1. Intro Chem Grading and Class Policy handout Book : Introductory Chemistry, 5 th ed., Zumdahl Website: http://schuder.wikispaces.comhttp://schuder.wikispaces.com Contact: Mr. Schuder, 302-545-2116 cell rschuder@charterschool.org Safety contract Right to Know 1

2. Lab Safety Summary Wear goggles in the lab. Do not wear open toed shoes in the lab. No horseplay. Label glassware. Remember that hot and cold glassware appears the same. Handle chemicals with care. Keep your work space clean. 2

3. Unit 1 - Measurement Upon completion of this unit, you should be able to do the following: 1.Define chemistry. 2.Describe how the terms experiment, hypothesis, theory and law fit into the scientific method. 3.Understand safe procedures for working in the lab. 4.Convert numbers from decimal notation to scientific notation and vice-versa. Perform calculations using scientific notation. 5.List and define basic SI units and prefixes. 6.Differentiate between precision and accuracy. 7.Determine percentage error in a measurement. 8.Determine the number of significant figures in a measurement and a calculated result. 9.Demonstrate how dimensional analysis can be used to solve various types of problems 10.Define density and its units. Perform calculations using the relationship among density, mass and volume. 3

4. Chemistry is the science that deals with the materials of the universe, called matter, and the changes that matter undergoes. This course will show how the concepts of chemistry allow us to understand the nature of chemical changes and help us manipulate natural materials to our benefit. Most of what occurs in the world around us involves chemical changes, where one or more substances become different substances. Chemistry: An Introduction 4

5. Examples of chemical changes: Wood burns in air forming carbon dioxide and water The steel in a car rusts Grape juice ferments to form wine Emissions from a power plant lead to the formation of acid rain Chemistry: An Introduction 5

6. Scientific Method All scientific studies follow the same approach to examining a problem. The scientific method requires that we: 1.Make observations. A qualitative observation does not involve a number. It involves properties like color, odor or appearance. A quantitative observation involves a number, like mass or volume, and includes units. It is called a measurement. 2.Formulate a hypothesis, a possible explanation for the observations. 3.Perform experiments to test the hypothesis. 6

7. Scientific Method Once we have a hypothesis that agrees with our observations, we assemble them into a theory that is often called a model. A theory is a set of tested hypotheses that explains some part of nature. An observation can be witnessed and recorded. A theory is an interpretation, a possible explanation of the behavior of nature. Theories change as more information becomes available. Can you think of examples of theories that have changed? 7

8. Scientific Method Zumdahl, page 8 8

9. Scientific Method As we observe nature, we often see the same observations apply to different systems. For example, studies of innumerable chemical changes have shown that the total mass of materials involved is the same before and after the chemical change. We can formulate this observed behavior into a statement called a natural law. The observation that the total mass of a material is not affected by a chemical change is called the law of conservation of mass. 4 grams of hydrogen reacts with 32 grams of oxygen to form 36 grams of water 2H 2 + O 2 → 2H 2 O 9

10. Scientific Method A law is a summary of observed behavior. A theory is an explanation of the behavior. A law tells what happens. A theory tells why it happens. Homework Read Zumdahl pages 1-9 10

11. Lab 1 - Density of Salt Hypothesis – Salt Sense® has 33% less sodium than regular table salt. Experiment - Measure mass of stated volumes of Salt Sense® and table salt. Data - Density of Salt Sense = Density of table salt = Conclusion - 11

12. Measurement Chapter 2 1.Scientific Notation 2.Units 3.Measurements of Length, Volume and Mass 4.Uncertainty in Measurements 5.Significant Figures 6.Dimensional Analysis 7.Temperature 8.Density 12

13. Scientific Notation Numbers associated with scientific measurements and calculations are often very small or very large. Scientific notation is a method used to make writing very large or very small numbers more compact and easier to use. 13

14. Scientific Notation Consider the number 125. It can be written as 1.25 x 100. Since 100 = 10 x 10 = 10 2, we can write 125 = 1.25 x 10 2 1.25 x 10 2 is in scientific notation, a format that expresses a number as the product of a number between 1 and 10 and a power of 10. 14

15. Scientific Notation One way to remember scientific notation is to start with the number and count how many places the decimal must be moved to obtain a number between 1 and 10. For 125, it is 2 places to the left. For 0.0036 it is 3 places to the right. When the decimal point is moved to the left, the power of 10 is positive. When the decimal point is moved to the right, the power of 10 is negative. 125 = 1.25 x 10 2 0.0036 = 3.6 x 10 -3 15

16. Units Units define the scale of measurement being used. Units are critically important in science. If you measure a length in feet and record only the number without the units, then someone reading your data assumes the units are mm, there will be confusion. 16

17. Units Two widely used systems of measurement are the English system and the metric system. The International System or SI (le Système Internationale in French) is based on the metric system and is used for much scientific work. Physical Quantity Name of SI UnitAbbreviationName of English Unit Abbreviation MassKilogramkgPoundlb LengthMetermFootft TimeSeconds s TemperatureKelvinKFahrenheitF 17

18. Metric prefixes Changing the prefix alters the size of a unit. PrefixSymbolMeaningSI gigaG100000000010 9 megaM100000010 6 kilok100010 3 millim.00110 -3 microμ.00000110 -6 nanon.00000000110 -9 18

19. Measurement Observations are a key component of science They can be qualitative or quantitative. A quantitative observation is called a measurement. Measurements have two parts: a number and units. 19

20. Measuring Length Zumdahl page 20 20

21. Measuring Volume Volume is the amount of space that an object occupies. The base metric unit is the liter (L). The common unit used in the lab is the milliliter (mL). One milliliter is exactly equal to one cm 3. Cubic centimeters (cm 3 ) is sometimes designated cc. The derived SI unit for volume is the m 3 which is too large for convenient use. 21

22. Measuring Volume Zumdahl page 20 22

23. Measuring mass Mass is the quantity of matter in an object. The SI unit of mass is the kilogram (kg). However, in the lab, the gram (g) is more commonly used. 23

24. Uncertainty in Measurement When measurements are made, an estimate is required. See example below. Zumdahl page 23 24

25. Uncertainty in Measurement Note the first two digits are the same, no matter who made the measurement. These are certain digits. The third digit is estimated and can vary. It is an uncertain digit. When making a measurement, record all the certain digits and the first uncertain digit. PersonPin Length 12.85 cm 22.84 cm 32.86 cm 42.85 cm 52.86 cm 25

26. Significant Figures The numbers recorded in a measurement are called significant figures. The number of significant figures for a given measurement is determined by the uncertainty of the measurement. The ruler use for measuring the pin can give results only to the hundredth of a cm. When we record the significant figures for a measurement, we give information about the uncertainty of the measurement. The ruler measurement of 2.86 cm is uncertain in the third digit. 26

27. Rules for Counting Significant Figures Nonzero integers always count as sigificant figures. A measured volume of 15.23 mL has 4 sig figs. Exact numbers have an unlimited number of significant figures. Counted numbers are exact. 3 trials of an experiment has unlimited sig figs. Leading zeros are NOT significant. A mass of 0.032 grams has 2 sig figs. Captive zeros count as significant figures. A mass of 10.08 grams has 4 sig figs. Trailing zeros are significant if written with a decimal point. A volume of 100 mL has 1 sig fig. 100. mL has 3 sig figs. 100.0 mL has 4 sig figs. 27

28. Significant Figures The number of significant digits is independent of the decimal point. 645. 64.5 6.45.645.0645 These numbers all have 3 significant figures. 28

29. Rounding Off Numbers After calculations, you may need to round off. Carry all digits until the final calculation. The round off to the proper number of significant figures. If the first insignificant digit is 5 or more, round up. If the first insignificant digit is less than 5, round down. 29

30. Significant Figures in Calculations For multiplication or division, the number of significant figures in the result is the same as the measurement with the smallest number of sig figs. This is the limiting term. 4.56 x 1.4 = 6.384 6.4 (2 sig figs) For addition or subtraction, the limiting term is the one with the smallest number of decimal places. 12.11 + 18.0 + 1.013 = 31.123 31.1 (1 decimal place) 30

31. Accuracy and Precision All measurements have some uncertainty. Uncertainty is measured with accuracy and precision. Accuracyis a measure of how close the measurement is to the true value. Precision is a measure of how close measurements are to each other. 31

32. Accuracy and Precision Slide courtesy of Mark Case, Emmaus HS 32

33. Homework Read Pages 14-28, as needed Pages 46 - 48; problems 8, 10a-f, 18, 20, 26, 36, 38, 40, 42 and 52. 33

34. Dimensional Analysis There are times when the units of a measurement are not the same as the units in which you seek an answer. For example, medicine is dosed based on mass in kg, but the English unit for mass is pounds. How many kg is 180 pounds? An equivalence statement defines the relationship between different units. 1 kg = 2.205 lbs 34

35. Dimensional Analysis A conversion factor is a ratio of the two parts of the statement that relate the two units. Examples: 2.205 lbs 1 kg 2.205 lbs 35

36. Dimensional Analysis Choose the conversion factor that cancels the units you don’t want. (factor label method) How many kg is 180 pounds? 180 lbs x = 81.63 kg In equations, units cancel just as numbers do. Always check that your answer makes sense. The process of converting from one unit to another is called dimensional analysis. 1 kg 2.205 lbs 36

37. Dimensional Analysis Practice Page 45 – In Class Question 1, parts a, b and c Pages 48 and 49 – Problems 59 and 63 Homework Read pages 28-33 as needed Pages 48 and 49; Problems 58, 60, 64 and 66 37

38. Temperature English units are degrees Fahrenheit ( o F) Metric units are degrees Celsius ( o C) Temperature in Sardegna, Italy 22 o C Is it hot, cold, moderate? 38

39. Temperature Conversion o C = ( o F – 32) * 5/9 o F = ( o C * 9/5) + 32 Freezing point of water 0 o C32 o F Boiling point of water 100 o C212 o F Typical body temperature 37 o C98.6 o F Equal point of Celsius and Fahrenheit scales -40 o C-40 o F 39

40. Abolute Temperature Scale Kelvin is an absolute temperature scale. It does not use the degree notation, just K. 0 K = -273.15 o C Absolute scale is needed when making thermodynamic calculations. 40

41. Density Extensive properties Depend on the quantity of sample measured. Example - mass and volume of a sample Intensive properties Independent of the sample size. Properties that are often characteristic of the substance being measured. Examples - density, temperature, melting and boiling points. 41

42. Density Density is an intensive property of a substance based on two extensive properties. It is defined as the amount of matter in a given volume. Common units are g/cm 3 or g/mL. Density = Mass Volume 42

43. Density What is the density of 5.00 mL of a fluid if it has a mass of 5.23 grams? d = mass / volume d = 5.23 g / 5.00 mL d = 1.05 g / mL What would be the mass of 1.00 liter of this sample? 43

44. Homework Read pages 29 - 47 as needed page 50; problems 92, 93, 94, 96, 98 In class practice page 50; problem 91 44