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Chemistry 100(02) Fall 2013

Chemistry 100(02) Fall 2013. Instructor: Dr. Upali Siriwardane e-mail : upali@coes.latech.edu Office : CTH 311 Phone 257-4941 Office Hours : M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F 8 :00 - 10:00 a.m.   Or by appointment Test Dates :.

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Chemistry 100(02) Fall 2013

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  1. Chemistry 100(02) Fall 2013 Instructor: Dr. UpaliSiriwardane e-mail: upali@coes.latech.edu Office: CTH 311 Phone257-4941 Office Hours: M,W, 8:00-9:30 & 11:30-12:30 a.m Tu,Th,F8:00 - 10:00 a.m.   Or by appointment Test Dates: September 30, 2013 (Test 1): Chapter 1 & 2 October 21, 2013 (Test 2): Chapter 3 & 4 November 13, 2013 (Test 3) Chapter 5 & 6 November 14, 2013 (Make-up test) comprehensive: Chapters 1-6 9:30-10:45:15 AM, CTH 328

  2. REQUIRED: Textbook:Principles of Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro - Pearson Prentice Hall and also purchase the Mastering Chemistry Group Homework, Slides and Exam review guides and sample exam questions are available online: http://moodle.latech.edu/ and follow the course information links. OPTIONAL: Study Guide: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro 2nd Edition Student Solutions Manual: Chemistry: A Molecular Approach, 2nd Edition-Nivaldo J. Tro2nd Text Book & Resources

  3. Chapter 1.Matter, Measurement, and Problem Solving 1. 1 Atoms and Molecules………………………………….. 1 1 .2 The Scientific Approach to Knowledge…………….. 3 1 .3 The Classification of Matter…………………………… 5 1 .4 Physical and Chemical Changes and Physical and Chemical Properties…………………………………….. 9 1 .5 Energy: A Fundamental Part of Physical and Chemical Change…………………………………………………….. 12 1 .6 The Units of Measurement……………………………... 13 1 .7 The Reliability of a Measurement……………………… 20 1 .8 Solving Chemical Problems……………………………. 27

  4. Chapter 1. KEY CONCEPTS

  5. Units

  6. SI UNITS OF MEASUREMENT Five Basic Units • Length meter (m) • Mass kilogram (kg) • Time second (s) • Temperature kelvin (K) • Amount mole (mol) • 6.02 x 1023 units

  7. The Units of Measurement 1) Give the name and abbreviation of the SI Unit for: • Length b) Mass c) Time d) Amount of substance e) Temperature

  8. Metric System How to change measurements to reasonable numbers

  9. Macroscale, Microscale, and Nanoscale

  10. 2) Give the abbreviation for the following units and describe what they are used to measure: • cubic centimeter b) micrometer c) nanoseconds d) millimole

  11. 3) Give the name and the abbreviation (without looking in the book) of the SI or metric prefix for: • 10-12b) 106 c) 10-9 d) 10-2 e) 10-3f) 109g) 103i) 10-6

  12. Uncertainty in a Measurement types of errors, random and systematic. • The last digit is an estimate. • . • Careless measurements • Low resolution instruments • Calibration errors Measurement ≈ 26.13 cm

  13. Uncertainty and Significant Figures All measurements involve some uncertainty. Scientists write down all the digits that have no uncertainty plus one additional uncertain digit. If an object is reported to have a mass = 6.3492 g, the last digit (“2”) is uncertain ( it is probably close to 2, but may be 4, 1 …). There are five significant figures in this number. All the digits are meaningful.

  14. Uncertainty and Significant Figures placeholders To find the number of significant figures: Read a number from left to right and count all digits, starting with the first non-zero digit. Alldigits are significant except those zeros that are used to position a decimal point (“placeholders”). 0.00034050 5 sig. figs. Scientific Notation (3.4050 x 10-4) significant significant

  15. Significant Figure Primer • All non-zero values are significant. • Example: 536 has three significant figures. • Zeros between non-zero digits are also significant. • Example: 6703 has four significant figures. • Place holder zeros: • Zeros to the left of a non-zero digit are NOT significant. • Example: 0.0043 has two significant figures. • Example: 0.0600 has three significant figures. • Zeros to the right or after a non-zero digit to the decimal point are NOT significant. • Example: 7000 has only one significant figure. • Example: 32040 has four significant figures. • Example: 50.0 has three significant figures; all the zeros are significant.

  16. Uncertainty and Significant Figures Examples Number Sig. figs. Comment 2.12 3 4.500 4 The zeros are not placeholders. They are significant. 0.002541 4 The zeros are placeholders (not significant). 0.00100 3 Only the last two zeros are significant. 500 1, 2, 3 ? Ambiguous. If a number lacks a decimal point the zeros may be placeholders or may be significant. 500. 3 Adding a decimal point is one way to show that the zeros are significant. 5.0 x 102 2 No ambiguity.

  17. The Reliability of a Measurement 4) Express the following numbers in scientific notation with the appropriate number of significant figures: • 10980000000 b) 414100 c) 0.000095162 d) 746.5 x 107

  18. Rounding Look at the 1st non-significant digit (the digit after the last one retained). If it: • is > 5, round the last retained digit up by 1. • is < 5, make no change. • equals 5, and the 2nd non-significant digit is: • absent, round the last retained digit up by 1. • odd, round the last retained digit up by 1. • even, make no change. Consider rounding 37.663147 to 3 significant figures. last retained digit 2nd non-significant digit 1st non-significant digit It rounds up to 37.7

  19. Rounding Examples Round the following numbers to 3 significant figures: 1st non-sig. 2nd non-sig. Rounded Number digit digit Number 2.123 2.123 - 2.12 51.372 51.372 51.372 51.4 131.5 131.5 - 132. 24.752 24.752 24.752 24.7 24.751 24.751 24.751 24.8 0.067440.06744-0.0674

  20. Can you define accuracy vs. precision? Precision = the degree of exactness of a measurement that is repeatedly recorded. Accuracy = the extent to which a measured value agrees with a standard value • Which set is more precise? 18.2 , 18.4 , 18.35 17.9 , 18.3 , 18.85 16.8 , 17.2 , 19.44

  21. Can you hit the bull's-eye? Three targets with three arrows each to shoot. Both accurate and precise Precise but not accurate Neither accurate nor precise How do they compare?

  22. Unit Conversion Calculation • Speed of light is 3.00 x 108 m s-1 . Convert the speed of light to miles per year (1 mile = 1.61 km).

  23. Using Conversion Factors 5) Convert 78.01 inches into: a) feet b) meters 6) Convert 15.42 meters into: a) kilometers b) micrometers

  24. Example: 301 in scientific notation is 3.01 x 102. NOTE: The decimal point was moved two places to the left. Example: 0.0301 in scientific notation is 3.01 x 10-2.NOTE: The decimal point was moved two places to the right. Both of these value indicate THREE significant figures. Significant Figures and Scientific Notation

  25. Significant Figures in Add/Sub adp = 3 adp = 4 The answer you report in a problem should only include significant digits. Addition and subtraction Find the number of digits after the decimal point (adp) in each number. answer adp = smallest input adp. Example Add: 17.245 +0.1001 17.3451 Rounds to: 17.345 (adp = 3)

  26. Addition and Subtractions Examples: • Examples: 2 3 .4 6 7 in • + 3 1 3 .2 1 in • 3 3 6 .6 7 7 in but you would report 336.68 in • 4 5 7 cm • - 0 . 6 8 cm • 4 5 6 . 3 2 cm but you would report 456 cm

  27. Significant Figures Add/Sub adp = 2 adp = 4 Example Subtract 6.72 x 10-1 from 5.00 x 101 Write the numbers down with the same power of 10: 5.00 x 101 – 0.0672 x 101 4.9328 x 101 Rounds to: 4.93 x 101 adp= 2

  28. Significant Figures Mul/Div sig. fig. = 5 sig. fig. = 4 Multiplication and Division Find the number of significant figures (sig. fig.) in each number. Answer has sig. fig = smallest input sig. fig. Example Multiply 17.425 and 0.1001 17.245 x 0.1001 1.7262245 Rounds to: 1.726 sig. fig. = 4 Example Multiply 2.346 x 12.1 x 500.99 Rounds to: 1.42 x 104(3 sig. fig.) = 14,221.402734

  29. 7) Perform the following calculations and give the answer in the correct number of significant figures. • 30.84 + 9.74 Answer: Scientific notation: b) 30.84 + 9.74486 Answer: Scientific notation: c) 145 + 1.54 x 10-6 Answer: Scientific notation:

  30. 7) Perform the following calculations and give the answer in the correct number of significant figures. d) 40.79 - 1.18432 Answer: Scientific notation: e) 1.43 x 0.848 Answer: Scientific notation: f) (7.601x107) x (8.09x10-4) Answer Scientific notation:

  31. Temperature Scales: Comparison

  32. Temperature Conversions Human body temperature is 98.6 oF. Convert this temperature to oC and K scale oC = 5/9 (98.6 - 32) = 5/9 (66.6) = 37.0 oC--> K = 37.0 oC +273.15 = 310.2 K

  33. Temperature Convesions oF -- > oC ; C = 5/9 (F - 32) oC -- > oF ; F =9/5 C + 32 oC -- > K ; K = C + 273.15 8) Convert 98.6 °F into: a) °C b) K.

  34. Problem Solving by Factor Label Method • State question in mathematical form • Set equal to piece of data specific to the problem • Use conversion factors to convert units of data specific to problem to units sought in answer • Other names used Unit Conversion Method or dimensional (Unit) Analysis

  35. Common Conversion Factors Length1 kilometer = 1000 m = 0.62137 mile 1 inch = 2.54 cm (exactly) 1 angstrom (Å) = 1 x 10-10 m Volume1 liter (L) = 1 x 10-3 m3 = 1000 cm3 = 1000 mL = 1.056710 quarts 1 gallon = 4 quarts = 8 pints Mass1 amu = 1.6606 x 10-24 g 1 pound = 453.59237 g = 16 ounces 1 ton (metric) = 1000 kg 1 ton (US) = 2000 pounds

  36. Solving Chemical Problems 9) Convert 75 miles per hour into: m s-1.

  37. Solving Chemical Problems 10) Convert 100 m2 into cm2 11) Convert 1 m3 into cm3 .

  38. Significant Figures and Mathematical Operations Mathematical operations dictate the reporting of significant figures in an answer. • Multiplication and Division The least precise measured value determines the number of significant figures in the reported answers. • Addition and Subtraction The value with the smallest decimal measurement determines the answer’s significant figure.

  39. Solving Chemical Problems 13) Perform the following mathematical operations and give the answer with the correct number of significant figures :

  40. Density • Density = mass (g) • volume (cm3) • An INTENSIVE physical property • The physical property does not depend on amountof substance. • The physical properties of mass and volume that determine a substance’s density are EXTENSIVE. • Extensive physical properties are dependent on amount. • Densities of liquid and gases are affected by temperature.

  41. Density Calculations PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds? Strategy: 1. Convert mL to cm3. 2. Solve for mass (in grams) using density relationship. 3. Convert grams to pounds.

  42. A Density Calculation PROBLEM: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds? Density = mass (g) volume (cm3) Step 1: 95 mL x (1cm3/1mL) = 95 cm3 Step 2: 13.6 g/cm3 = x / 95 cm3 x = 1.29 x 103 g, but report 1.3 x 103 g Step 3: 1.3 x 103 g x (1 lb/454 g) = 2.9 lb

  43. Solving Chemical Problems 12) Aluminum block weighs 14.2 g and has a density of 2.70 g cm-3. Calculate the volume of the block.

  44. Problem: The density of octane (C8H18) is 7.00 lb/gal. a) What is density in mg/cm3? b) What is the mass in grams of 1.25 liters of octane? Strategy: 1. Convert 7.00 lbs to mg. 2. Change gallons to cm3. 3. Determine the density of octane in mg/cm3. 4. Convert 1.25 L to mL. 5. Determine the mass of octane in 1.25 L using density.

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