Chem 142, 152, and 162

Transcription

1 CHEMISTRY 142 A Summer 2011 Dr. Kari Pederson LECTURES: MON, WED, and FRI 9:40 10:40 am Prelabs due Tues Quiz Sections Tues Homework due Wed Lab Sessions Tues/Thurs 1 Chem 142, 152, and 162 Zumdahl, 6 th Edition, Custom Split Volumes All volumes include the complete Student Solutions Manual at no additional cost. The Chem 142 volume includes a lifetime-of-edition WebAssign access code. If you plan to take Chem 152 or 162 you may consider purchasing a full version of the textbook. 2 1

2 Chemistry 142 Text: Chemical Principles - By Steven S. Zumdahl Chapter #1 : Chemists and Chemistry Chapter #2 : Atoms, Molecules, and Ions Chapter #3 : Stoichiometry: Mole - Mass Relationships in Chemical Systems Chapter #4 : Chemical Reactions and Solution Stoichiometry Chapter #5 : Gases and the Kinetic - Molecular Theory Chapter #6 : Chemical Equilibrium Chapter #7 : Acids and Bases Chapter #8 : Applications of Aqueous Equilibria 3 Weight of Course Components Participation (worksheets in Quiz section) 5% Clickers (1st week dropped) 5% Homework (ALEKS) 20% Laboratory 15% 2 midterm exams (1 hr each) 30% Final exam (1 hr) 25% Total 100% 4 2

3 Exams Midterm Exam #1 Friday, July 15 th Midterm Exam #2 Friday, August 5 th Final Exam Friday, August 19 th Exams begin at 9:40 am. Be in your seat at least 5 minutes before exam starts. Bring a prepared Scantron form, a few pencils, a non graphing calculator, and a photo I.D. to every exam. All exams are cumulative. 5 Success!! Depends on each of us: lecturer, TA, YOU! Policies/Expectations Read the syllabus and front pages of lab manual Structure Homework (every week) Clickers and Worksheets (every week) Exams (2 Midterms and 1 Final) Labs (4 labs, most reports due in lab) (DON T miss the deadlines for HWs, quizzes, or prelabs!) Resources for help with concepts/assignments TA and instructor office hours Reading - DO IT! 6 3

4 Hard Practice, Easier Exams Practice what you are to do on the exams. Work many problems. Develop skills for problem analysis. You must develop your own problem-solving skills and learn how to approach problems yourself! 7 Resources to Help You Navigate the General Chemistry Series Help with Lecture, Exams, Homework Your TA Course Instructor (see front page of syllabus for office hours & contact info.) Help with Prelabs and Lab Reports (lab content) Your TA Lab Instructor (see front page of the syllabus for contact info.) Missed Exams Tracy Harvey (harvey@chem.washington.edu) Catalyst Tools (Prelabs) Your TA Online Catalyst help No extensions for missed WebQ deadlines! Add/Drop Codes UG Services BAG 303 (in person only, Mon Fri 8am to 4pm) YOU ALEKS (HW) Your TA Online Support No extensions for missed deadlines! Missed Labs UG Stockroom to reschedule UGlabs@chem.washington.edu BAG 271 Tracy Harvey (if you can t reschedule) (harvey@chem.washington.edu) Your TA Clickers UG Stockroom to test clicker function * communication is best reserved for administrative issues. For help with course or lab content, please try to attend TA or instructor office hours. 8 4

5 Chapter# 1 : Chemists and Chemistry 1.1 Thinking Like a Chemist 1.2 A Real-World Chemistry Problem 1.3 The Scientific Model 1.4 Industrial Chemistry 1.5 Polyvinyl Chloride (PVC) Real-World Chemistry 9 Chemistry as a Central Science Oceanography Atmospheric Sciences Physics Medicine Economics Governments Chemistry Geology Anthropology Biology Astronomy Politics People 10 5

6 Figure 1.1: Chemists interact 11 Basic Skills of a Chemist 1. Understand atomic symbols (the basic vocabulary!) B for boron, C for carbon 2. Know how to describe everything with chemical formulas: air (N 2, O 2, CO 2, CO, He, Ne, ); water (H 2 O); table salt (NaCl, NaI, MgCO 3 ); baking soda (Na 2 CO 3 ); vinegar (CH 3 COOH); glucose (C 6 H 12 O 6 ); and soda (H 2 O, CO 2, H 2 CO 3, HCO 3 -, CO 3 2-, sugars, ) 3. Know the structure, reactivity & properties of molecules. 4. Know how to convert between units of measurement (mass, volume, moles,...) 5. Have lab skills: measure, analyze, synthesize, critically evaluate

7 The Scientific Method Information or data is gathered by careful observation of the phenomenon being studied. On the basis of that information a preliminary hypothesis is formed. A series of experiments is devised to test the predictive power of the hypothesis. 13 The Scientific Method (cont.) On the basis of the experimental tests, the hypothesis may be (a) Accepted as scientific law or theory. (b) Modified so that all results are adequately explained. (c) Discarded. 14 7

8 Demo 1, Part 1 Hydrogen Balloon What happened to the balloon? 1. It spontaneously popped all on its own. 2. Something sharp popped it. 3. The heated gas burst the balloon and then ignited. 15 2H 2 (g) + O 2 (g) 2 H 2 O (g) + Energy Hydrogen and oxygen are diatomic gases! (H 2 and O 2 ) Water can be a gas! ENERGY was given off! This is characteristic of an exothermic reaction! This is a balanced chemical reaction! 16 8

9 Demo 1, Part 2 Helium Balloon Helium is less flammable, what will happen? 1. Bigger bang. 2. Smaller bang. 3. Same as hydrogen balloon. 17 Demo 1, Part 3 Hydrogen and Oxygen Mixture What will happen? 1. Bigger bang. 2. Smaller bang. 3. Same as hydrogen balloon. 4. Same as helium balloon 18 9

10 Hydrogen burns, releasing energy! Physics Chemical Reaction He H 2 H 2 +O 2 O 2 19 Figure 1.4: Parts of the scientific method (Qualitative and/or quantitative) (Possible explanation) (Test explanation) (Attempts to explain why something happens) (Summarizes what happens) 20 10

11 CHEMISTRY The Study of Matter and its Properties, the Changes that Matter Undergoes, and the Energy Associated with those Changes

12 The Periodic Table of the Elements 23 Definitions Matter - The stuff of the universe: books, planets, trees, animals - anything that has mass and volume. Composition - The types and amounts of simpler substances that make up matter. Properties - The characteristics that give each substance a unique identity. Physical Properties - those the substance shows by itself, without interacting with another substance (color, melting point, boiling point, density, etc.) Chemical Properties - those that the substance shows as it interacts with, or transforms into, other substances (flammability, corrosiveness, etc.) 24 12

13 Definitions Energy - The capacity to do work! Potential Energy - The energy due to the position of the object or the energy from a chemical reaction. Kinetic Energy - The energy due to the motion of the object. 25 Energy Involved in Phase Changes Gas Liberates Energy Evaporation Condensation Liquid Melting Freezing Requires Energy Solid 26 13

14 All Measured Quantities Consist of a Number and a Unit Length : a car is 12 feet long, not 12! a person is 6 feet tall, not 6! Area : a carpet measuring 3 feet (ft) by 4 ft has an area of: ( 3 x 4 )( ft x ft ) = 12 ft 2 Speed and Distance : A car traveling 350 miles (mi) in 7 hours (hr) has a speed of 350 mi = 50 mi 7hr hr In 3 hours the car travels: 3 hr x 50 mi = 150 mi hr 27 FUNDAMENTAL UNITS (Systeme International) All other units are derived from this system

15 FYI.. Derived SI Units Quantity Definition of Quantity SI unit Area Length squared m 2 Volume Length cubed m 3 Density Mass per unit volume kg/m 3 Speed Distance traveled per unit time m/s Acceleration Speed changed per unit time m/s 2 Force Mass times acceleration of object kg m/s 2 ( =newton, N) Pressure Force per unit area kg/(m 2 ) ( = pascal, Pa) Energy Force times distance traveled kg m 2 /s 2 29 ( = joule, J) 30 15

16 FYI.. 31 Temperature Scales and Conversions Kelvin ( K ) - The Absolute temperature scale begins at absolute zero and has only positive values. Celsius ( o C ) - The temperature scale used in the sciences, formally called centigrade and most commonly used scale around the world; water freezes at 0 o C and boils at 100 o C. Fahrenheit ( o F ) - Commonly used scale in the U.S. for our weather reports; water freezes at 32 o F and boils at 212 o F

17 33 Temperature Conversions Know these!! T (in K) = T (in C) T (in C) = T (in K) T (in F) = 9/5 T (in C) + 32 T (in C) = [ T (in F) - 32] 5/

18 35 General Strategy for Unit Conversions 1. Unity factor 2. Rearrange to form: What you want What you ve got 3. Multiply numbers 4. Cancel units 36 18

19 Conversion Factors : Unity Factors - I Equivalent factors can be turned into conversion factors by dividing one side into the other! 1 mile = 5280 ft 1 mile = 1 or 5280 ft = ft 1 mi 1 in = 2.54 cm 1 in = 1 or 2.54 cm = cm 1 in In converting one set of units to another, and the old one is canceled out! Convert 29,141 feet into miles: 29,141 ft 1 mile 5280 ft = miles 37 Conversion Factors - II 1.61 km = 1 mi or 1.61 km = 1 1 mi Convert miles in to kilometers: mi 1.61 km 1 mi = 8.89 km Conversions in the metric system are easy: 1 km = 1000 m 1 m = 100 cm 1 cm = 10 mm Example: convert 8.89 km into m and cm! 8.89 km 1000 m = 8890 m 1 km 8890 m 100 cm 1 m = cm 38 19

20 Conversion Factors - III Convert 3.56 lbs/hr into units of mg/s Multiple conversion factors! 1 kg = lbs, 1 hr = 60 min, 1 min = 60 s Look at it as two parts: 1) mass and 2) time 3.56 lbs 1 hr 1 kg 2.21 lbs 1000 g 1 kg 1000 mg 1 g 1 hr 60 min 1 min 60 sec = 448 mg/s

21 Steps for Determining Significant Digits All digits are significant, except zeros that are used only to position the decimal point. 1. Look for the decimal point. 2. Start at the left of the number and move right until you reach the first nonzero digit. 3. Count that digit and every digit to it s right as significant. 4. Zeros that end a number and lie either after or before the decimal point are significant. (1.030 ml and 5,343 L each have four significant figures) 41 Examples of Significant Digits in Numbers Number Sig. digits Number Sig. digits L two x 10 7 nm six g four 5600 ng two kg two 87,000 L two L six 78,002.3 ng six g four g four 875,000 oz three x 10-6 L four 30,000 kg one oz five m 3 five kg six 23, lbs seven ml nine g three kg eight 1,470,000 L three 1,000,000,000 g one 42 21

22 Rules for Sig Figs in Calculations For multiplication and division. The answer contains the same number of significant figures as the measurement with the fewest significant figures. Multiply the following numbers: 9.2 cm x 6.8 cm x cm = cm 3 = 23 cm 3 For addition and subtraction. The answer has the same number of decimal places as the measurement with the fewest decimal places. Example, adding two volumes 83.5 ml ml = ml = ml Example subtracting two volumes: ml ml = ml = ml Rules for Rounding off Numbers Calculations: Carry as many digits as possible until the final answer. If the rounding digit is less than 5 the preceding digit remains the same rounds to 8.31 If the rounding digit is equal to or greater than 5 the preceding digit increases by rounds to rounds to

23 How to Solve Chemistry Problems 1) Problem: State all of the information needed to solve the problem. Write down the knowns and unknowns. Include as known any definitions or laws that might mathematically relate the values. 2) Plan: Suggest the steps needed to find the solution. Develop a roadmap solution. Look for additional information / laws as needed. 3) Solution: Do the algebraic manipulations & calculations in the order planned. 4) Check: Is the result what you expect or at least in the same order of magnitude! Does it have the right units? 45 An Example Problem! The volume of an irregularly shaped solid can be determined from the volume of water it displaces. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of iron, is submerged in the water, the volume increases to ml. What is the volume of the piece of Pyrite in cm 3 and in liters? 46 23

24 STEP ONE EXTRACT RELEVANT INFORMATION FROM THE QUESTION!!!! THIS IS ALWAYS STEP ONE! 47 An Example Problem! The volume of an irregularly shaped solid can be determined from the volume of water it displaces. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of Iron, is submerged in the water, the volume increases to ml. What is the volume of the piece of Pyrite in cm 3 and in L? Vol (ml) = ml ml = 70.8 ml Vol (cm 3 ) = ml 1 cm 3 = 70.8 cm 3 1 ml Vol (L) = ml 1 L 1000 ml = 7.08 x 10-2 L 48 24

25 A Review Problem! The mass of object can also be determined from the displacement of water. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of iron, is submerged in the water, the volume increases to ml. The density of Pyrite is 2.4 g/ml. What is the mass of the piece of Pyrite in kg and lbs? 49 STEP ONE EXTRACT RELEVANT INFORMATION FROM THE QUESTION!!!! THIS IS ALWAYS STEP ONE! 50 25

26 A Review Problem! The mass of object can also be determined from the displacement of water. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of iron, is submerged in the water, the volume increases to ml. The density of Pyrite is 2.4 g/ml. What is the mass of the piece of Pyrite in kg and lbs? 51 The volume of the piece of pyrite in ml is: ml ml ml ml 52 26

27 A Review Problem! The mass of object can also be determined from the displacement of water. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of iron, is submerged in the water, the volume increases to ml. The density of Pyrite is 2.4 g/ml. What is the mass of the piece of Pyrite in kg and lbs? Vol (ml) = ml ml = 71.7 ml 53 The mass of the Pyrite in kg is: kg kg kg kg 54 27

28 A Review Problem! The mass of object can also be determined from the displacement of water. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of iron, is submerged in the water, the volume increases to ml. The density of Pyrite is 2.4 g/ml. What is the mass of the piece of Pyrite in kg and lbs? Mass (g) = Mass (kg) = 71.7mL 2.4 g 1 ml g 1 k g 1000 g = g = kg Sig Figs! = 0.17 kg 55 The mass of the Pyrite in lbs is: lbs lbs lbs lbs 56 28

29 A Review Problem! The mass of object can also be determined from the displacement of water. A graduated cylinder contains ml water. When a small piece of Pyrite, an ore of iron, is submerged in the water, the volume increases to ml. The density of Pyrite is 2.4 g/ml. What is the mass of the piece of Pyrite in kg and lbs? Mass (lbs) = kg 2.21 lbs = kg Sig Figs! = 0.38 lbs 57 Precision and Accuracy Errors in Scientific Measurements Precision - Refers to reproducibility or how close the measurements are to each other! Accuracy - Refers to how close a measurement is to the real value! Systematic error - produces values that are either all higher or all lower than the actual value. Random Error - in the absence of systematic error,produces some values that are higher and some that are lower than the actual value

30 Precise AND Accurate Precise NOT Accurate Accurate NOT Precise NEITHER Precise NOR Accurate 59 30