almost everything else has positive slope. The behavior of water under pressure is an example of LE CHATELIER’S PRINCIPLE At Solid/Liquid equilibrium, raising P squeezes the solid. It responds by going to phase with greater density, i.e., the liquid phase. In any system, if you increase P the DENSITY will go up. Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram).

  37. Changes of State • Changes of state are examples of phase changes, or phase transitions • Any of the three forms of matter (gas, liquid, solid) is converted to either of the other two • Six most common phase changes: 1. melting solid → liquid 2. freezing liquid → solid 3. vaporization liquid → gas 4. condensation gas → liquid 5. sublimation solid → gas 6. deposition gas → solid

  38. Temperature Curves Heating Curve – A plot of temperature versus time where heat is added • Sample is ice at -23ºC • As heat is added, the temperature of the ice increases linearly with time • Slope of the line depends on both the mass of the ice and the specific heat of ice, the number of joules required to raise the temperature of 1 g of ice by 1ºC • As the temperature of the ice increases, the water molecules in the ice crystal absorb more and more energy and at the melting point, they have enough kinetic energy to overcome attractive forces and to move • As more heat is added, the temperature of the system does not increase further, but remains constant at 0ºC until all the ice has melted • Once all the ice has been converted to liquid water, the temperature of the water begins to increase – temperature increases more slowly than before because the heat capacity of water is greater than that of ice • At 100ºC, water begins to boil; temperature remains constant until all the water has been converted to steam • Temperature again begins to rise, but at a faster rate because the heat capacity of steam is less than that of ice or water

  39. Temperature Curves • Cooling curves – A plot of temperature vs. time when heat is removed 1. As heat is removed from steam at 200ºC, the temperature falls until it reaches 100ºC, where the steam begins to condense to liquid water 2. No further temperature change occurs until all the steam is converted to the liquid; then the temperature again decreases as the water is cooled 3. Temperature drops below the freezing point for some time — region corresponds to an unstable form of the liquid, a supercooled liquid that will convert to a solid 4. As water freezes, temperature increases due to heat evolved and then holds constant at the melting point

  40. Units of Concentration • There are several different ways to quantitatively describe the concentration of a solution, which is the amount of solute in a given quantity of solution. 1. Molarity– Useful way to describe solution concentrations for reactions that are carried out in solution or for titrations. Volume of a solution depends on its density, which is a function of temperature Molarity = moles of solute = mol/L = mmol/mL liter of solution 2. Molality– Concentration of a solution can also be described by its molality (m), the number of moles of solute per kilogram of solvent. Depends on the masses of the solute and solvent, which are independent of temperature. Molality = moles of solute kilogram solvent • Mole fraction – Used to describe gas concentrations and determine the vapor pressures of mixtures of similar liquids. Depends on only the masses of the solute and solvent and is temperature independent. Mole fractions sum to one for a given mixture. Mole fraction () = moles of component . total moles in the solution 4. Mass percentage (%) – The ratio of the mass of the solute to the total mass of the solution. Result can be expressed as mass percentage, parts per million (ppm), or parts per billion (ppb). ppm and ppb are used for highly dilute solutions, and correspond to milligrams (10-3) and micrograms (10-6) of solute per kilogram of solution, respectively mass percentage = mass of solute 100% mass of solution parts per million (ppm)= mass of solute 106 mass of solution parts per billion (ppb)= mass of solute 109 mass of solution

  41. Units of Concentration – Mass percentage and parts per million or billion can express the concentrations of substances even if their molecular mass is unknown because these are simply different ways of expressing the ratios of the mass of a solute to the mass of the solution

  42. Molality and Mole Fraction • Calculate the molarity and the molality of an aqueous solution that is 10.0% glucose, C6H12O6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C6H12O6 = 180 g

  43. Molality and Mole Fraction • Calculate the molality of a solution that contains 7.25 g of benzoic acid C6H5COOH, in 2.00 x 102 mL of benzene, C6H6. The density of benzene is 0.879 g/mL. 1 mol C6H5COOH = 122 g

  44. Molality and Mole Fraction • What are the mole fractions of glucose and water in a 10.0% glucose solution?

  45. Vapor Pressure of Solutions andRaoult’s Law • The addition of a nonvolatile solute, one whose vapor pressure is too low to measure readily, to a volatile solvent decreases the vapor pressure of the solvent. • The relationship between solution composition and vapor is known as Raoult’s law, PA = AP0A, where PA is the vapor pressure of component A of the solution (the solvent), A is the mole fraction of A in solution, and P0A is the vapor pressure of pure A. • This equation is used to calculate the actual vapor pressure above a solution of a nonvolatile solute.

  46. Vapor Pressure of Solutions andRaoult’s Law • Plots of the vapor pressures of both components versus the mole fractions are straight lines that pass through the origin. • A plot of the total vapor pressure of the solution versus the mole fraction is a straight line that represents the sum of the vapor pressures of the pure components. • The vapor pressure of the solution is always greater than the vapor pressure of either component. • Solutions that obey Raoult’s law are called ideal solutions, in which the intermolecular forces in the two pure liquids are almost identical in both kind and magnitude and the change in enthalpy on solution formation is essentially zero (Hsoln 0). • Real solutions exhibit positive or negative deviations from Raoult’s law because the intermolecular interactions between the two components A and B differ. • Negative deviation • If the A–B interactions are stronger than the A–A and B–B interactions, each component of the solution exhibits a lower vapor pressure than expected for an ideal solution, as does the solution as a whole. • The favorable A–B interactions stabilize the solution compared with the vapor. • Positive deviation • If the A–B interactions are weaker than the A–A and B–B interactions yet the entropy increase is enough to allow the solution to form, both A and B have an increased tendency to escape from the solution into the vapor phase. • The result is a higher vapor pressure than expected for an ideal solution.

More Related