Course Objectives. Problems, Exercises Exams. Grading. Semester Schedule (later slide) Chem 1310 Summer 2003

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1 Subject: new element A major research institution has recently announced the discovery of the heaviest element yet known to science. This new element has been tentatively named "Administratium". Administratium has I neutron, 12 assistant neutrons, 75 deputy neutrons, and III assistant deputy neutrons, giving it an atomic mass of 312. These 312 particles are held together by a force called "morons", which are surrounded by vast quantities of lepton-like particles called "peons". Since Administratium has no electrons, it is inert. However, it can easily be detected as it impedes every reaction with which it comes into contact. A minute amount of Administratium causes one reaction to take over 4 days to complete when it would normally take less than an hour. Administratium has a normal half-life of 3 years; it does not decay but instead undergoes a reorganization, in which a portion of the assistant neutrons and deputy neutrons and assistant deputy neutrons exchange places. In fact, Administratium's mass will actually increase over time, since each reorganization causes some morons to become neutrons forming "isodopes." This characteristic of moron-promotion leads some scientists to speculate that Administratium is formed whenever morons reach a certain quantity in concentration. This hypothetical quantity is referred to as "Critical Mass". You will know it when you see it. OFB Chapter 1 1 Chem 1310 Summer 2003 William J. Baron, Ph.D. GaTech School of Chemistry and Biochemistry Director of the Freshman Chemistry Program 1310, 1311, 1312, 1313 Office Boggs Office Chem Annex Room 47 preferred bill.baron@chemistry.gatech.edu Phone Mailbox in Chemistry Office Lucent Technologies, Bell Laboratories (retired) Director, Optical Products Technology Norcross Fiber Optics, R&D, product design and development Director, Advanced Materials Development New Jersey Semiconductor crystals and devices Circuit board fabrication and assembly Ph.D. Princeton, Postdoctoral work at Ruhr Universität Bochum and Columbia OFB Chapter 1 2 Chem 1310 Summer 2003 Sections EF Course Objectives Semester Schedule (later slide) Problems, Exercises Exams Friday, June 6 Friday, June 27 Friday, July 18 Monday, April 14 Tuesday, July 29, 8:00 AM to 10:40 AM in Room CoC 16) Grading Hour Exams 40% (10%, 15%, 15%) In-class Quizzes 5% Final Exam 25% Homework 10% Laboratory 20% Laboratory Grade: You must pass Laboratory to pass the overall course. Teaching Assistants will have the responsibility for establishing laboratory grades. Students are graded on pre-lab quizzes, formal lab reports, summary reports, report accuracy, lab technique and attitude, Lab Midterm exam, and Lab Final exam. A grade of 70% or better in the lab is considered passing. Grades between 60% and 65% will be considered, if documentation is provided for any extenuating circumstances. At the beginning of the term, you will be asked to read and sign a pledge that laboratory reports and data gathering are your own, even when using partners. OFB Chapter 1 3 OFB Chapter 1 4

2 Chem 1310 Summer 2003 Sections EF Lectures MWF, 1:20 2:30 Will use ppt, suggest downloading lecture notes, print 1, 4, 6 per slide Recitations and Labs Recitations: (70 minutes) Either Tuesdays or Thursdays, 1:20-2:30 PM Labs: (3 hours 40 minutes) Tuesdays or Thursdays, 12:00-3:40 PM Office Hours MWF following lecture in Chem Annex 47 office Or by appointment Study ask questions, memorization not the whole answer, study everyday, stay limber, can t cram OFB Chapter 1 5 Instructor Teaching Assistants Required Course materials Course Description Syllabus and Schedule Textbook errors Exams Exam 1 Exam 2 Exam 3 Exam 4 Final Exam Homework Lecture Notes WebCT old exams answer key OFB Chapter 1 6 Reading OFB Chapter 1 Appendices A, B, and C Read, Read, Read Work Problems, Work Problems, Work Problems Chapter 1 The Atomic Nature of Matter 1-1 Chemistry: Science of Change 1-2 The Composition of Matter 1-3 The Atomic Theory of Matter 1-4 Chemical Formulas and Relative Atomic Masses 1-5 The Building Blocks of the Atom 1-6 Finding Atomic Masses the Modern Way 1-7 The Mole Concept: Counting and Weighing Atoms and Molecules 1-8 Finding Empirical and Molecular Formulas the Modern Way 1-9 Volume and Density OFB Chapter 1 7 OFB Chapter 1 8

3 Definitions Analysis (Take things apart) Synthesis (Put things together) Physical Properties (Color, odor, taste, boiling point, etc.) Chemical properties (with respect to other materials, e.g., uniformity) Substance (refer to elements and compounds, never mixtures) Elements (cannot be decomposed into simpler substances) Compounds (contain two or more elements bonded together, e.g., NaCl) Molecule (a few atoms connected together, e.g., CO 2 ) Mixtures (can be separated into two or more substances) Homogenous (uniform throughout, solutions) Heterogeneous (properties vary from region to region) Phase (liquid, gas, solid) Examples Table Salt Heterogeneous (mixture of NaCl and small amounts of other substances) Wood Heterogeneous (mixture of tree cells, and thousands of other substances) Mercury A substance and an element Air Homogenous mixture of gases Also a Heterogeneous mixture of gases and dust Water A substance that is a compound with molecular formula H 2 O What is an example of a homogenous sample that would gradually become heterogeneous if left to itself? E.g., A solution of sugar in water 1 st Homogeneous 2 nd allow evaporation to start, becomes heterogeneous 3 rd, complete evaporation, becomes homogeneous OFB Chapter 1 9 OFB Chapter 1 10 Atomic Theory of Matter Law of conservation of mass: Mass is neither created nor destroyed in a chemical reaction Dalton s Atomic Theory of Matter (1808): 1. All matter consists of solid and indivisible atoms 2. All atoms of a given chemical element are identical in mass and in all other properties 3. Different elements have different kinds of atoms; these atoms differ in mass from element to element 4. Atoms are indestructible and retain their identity in all chemical reactions 5. The formation of a compound from its elements occurs through the combination of atoms of unlike elements in small whole-number Chemical Formulas and Relative Atomic Masses Chemical Formulas display symbols for the elements and the relative number of atoms E.g., NH 3, CO 2, CH 3 CO 2 H or C 2 H 4 O 2 Molecules are groupings of two or more atoms bound closely together by strong forces that maintain them in a persistent combination ratio. OFB Chapter 1 11 OFB Chapter 1 12

4 Exercise 1-1 Problem: Every g of the compound SiH 4 contains g of Si and g of H. Find the ratio of the atomic mass of S to the atomic mass of H. Strategy: 1. Take the ratio of S/H 2. Account for the 1:4 ratio of S:H Exercise 1-1 Problem: Every g of the compound SiH 4 contains g of Si and g of H. Find the ratio of the atomic mass of S to the atomic mass of H OFB Chapter 1 13 OFB Chapter 1 14 Building Blocks of the Atom Electrons, Protons and Neutrons Electrons discovered in 1897 by Thomson Rutherford proposed that the atomic nucleus was composed of neutral particles called Neutrons and positively charged particles called protons Neutron number = N Atomic number = Z = number of Protons Atomic mass number = A A = Z + N OFB Chapter 1 15 OFB Chapter 1 16

5 C N Atomic Mass C Carbon Atomic Number Carbon Nitrogen A = Z + N Si Silicon P Phosphorus Atomic Mass = # Protons + # Neutrons For Carbon, 12 = 6 + Neutrons Neutrons = 6 Every Carbon atom has 6 electrons, 6 protons and 6 neutrons OFB Chapter 1 17 OFB Chapter 1 18 Mass Spectrometry and Isotopes Mass Spectrometer accelerates ions (or molecular ions) in an electric field and then separates those ions by relative mass in a magnetic field Relative Amount Mass Spectrometer Separation of Chlorine Relative Mass Cl Chlorine Atoms Avogadro s Number is the number of 12 C atoms in exactly 12 grams of carbon N 0 = X The mass, in grams, of Avogadro's number of atoms of an element is numerically equal to the relative atomic mass of that element OFB Chapter 1 19 OFB Chapter 1 20

6 Molecules Relative Molecular Mass of a molecule equals the sum of the relative atomic masses of all of the atoms making up the molecule Moles A mole measures the chemical amount of a substance Mole is an abbreviation of gram molecular weight One mole of a substance equals the amount that contains Avogadro's number of atoms, molecules. One mole = Molar mass (M) of that element or molecule OFB Chapter 1 21 OFB Chapter 1 22 Exercise 1-6 Molecules of isoamyl acetate have the formula C 7 H 14 O 2. Calculate (a) how many moles and (b) how many molecules are present in 0.250g of isoamyl acetate. Strategy: 1. Calculate molar mass of C 7 H 14 O 2 2. Calculate the number of moles in grams 3. Using Avogadro s number to calculate the number of molecules in X moles of C 7 H 14 O 2 Exercise 1-6 Molecules of isoamyl acetate have the formula C 7 H 14 O 2. Calculate (a) how many moles and (b) how many molecules are present in 0.250g of isoamyl acetate. 1. Calculate molar mass of C 7 H 14 O 2 2. Calculate the number of moles in grams 3. Using Avogadro s number calculate the number of molecules in X moles of C 7 H 14 O 2 OFB Chapter 1 23 OFB Chapter 1 24

7 Percentage Composition from Empirical or Molecular Formula Exercise 1-8 Tetrodotoxin, a potent poison found in the ovaries and liver of the globefish, has the empirical formula C 11 H 17 N 3 O 8. Calculate the mass percentages of the four element in this compound. Strategy: 1. Calculate molar mass of C 11 H 17 N 3 O, by finding the mass contributed by each element 2. Divide the mass for each element by the total mass of the compound. Exercise 1-8 Tetrodotoxin has the empirical formula C 11 H 17 N 3 O 8. Calculate the mass percentages of the four element in this compound. 1. Calculate molar mass of C 11 H 17 N 3 O, by finding the mass contributed by each element 2. Divide the mass for each element by the total mass of the compound. OFB Chapter 1 25 OFB Chapter 1 26 Finding an Empirical Formula Exercise 1-9 Heating a 150.0mg dose of a compound used to treat rheumatism decomposes it to its constituent elements, which are separated. There are mg of gold, mg of sodium, mg of oxygen, and mg of sulfur. Determine the empirical formula of this compound. Strategy: 1. Calculate the chemical amount (in moles) of each element in the sample using the table of atomic masses. 2. Find the ratios of the moles for each element by dividing each by the smallest one, i.e., normalize to the smallest. 3. If necessary, multiply smallest factor that clears any fractions that they contain. OFB Chapter 1 27 Exercise 1-9 Heating a 150.0mg dose of a compound used to treat rheumatism decomposes it to its constituent elements, whic are separated. There are mg of gold, mg of sodium, mg of oxygen, and mg of sulfur. Determine the empirical formula of this compound. 1. Calculate the chemical amount (in moles) of each element in the sample using the table of atomic masses. 2. Find the ratios of the moles for each element by dividing each by the smallest one, i.e., normalize the smallest. 3. If necessary, multiply smallest factor that clears any fractions that they contain. OFB Chapter 1 28

8 Exercise 1-9 Determine the empirical formula of this compound. 1. Calculate the chemical amount (in moles) of each element in the sample using the table of atomic masses. Exercise 1-10 Moderate Heating of mg of a compound containing nickel, carbon and oxygen and no other elements drives off all of the carbon and oxygen in the form of carbon monoxide (CO) and leaves mg of metallic nickel behind. Determine the empirical formula of the compound. 2. Find the ratios of the moles for each element by dividing each by the smallest one, i.e., normalize to the smallest. 3. If necessary, multiply smallest factor that clears any fractions that they contain. OFB Chapter 1 29 Strategy: 1. Write the reaction 2. Use the conservation of mass to find the amount of CO 3. Find the number of moles of CO and Nickel 4. Find the ratios of the moles for each substance by dividing each by the smallest one, i.e., normalize to the smallest. OFB Chapter 1 30 Exercise 1-10 Moderate Heating of mg of a compound containing nickel, carbon and oxygen and no other elements drives off all of the carbon and oxygen in the form of carbon monoxide (CO) and leaves mg of metallic nickel behind. Determine the empirical formula of the compound. 1. Write the reaction Exercise 1-10 Moderate Heating of mg of a compound containing nickel, carbon and oxygen and no other elements drives off all of the carbon and oxygen in the form of carbon monoxide (CO) and leaves mg of metallic nickel behind. Determine the empirical formula of the compound. 1. Write the reaction 2. Use the law of conservation of mass to find the amount of CO 2. Use the law of conservation of mass to find the amount of CO 3. Find the number of moles of CO and Nickel 3. Find the number of moles of CO and Nickel 4. Find the ratios of the moles for each substance by dividing each by the 4. Find the ratios of the moles for each substance by dividing each by the smallest one, i.e., normalize to the smallest. smallest one, OFB Chapter i.e., normalize 1 to the 31 OFB Chapter 1 32

9 Combustion Analysis Exercise 1-11 A sample of a liquid hydrocarbon weighing mg is burned in a combustion train to give mg of carbon dioxide, mg of water and no other products. What is the empirical formula of this hydrocarbon? Strategy: 1. Calculate the chemical amount (in moles) of carbon dioxide and water using the table of atomic masses. 2. Find the ratios of the moles for each substance by dividing each by the smallest one, i.e., normalize to the smallest. 3. If necessary, multiply smallest factor that clears any fractions that they contain. Exercise 1-11 A sample of a liquid hydrocarbon weighing mg is burned in a combustion train to give mg of carbon dioxide, mg of water and no other products. What is the empirical formula of this hydrocarbon? 1. Calculate the chemical amount (in moles) of carbon dioxide and water using the table of atomic masses. 2. Find the ratios of the moles for each substance by dividing each by the smallest one, i.e., normalize to the smallest. 3. If necessary, multiply smallest factor that clears any fractions that they contain. OFB Chapter 1 33 OFB Chapter 1 34 Exercise Calculate the chemical amount (in moles) of carbon dioxide and water using the table of atomic masses. 2. Find the ratios of the moles for each substance by dividing each by the smallest one, i.e., normalize to the smallest. 3. If necessary, multiply smallest factor that clears any fractions that they contain. Volume and Density Exercise 1-13 The density of liquid mercury at 20 deg C is g cm -3. A chemical reaction requires mol of mercury. What volume (in cubic centimeters) of mercury should be measured out at 20 C? Strategy: 1. Use density and mass to find volume. Rearrange m m d = V = V d 2. Density is given, can find mass from the number of moles of mercury which is given 3. Solve for volume. OFB Chapter 1 35 OFB Chapter 1 36

10 Exercise 1-13 The density of liquid mercury at 20 deg C is g cm -3. A chemical reaction requires mol of mercury. What volume (in cubic centimeters) of mercury should be measured out at 20 C? 1. Use density and mass to find volume. m V = d 2. Density is given, can find mass from the number of moles of mercury which is given Chapter 1 The Atomic Nature of Matter Examples / Exercises All (1-1 thru 1-13) Problems 9, 10, 11, 17, 19, 20, 29, 37, 38, 55, Solve for volume. OFB Chapter 1 37 OFB Chapter 1 38