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Dr. C. Weldon Mathews Chem 122 Office: 0042 Evans Lab Telephone: 292-1574

Dr. C. Weldon Mathews Chem 122 Office: 0042 Evans Lab Telephone: 292-1574 email: mathews.6@osu.edu course web site: http://www.chemistry.ohio-state.edu/~mathews/chem122wi07/ Office hours: TR 12:30 - 2:00 pm TR 4:00 - 5:00 pm

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Dr. C. Weldon Mathews Chem 122 Office: 0042 Evans Lab Telephone: 292-1574

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  1. Dr. C. Weldon Mathews Chem 122 Office: 0042 Evans Lab Telephone: 292-1574 email: mathews.6@osu.edu course web site: http://www.chemistry.ohio-state.edu/~mathews/chem122wi07/ Office hours: TR 12:30 - 2:00 pm TR 4:00 - 5:00 pm or by appointment Chapters we’ll cover in Chem 122: 10, 11, 13, 14, 15, 16, 17 (17.1-17.3) First Week: 10.0-10.6 Second Week: 10.7-10.9 and 11.1-11.5 First Quiz: Week of Jan 8 (second week) Review Chem 121, especially Chaps 8 and 9

  2. 1st Lecture & 1st Quiz

  3. Bloom’s Taxonomy– [Approaches to learning Chemistry.] Knowledge – Simple recall of facts Comprehension – Translate into your own words or equations. Application – Apply concepts to specific situations; recognizing and solving a problem when the equations are not given. Analysis – Application plus recognition of important parts of problem. Synthesis – Assemble components into a form new to them, i.e. design a research plan or devise a synthetic scheme. Evaluation – Judge the value of materials in terms of internal and external criteria. This is a grossly abbreviated adaptation from Bloom, B. S. (Ed.) (1956) Taxonomy of educational objectives: The classification of educational goals: Handbook I, cognitive domain. New York; Toronto: Longmans, Green. Use Google to find other references. see also http://www.coun.uvic.ca/learn/program/hndouts/bloom.html http://www.officeport.com/edu/blooms.htm http://www.kurwongbss.eq.edu.au/thinking/Bloom/blooms.htm Expected in this course. Expected in Research.

  4. Study Habits and Study Resources: • “Lectures” and “Reading” - minimal impact by themselves • “Chemistry is not a Spectator Sport!” Prof. Janet Tarino, OSU Mansfield • Recitation and Laboratory TAs • Ask questions and seek help whenever you need it! • EXPECTATIONS • Web resources: • http://www.chemistry.ohio-state.edu/ • http://www.chemistry.ohio-state.edu/~mathews/chem122wi07/ • /~rbartosz/ • /~rzellmer/ • chemistry ->Undergraduate Program->Interactive Tutorials

  5. First Lab Experiment: Calorimetry and Hess’s Law You’ll need to review material from Chem 121: Calorimetry, Chap 5, pp 182-187 Hess’s Law, Chap 5, pp 187-191 Enthalpies of Formation, pp 191-196 (and of Reactions)

  6. Some other things you need to know: Your Carmen Student ID. If you can’t find it your TA will be able to help in lab or recitation. Call Number for your section (eg 04499-3) and your Chemistry Dept. Section Number (eg 109), as shown on one of the following slides.

  7. Characteristics of Gases • Unlike liquids and solids, they • Expand to fill their containers. • Are highly compressible. • Have extremely low densities • Mix homogeneously with other gasses

  8. Characteristics of Gases

  9. Pressure • Pressure is the force acting on an object per unit area: • Gravity exerts a force on the earth’s atmosphere • A column of air 1 m2 in cross section exerts a force of 105 N, with a mass of about 104 kg or 2.2 x 104 lbs. • The pressure of a 1 m2 column of air is about 100 kPa.

  10. Pressure F = ma = (104 kg)(9.8m/s2) = 1 x 105 kg-m/s2 = 1 x 105 N = Newton P = 1 x 105 N/m2 = 1 x 105 Pa = 1 x 102 kPa

  11. Pressure • Atmosphere Pressure and the Barometer • SI Units: 1 N = 1 kg.m/s2; 1 Pa = 1 N/m2. • Atmospheric pressure is measured with a barometer. • Standard atmospheric pressure is the pressure required to support 760 mm of Hg in a column. • Units: 1 atm = 760 mmHg = 760 torr = 1.01325  105 Pa = 101.325 kPa.

  12. Pressure Atmosphere Pressure and the Barometer Notice that the top end of the BAROMETER is closed and that a “Torricelli VACUUM” exists above the mercury.

  13. Pressure • Atmosphere Pressure and the Barometer • The pressures of gases not open to the atmosphere are measured in manometers. • A MANOMETER consists of a bulb of gas attached to a U-tube containing Hg or other liquids: • If Pgas < Patm then Pgas + Ph2 = Patm. • If Pgas > Patm then Pgas = Patm + Ph2. Here Ph may be positive or negative!

  14. Pressure Notice again the various units that may be used to measure Pressure: Easiest to refer to the ‘Standard Atmospheric Pressure’ which is defined as 1 atm 1 atm = 101.325 kPa = 1.01325 x 105 Pa = = 760 mmHg = 760 torr Memorize these relations! They will be useful and help you with conversion factors between the various units.

  15. The Gas Laws • The Pressure-Volume Relationship: Boyle’s Law • Weather balloons are used as a practical consequence to the relationship between pressure and volume of a gas. • As the weather balloon gets further from the earth’s surface, the atmospheric pressure decreases. • As a consequence, the volume of the balloon increases. • Boyle’s Law: the volume of a fixed quantity of gas is inversely proportional to its pressure. • Boyle used a manometer to carry out the experiment.

  16. In our discussions about gases, we’ll often use the gas cylinder with a movable piston as a helpful analogy. You also may think of a bicycle pump as an example.

  17. Let’s run an “experiment” animation

  18. The Gas Laws The Pressure-Volume Relationship: Boyle’s Law V = f(P) = ? find VP = cnst = k V = f(1/P) = k(1/P)

  19. The Gas Laws • The Pressure-Volume Relationship: Boyle’s Law • Mathematically: • A plot of V versus P is a hyperbola. • Similarly, a plot of V versus 1/P must be a straight line passing through the origin.

  20. The Gas Laws • The Temperature-Volume Relationship: Charles’s Law • We know that hot air balloons expand when they are heated. • Charles’s Law: the volume of a fixed quantity of gas at constant pressure increases as the temperature increases. • Mathematically:

  21. These are typical of observations you might make in the lab. Notice that on this plot the equation is of the form y = a + b x and the volume does NOT go to 0 at x = 0 !!!

  22. The Gas Laws • The Temperature-Volume Relationship: Charles’s Law • A plot of V versus T is a straight line. • When T is measured in C, the intercept on the temperature axis is -273.15C. • We defineabsolute zero, 0 K = -273.15C. • Note the value of the constant reflects the assumptions: of a constant amount of gas and pressure. • And now the equation is of the form y = bx , i.e. a = 0.

  23. The Gas Laws • The Quantity-Volume Relationship: Avogadro’s Law • Gay-Lussac’s Law of combining volumes: at a given temperature and pressure, the volumes of gases which react are ratios of small whole numbers.

  24. The Gas Laws • The Quantity-Volume Relationship: Avogadro’s Law • Avogadro’s Hypothesis: equal volumes of gas at the same temperature and pressure will contain the same number of molecules. • Avogadro’s Law: the volume of gas at a given temperature and pressure is directly proportional to the number of moles of gas.

  25. The Gas Laws • The Quantity-Volume Relationship: Avogadro’s Law • Mathematically: • We will show that 22.4 L of any gas at 0 C and 1 atm contain 6.02  1023 gas molecules = 1 mole of molecules.

  26. The Gas Laws The Quantity-Volume Relationship: Avogadro’s Law

  27. Ideal Gas Equation • Consider the three gas laws. • We can combine these into a general gas law: • Boyle’s Law: • Charles’s Law: • Avogadro’s Law:

  28. Ideal Gas Equation • If the k is now defined as R, the proportionality constant of (called the gas constant), then • The ideal gas equation is: • R = 0.08206 L·atm/mol·K = 8.314 J/mol·K But how can you derive the value of R?

  29. Ideal Gas Equation Recall this slide and how to convert grams to moles. The Quantity-Volume Relationship: Avogadro’s Law

  30. Ideal Gas Equation • We define STP (standard temperature and pressure) = 0C, 273.15 K, 1 atm. • And now we see the volume of 1 mol of gas at STP is:

  31. ( notice the “>” should be a “ / “ )

  32. Fig. 10.12 shows the actual results for a number of gases. Notice that the “ideal gas model” does a pretty good job.

  33. Ideal Gas Equation • The Ideal-Gas Equation and the Gas Laws • If PV = nRT and n and T are constant, then PV = constant and we have Boyle’s law. • Other laws can be generated similarly. • In general, if we have a gas under two sets of conditions, then

  34. Applications of The Ideal Gas Equation • Gas Densities and Molar Mass • Density has units of mass over volume. • Rearranging the ideal-gas equation with M as molar mass we get has units of mol·L-1 And now the units are g·L-1

  35. Ideal Gas Equation • Gas Densities and Molar Mass • The molar mass of a gas can be determined as follows: • Volumes of Gases in Chemical Reactions • The ideal-gas equation relates P, V, and T to number of moles of gas. • The n can then be used in stoichiometric calculations.

  36. Ideal Gas Equation Volumes of Gases in Chemical Reactions Consider now the application of these ideas to chemical reactions. eg, 2 NaN3 (s)  2 Na + 3 N2 (g) Given the mass of sodium azide that reacts, the number of moles of nitrogen gas generated may be calculated. From this, the volume may be calculated at a given temperature and pressure.

  37. What volume of gas at 760 torr and 0 oC would be generated from 32.51 g of sodium azide which decomposes according to the equation 2 NaN3 (s)  2 Na + 3 N2 (g) ? In order to answer this question, we need to know how many moles of nitrogen will be generated (see Chem 121). Then we apply the ideal gas law.

  38. Gas Mixtures & Partial Pressures • Since gas molecules are so far apart, we can assume they behave independently. • Dalton’s Law: in a gas mixture the total pressure is given by the sum of partial pressures of each component: • Each gas obeys the ideal gas equation:

  39. Gas Mixtures & Partial Pressures • Combing the equations • Partial Pressures and Mole Fractions • Let ni be the number of moles of gas i exerting a partial pressure Pi, then • where i is the mole fraction (ni/nt).

  40. A mixture of gases containing 0.538 mol He(gas 1), 0.315 mol Ne (gas 2), • and 0.103 mol Ar (gas 3) is confined in a 7.00-L vessel at 25 0C. • Calculate the partial pressure of each gas. (b) Calculate the total • pressure in the vessel. (c) Calculate the mole fraction of each gas. P1 = n1 RT / V = (0.538 mol)(0.0821 L-atm/mol-K)(298 K) / (7.00 L) = 1.88 atm of He similarly, P2 = 1.10 atm of Ne, and P3 = 0.360 atm Ar The total pressure is just PT = P1 + P2 + P3 = ∑ Pi = 1.88 + 1.10 + 0.360 = 3.34 atm The mole fraction of He is X1 = P1 / PT = 1.88/3.34 = 0.563 (This also could be obtained from X1 = n1 / nT = 0.538/0.956 Likewise, X2 = 0.329 and X3 = 0.108 note that ∑ Xi = 1.00 always (within error limits): 0.563 + 0.329 + 0.108 = 1.00

  41. Gas Mixtures & Partial Pressures • Collecting Gases over Water • It is common to synthesize gases and collect them by displacing a volume of water. • To calculate the amount of gas produced, we need to correct for the partial pressure of the water:

  42. Gas Mixtures & Partial Pressures Collecting Gases over Water

  43. 10.7 Kinetic Molecular Theory • Theory developed to explain gas behavior. • Theory based on properties at the molecular level. • Assumptions: • Gases consist of a large number of molecules in constant random motion. • Volume of individual molecules negligible compared to volume of container. • Intermolecular forces (forces between gas molecules) negligible.

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