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Syllabus Chemistry 101 Fall 2006 Sec. 528-538 (MWF 12:40-1:30) RM 100 HELD

Syllabus Chemistry 101 Fall 2006 Sec. 528-538 (MWF 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone Office: Room 408 Heldenfels (HELD) Telephone: 845-3010 (no voice mail) or leave a message at 845-2356 email: stone@mail.chem.tamu.edu

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Syllabus Chemistry 101 Fall 2006 Sec. 528-538 (MWF 12:40-1:30) RM 100 HELD

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  1. Syllabus Chemistry 101 Fall 2006 Sec. 528-538 (MWF 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone Office: Room 408 Heldenfels (HELD) Telephone: 845-3010 (no voice mail) or leave a 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: Wed. 8:00-10:50 AM I.A. TBA S.I. Leader: TBA CHEM 101 and 102 are the first-year chemistry sequence in the core curriculum. These are 4-credit courses. The sections in this lecture are a part of a much larger program. Each grouping of sections of this course is independent of the other instructors’ sections, but we strive to cover common content. The lecture component of Chemistry 101 covers stoichiometry, atomic and molecular structure, organic nomenclature and structure, chemical bonding, acid/base chemistry, solution chemistry, properties of liquids and solids, and the gas laws.

  2. Required Course Materials: “Chemistry and Chemical Reactivity, 6th Edition”, by Kotz, Treichel, and Weaver Nonprogrammable calculator suitable to use on lecture exams. The lab notebook, 8 1/2” x 11”, perforated, numbered, alternating white and yellow pages. Approved eye protection. Prior to the first laboratory meeting, purchase four standard (8½ x11) gray scantron sheets (Form No. 0-101607-TAMU) from the bookstore and turn them in UNMARKED to your laboratory TA.

  3. Grading: • Your grade will be based on • Three one-hour examinations (each worth 180 points) • A final examination (360 points) • The laboratory component (300 points)  • Major Examination Schedule Fall 2005: • Mon. Sept. 25 Major Exam No.1 • Mon. Oct. 23 Major Exam No.2 • Mon. Nov. 20 Major Exam No.3 • Mon. Dec. 5 Make-up for University approved absences • Mon. Dec 11 Final Exam 10:30 to 12:30

  4. 80% 63% 70% 90% 46% 80% 60% 100% 97% 29% D,F,Q,W C B A 2% 16% 50% 84% 98% Percentile Rank A = > average + 1 s B = > average C = > average - 1 s

  5. Matter and Energy - Vocabulary • Chemistry • Science that describes matter – its properties, the changes it undergoes, and the energy changes that accompany those processes • Matter • Anything that has mass and occupies space. • Energy • The capacity to do work or transfer heat. • Scientific (natural) law • A general statement based the observed behavior of matter to which no exceptions are known.

  6. Scientific Method 1 • Observation • Hypothesis • Experiment • Theory • Law 2 3 4 5

  7. Natural Laws • Law of Conservation of Mass • Law of Conservation of Energy • Law of Conservation of Mass-Energy • Einstein’s Relativity • E=mc2

  8. Dr. Stone’s patent pending chemistry problem solver • Write down everything you are given including units • Vocabulary • Write down what you want to know including the units • Write down mathematical equation(s) that includes these • values and units • Write a balanced stoichiometric equation • Mole concept • Convert everything to moles • Dimensional analysis • Convert everything to the unknown’s units • Rounding, significant figures, accuracy and precision

  9. Chemical and Physical Properties • Chemical Properties - chemical changes • rusting or oxidation • chemical reactions • Physical Properties - physical changes • changes of state • density, color, solubility • Extensive Properties - depend on quantity • Intensive Properties - do not depend on quantity Units of Measurement Use SI units — based on the metric system • Mass - measure of the quantity of matter in a body - grams • Weight - measure of the gravitational attraction for a body • Length - meters • Time - seconds • Temperature - oC, K

  10. Use of Numbers • Exact numbers • 1 dozen = 12 things for example • Accuracy • how closely measured values agree with the correct value • Precision • how closely individual measurements agree with each other • Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

  11. Use of Numbers • Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends.

  12. States of Matter • Change States • heating • cooling

  13. States of Matter • Illustration of changes in state • requires energy

  14. Compounds & Molecules • COMPOUNDS are a combination of 2 or more elements in definite ratios by mass. • The character of each element is lost when forming a compound. • MOLECULES are the smallest unit of a compound that retains the characteristics of the compound. Composition of molecules is given by a MOLECULAR FORMULA C8H10N4O2 - caffeine H2O

  15. MOLECULAR FORMULAS • Formula for glycine is C2H5NO2 • In one molecule there are • 2 C atoms • 5 H atoms • 1 N atom • 2 O atoms

  16. WRITING FORMULAS • Can also write glycine formula as • H2NCH2COOH to show atom ordering • or in the form of a structural formula

  17. MOLECULAR MODELING Structural formula of glycine Ball & stick Space-filling

  18. Counting Atoms Chemistry is a quantitative science—we need a “counting unit.” MOLE - 1 mole is the amount of substance that contains as many particles (atoms, molecules) as there are in 12.0 g of 12C. Avogadro’s Number 6.02214199 x 1023 There is Avogadro’s number of particles in a mole of any substance. 518 g of Pb, 2.50 mol Amedeo Avogadro 1776-1856

  19. Molar Mass 1 mol of 12C = 12.00 g of C = 6.022 x 1023 atoms of C 12.00 g of 12C is its MOLAR MASS Taking into account all of the isotopes of C, the molar mass of C is 12.011 g/mol

  20. One-mole Amounts

  21. Atomic Weights • Weighted average of the masses of the constituent isotopes if an element. • Tells us the atomic masses of every known element. • Lower number on periodic chart. • How do we know what the values of these numbers are?

  22. PROBLEM: What amount of Mg is represented by 0.200 g? How many atoms? Mg has a molar mass of 24.3050 g/mol. How many atoms in this piece of Mg?

  23. MOLECULAR WEIGHT & MOLAR MASS Molecular weight = sum of the atomic weights of all atoms in the molecule. Molar mass = molecular weight in grams Problem: What is the molar mass of ethanol – C2H6O? 1 mol contains 2 mol C (12.01 g C/1 mol) = 6 mol H (1.01 g H/1 mol) = 1 mol O (16.00 g O/1 mol) = TOTAL = molar mass =

  24. Tylenol • Formula = • Molar mass = C8H9NO2 151.2 g/mol

  25. Molar Mass

  26. PROBLEM: How many moles of alcohol are there in a “standard” can of beer if there are 21.3 g of C2H6O? (a) Molar mass of C2H6O = 46.08 g/mol (b) Calc. moles of alcohol

  27. PROBLEM: How many moleculesof alcohol are there in a “standard” can of beer if there are 21.3 g of C2H6O? We know there are 0.462 mol of C2H6O.

  28. PROBLEM: How many atoms of C are there in a “standard” can of beer if there are 21.3 g of C2H6O?

  29. Molecular & Ionic Compounds NaCl Heme N Fe

  30. IONS AND IONIC COMPOUNDS • IONS are atoms or groups of atoms with a positive or negative charge. • Taking away an electron from an atom gives a CATION with a positive charge • Adding an electron to an atom gives an ANION with a negative charge. A CATION forms when an atom loses one or more electrons. An ANION forms when an atom gains one or more electrons Mg --> Mg2+ + 2 e- F + e- --> F- In general • metals (Mg) lose electrons ---> cations • nonmetals (F) gain electrons ---> anions

  31. Charges on Common Ions -1 -3 -2 -4 +1 +2 +3 By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

  32. POLYATOMIC IONS Groups of atoms with a charge. MEMORIZE the names and formulas in Table 3.1, page 107. NO3- nitrate ion HNO3 nitric acid

  33. COMPOUNDS FORMED FROM IONS CATION + ANION ---> COMPOUND Na+ + Cl- --> NaCl A neutral compd. requires equal number of + and - charges.

  34. Some Ionic Compounds Mg2+ + NO3- ----> Mg(NO3)2 magnesium nitrate Fe2+ + PO43- ----> Fe3(PO4)2 iron(II) phosphate (See CD, Screen 3.11 for naming practice) Ca2+ + 2 F- ---> CaF2 calcium fluoride

  35. Properties of Ionic CompoundsForming NaCl from Na and Cl2 • A metal atom can transfer an electron to a nonmetal. • The resulting cation and anion are attracted to each other by electrostatic forces. These forces are governed by COULOMB’S LAW.

  36. Molecular CompoundsCompounds without Ions CO2 Carbon dioxide BCl3 boron trichloride CH4 methane

  37. Naming Molecular Compounds All are formed from two or more nonmetals. CO2 Carbon dioxide Ionic compounds generally involve a metal and nonmetal (NaCl) BCl3boron trichloride CH4 methane

  38. Empirical & Molecular Formulas A pure compound always consists of the same elements combined in the same proportions by weight. Therefore, we can express molecular composition as PERCENT BY WEIGHT Ethanol, C2H6O 52.13% C 13.15% H 34.72% O

  39. Percent Composition Consider NO2, Molar mass = ? What are the weight percentages of N and O in NO?

  40. Determining Formulas In chemical analysis we determine the % by weight of each element in a given amount of pure compound and derive the EMPIRICAL or SIMPLEST formula. PROBLEM: A compound of B and H is 81.10% B. What is its empirical formula?

  41. DETERMINE THE FORMULA OF A COMPOUND OF Sn AND I Sn(s) + some I2(s) ---> SnIx

  42. Data to Determine the formula of a Sn—I Compound • Reaction of Sn and I2 is done using excess Sn. • Mass of Sn in the beginning = 1.056 g • Mass of iodine (I2) used = 1.947 g • Mass of Sn remaining = 0.601 g

  43. Tin and Iodine Compound Now find the number of moles of I2 that combined with 3.83 x 10-3 mol Sn. Mass of I2 used was 1.947 g. How many mol of iodine atoms? Now find the ratio of number of moles of moles of I and Sn that combined.

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