Welcome to Chemistry! with Mrs. Strain Rm. 403

1. Welcome to Chemistry!with Mrs. Strain Rm. 403 • Do Now:.Findyourseat & distribute papers on desk. • HWK for all classes: • Read “Welcome to Chem A” page on my website to understand procedures for this class. Complete hand out and return on Tuesday, 9/8/14 • Start “Intro to WebAssign”: due Monday • Supplies & Signature Sheet due ASAP but no later than Tuesday, 9/8/14 • Check board for other HW assignments based on class period • Safety Quiz – on lab day AFTER completing safety lab. Review safety contract info • Math Assessment on first non-lab day starting Monday(no need to study )

2. What is Chemistry? • Your Task: on a piece of paper answer this question... What is Chemistry? • Does your answer sound like any of these responses?

3. What is Chemistry? • The definition we’ll work off of this year: Chemistry is the study of matter& of the changesit undergoes • Composition • Structure • Properties • Energy changes

4. A Quick Demo • If we want to describe matter & its changes, there is a certain language we need to become familiar using. • There are good observations & there are bad observations. • During the demo: Write down what you see happening. Imagine you were trying to explain this to someone who is not present in the room.

5. Two Types of Measurements • In science we can take two different kinds of observations: • Qualitative • Quantitative • Qualitative (think “quality”): observations using words • Quantitative (think “quantity”): observations using numbers and units.

6. Here’s what I am hoping to see… • Qualitative observations: • States of matter • Color • Texture • Smell • Viscosity • Quantitative observations: • Amount of substances present • Step by step procedure!

7. Here’s what I don’t want to see… • Opinionated language • “I feel” • “I like” • Non-specific wording • “sort of…”, “lots of…”, ”kinda” • Descriptions that sound like a kindergartener wrote them • “It was all bouncy and …” • describing something as “chunky”

8. Taking Measurements in Chemistry Ch. 2 The SI or Metric System

9. The SI System • Around 1793, scientists all over the world began to agree upon a single measurement system called • Le Systeme International d’ Unites or SI System • 7 base units • The idea was to create a unifying system of weights and measurements

10. Crash Course: Units • Where’s volume??

11. Derived Units m m D D V V • Combinations of base units • Volume: amount of space taken up by an object • Derived SI unit is cubic meter, m3 • More often we use cm3 = mL • Density: ratio of mass to volume • g/cm3 of g/mL or g/L • Does not change for a given substance D = m V D = m V

12. Other Derived Units

13. Larger quantities

14. Using SI prefixes: Number Line Method Conversions from one SI prefix to another (within 1 of the 7 base units) can easily be preformed by moving the decimal place of a quantity by 1 space or 3, left or right.

15. Practice Problems • 5.6 cm to m • 56 mg to g • 340 mm to cm • 1.2 ML to L 0.056 m 0.056 g 34 cm 1,200,000 L

16. Using SI prefixes: Factor-Label Method (Dimensional Analysis) • Method requires translating two equal quantities into a ratio or conversion factor • Ex: 16 oz = 1 lb can be written 16 oz or 1 lb 1 lb 16 oz • Notice: a conversion factor can be represented 2 ways! • This can be done with any 2 equal quantities • 2 grand slams = 8 R.B.I.’s • 1 fortnight = 14 days • 100 cm = 1 m

17. Using SI prefixes: Factor Label Method • Using the factor label method to solve problems • Ex: How many dimes are in 14 dollars? • Write the given • Write conversion factor • Solve, crossing out units that have divided out 14 dollars x 10 dimes = 14o dimes 1 dollar

18. Using Factor-Label Method • Sample Problems: Converting 9.8 g to kg 9.8 g    x   1 kg        = 0.0098 kg               1000. g Converting 9.8 kg to g 9.8 kg    x   1000. g        = 9800 g                   1 kg “1” goes in front of larger unit!

19. Practice Problems • Try these practice problems, but now using the Factor-Label Method • (I realize this seems like more work than the number line method…but there’s a reason why we have to learn this) • 5.6 cm to m • 1.2 L to ML • 100 mm to cm • 25 kg of water to mL 0.056 m 1.2 x 10-6 ML 10 cm 2500 mL

20. Do Now: Test your Metric System “With-it-ness” • For each of the measurements on your worksheet, decide the appropriate quantity that should be assigned to it.

21. Density Practice m D V • Density Formula • Use Density Pyramid as a short cut D = m V m D V

22. Taking Measurements in Chemistry Accuracy vs. Precision

23. Accuracy & Precision in Measurements • Accuracy: closeness of measurements to correct value • Precision: closeness of a set of measurements to each other (assuming they’re made in the same way)

24. Low accuracy Low precision High accuracy High precision Low accuracy High precision

25. Accuracy vs. Precision • Example: A student measures the density of a sample of nickel. • The density of nickel is 8.9 g.mL-1 • So the results were: Precise, but not accurate

26. Accuracy & Precision (continued) • Some error always exists in measurements • Skill of measurer • Conditions of measurements • Limitation of instruments

27. Percentage Error • Accuracy of an individual value (or average) can be compared to the correct/accepted value % Error = Experimental – Accepted x 100 Accepted

28. Percentage Error • What is the percentage error for a mass measurement of 17.7 g, given that the correct value is 21.2 g? • A volume is measured experimentally as 4.26 mL. What is the percentage error, given that the accepted value is 4.15 mL?

29. Taking Measurements in Chemistry Significant Figures

30. Exploring Uncertainty and PrecisionThe Paper Clip Activity • Measuring always involves some degree of estimation (i.e. uncertainty) Ruler #3 required the least amount of estimation because instrument had greater precision (more markings)

31. Significant Figures • Certain digits: digits that represent a marking on a scale or non-blinking number of a display • Uncertain (estimated) digits: digits that represents the space between the marks on a scale or the blinking number on a display • Sig Figs – all digits of certainty + 1 estimated

32. Sig Figs: Using the Pacific/Atlantic Rule • Step 1: Ask yourself: is the decimal point present or absent? • Step 2: Determine which way to start counting • If the decimal point is present, start counting from the LEFT • If the decimal point is absent, start counting from the RIGHT A T L A N T I C bsent P A C I F I C resent

33. Pacific/Atlantic Rule • Step 3: Start counting on Pacific or Atlantic side from the first NON-ZERO number. Count all numbers after the first non-zero number including zeros.

34. Pacific/Atlantic Rule • Examples: • 1234 = ________ sig figs • 1204 = ________ sig figs • 0.00234 = _______ sig figs • 1230 = ______ sig figs • 1234.0 = ______ sig figs 4 Absent  4 Absent  3 Present  3 Absent  5 Present 

35. Pacific/Atlantic Rule 3 certain digits – indicated by lines on measuring device ; 1 estimated digit - in between lines • Examples: • 1234 = ________ sig figs • 1204 = ________ sig figs • 0.00234 = _______ sig figs • 1230 = ______ sig figs • 1234.0 = ______ sig figs 4 4 3 certain ; 1 estimated 3 2 certain ; 1 estimated (zero’s are place holders) 3 2 certain ; 1 estimated (zero is a place holder) 5 4 certain ; 1 estimated

36. Using Sig. Figs. In Calculations • Addition/Subtraction Rule • Answer should contain least # of decimal places • Multiplication/Division Rule • Answer should contain least sig figs.

37. Do Now: Precision of Lab Instruments • Record the following quantities to the correct number of decimal places. ________ L ________ mL _______ oC • Convert your answer in A to milliliters: ________ mL • Add your answer from A & B. Record using correct sig. figs. ________ mL

38. Scientific Notation • Some numbers are very large or very small, so we need a short hand notation. Too large: 602,200,000,000,000,000,000,000 6.022 x 1023 Too small: 0.0000000000000000000000199 1.99 x 10-23

39. Scientific Notation N x 10n N is a number between 1 and 10 n is a positive or negative integer if n is a negative number, the full number is a small decimal if n is a positive number, the full number is a large number 3.69 x 10-4 ________________ 1.245 x 105 ________________

40. Taking Measurements in Chemistry According to the Scientific Method

41. The Scientific Method • Scientific Method: logical approach to solving problems by… • Observing & collecting data • Formulating hypotheses • Testing hypotheses • Formulating theories • Publishing results

42. Two Types of Measurements • Remember: observations about matter can be categorized in two groups: • Qualitative Data • Quantitative Data • Qualitative (think “quality”): observations using words • Quantitative (think “quantity”): observations using numbers and units.

43. Studying a System • System: specific portion of matter in a given region of space that has been selected for study • Microscopic or macroscopic level • Variable: any condition that changes during an experiment • Independent: value being manipulated • Dependent: result

44. Studying a System • Experimental Control: conditions that remain constant throughout (i.e. don’t change) • Often many controlled portions of system • Model: Explanation of how phenomena occur and how data or events are related • Visual • Verbal • Mathematical • Ex: atomic model of matter

45. Studying a System • Theory: broad generalization that explains a body of facts or phenomena • Used to predict results of new experiments • Ex: kinetic molecular theory

46. Taking Measurements in Chemistry Graphing Measurements

47. Independent Variable Dependent Variable Fertilizer Growth Direct Relationship • Title • Appropriate scale • Axis labeled • “Best fit” line

48. Direct Relationships When 2 quantities divided by each other gives a constant value K (constant value) = Y/X Ex: Density

49. InverseRelationships When 2 quantities multiplied by each other gives a constant value K = X Y Ex: Boyle’s Law K = PV