Chapter 1: Matter and Measurements

1. Chapter 1: Matter and Measurements

2. What is Chemistry? Biology vs. Chemistry vs. Physics

3. What is Chemistry? Biology Physics Chemistry The study of living organisms The study of forces & motion The study of matter and its reactions and properties

4. What is Chemistry? Chemistry is the study of CHANGES in the “stuff” around us. (We formally define “stuff” as matter!)

5. What is Chemistry? Think, pair, share: What are all the chemicals you use in your daily life?

6. Review: Scientific Methods 1. Hypothesis • Suggested solution to a problem 2. Experiment • A controlled method of testing a hypothesis 3. Data • Organized observations a. Data is always reproducible.

7. Review: Scientific Methods 4. Scientific Law • Statement which summarizes results of many observations and experiment a. Scientific laws explain WHAT is observed. • Example of a scientific law: 5. Scientific Theory • Explanation that supports a hypothesis and which has been supported with repeated testing b. Scientific theories explain WHY something is observed. • Example of a scientific theory:

8. Review: Scientific Methods 6. Steps of the Scientific Method—Review a. b. c. d. e. f.

9. What is Chemistry? Biology Physics Chemistry The study of living organisms The study of forces & motion The study of matter and its reactions and properties

10. What is Chemistry? Chemistry is the study of CHANGES in the “stuff” around us. (We formally define “stuff” as matter!)

11. Matter

12. Matter

13. Elements • Type of matter that cannot be broken down into simpler, stable substances and is made of only one type of atom

14. Compounds • A pure substance that contains two or more elements whose atoms are chemically bonded

15. Compounds • Fixed compositions • A given compound contains the same elements in the same percent by mass

16. Compounds • The properties of a compound are VERY DIFFERENT from the properties of the elements they contain Ex.) Sodium Chloride (NaCl) vs. Sodium & Chlorine Sodium: http://www.youtube.com/watch?v=RAFcZo8dTcU http://www.youtube.com/watch?v=92Mfric7JUc

17. Electrolysis

18. Mixtures • A blend of two or more kinds of matter, each of which retains its own identity and properties • Homogeneous • Heterogeneous

19. Homogeneous Mixtures • Composition is the same throughout the mixture • Examples: salt water, soda water, brass • A.k.a. a solution • Solute in a solvent (salt dissolved in water)

20. Heterogeneous Mixtures • Non-uniform; composition varies throughout the mixture

21. Separating Mixtures • Filtration

22. Separating Mixtures • Distillation

23. Separating Mixtures • Chromatography

24. Scientific Measurements Chemistry is a quantitative science. • This means that experiments and calculations almost always involve measured values. Scientific measurements are expressed in the SI (metric) system. • This is a decimal-based system in which all of the units of a particular quantity are related to each other by factors of ten.

25. SI System • Definition: modernized version of metric system; uses decimals • All units derived from base units; larger and smaller quantities use prefixes with base unit • Must memorize prefixes from nano- (10-9) to tera- (1012)

26. Prefixes (see handout & Ebook) • You will need to memorize all of the prefixes (factors, names and abbreviations from 109 (giga-) to 10-9 (nano)! • One example of a memory device:

27. INSTRUMENTS & UNITS Use SI units — based on the metric system Length Mass Time Temperature Meter, m Kilogram, kg Seconds, s Celsius degrees, ˚C Kelvins, K

28. Length The standard unit of length in the metric system is the METER which is a little larger than a YARD. USING THE PREFIXES WITH LENGTH: cm – often used in lab km – Gm –

29. Length Base unit: METER Conversions: 1 km=1000 m 1 cm = 10 -2 m 1 Gm = 106 m

30. O—H distance = 9.58 x 10-11 m 9.58 x 10-9 cm 0.0958 nm Units of Length • 1 kilometer (km) = 1000 meters (m) • 10-2 meter (m) = 1 centimeter (cm) • 102 meter (m) = 1 hectometer (Hm) • 1 nanometer (nm) = 1.0 x 10-9 meter

31. Volume THE COMMON UNITS OF VOLUME IN CHEMISTRY ARE: the liter (milliliter) and cubic centimeter (cm3) THE COMMON INSTRUMENTS FOR MEASURING VOLUME IN CHEMISTRY ARE: graduated cylinder & buret Note that 1 cm3 = 1 mL (We will use this exact conversion factor throughout the year, so you will need to memorize it!)

32. Mass • THE COMMON UNIT OF MASS IN CHEMISTRY IS : the gram (g) — often used in lab • Mass IS A MEASURE OF THE AMOUNT OF MATTER IN AN OBJECT; • Weight IS A MEASURE OF THE GRAVITATIONAL FORCE ACTING ON THE OBJECT. CHEMISTS OFTEN USE THESE TERMS INTERCHANGEABLY. 1000 g= 1 kg 1 Mg = 10 6 g

33. Temperature Scales Fahrenheit Celsius Kelvin Anders Celsius 1701-1744 Lord Kelvin (William Thomson) 1824-1907 TEMPERATURE IS THE FACTOR THAT DETERMINES the direction of heat flow.

34. Temperature Scales 212 ˚F 100 ˚C 373 K 100 K 180˚F 100˚C 32 ˚F 0 ˚C 273 K Fahrenheit Celsius Kelvin Boiling point of water Freezing point of water Notice that 1 kelvin degree = 1 degree Celsius

35. Temperature Scales 100 oF 38 oC 311 K oF oC K

36. SI System English Units (inches, feet, degrees F, etc.) are NEVER used to take measurements in the lab!

37. Calculations Using Temperature Fahrenheit/Celsius T (F) = 1.8 t (˚C) + 32

38. Calculations Using Temperature • Some calculations are in kelvins (especially important for Ch 5!!) • T (K) = t (˚C) + 273.15 (273) • Body temp = 37 ˚C + 273 = 310 K • Liquid nitrogen = -196 ˚C + 273 = 77 K

39. Problem Example 1L.1 A baby has a temperature of 39.8oC. Express this temperature in oF and K.

40. SI System: Base Units ESTABLISHMENT OF THE INTERNATIONAL SYSTEM OF UNITS (SI)— SI UNITS AS ESTABLISHED BY THE SI: LENGTH – meter (m) VOLUME – cubic meter (m3) MASS – kilogram (kg) TEMPERATURE – Kelvin (K)

41. Time Base unit: SECOND (sec) • Conversions: only non-decimal base unit 60 sec = 1 min 60 min = 1 hr

42. Precision and Accuracy in Measurements Precision vs. Accuracy Definitions: Precision—how close answers are to each other (reproducibility) Accuracy—how close answer is to accepted (true) value (agreement to accepted value)

43. Precision and Accuracy in Measurements Percent Error - a way to calculate accuracy in the lab Equation: % Error = | Accepted Value – Exp. Value | x 100 Accepted Value

44. Precision and Accuracy in Measurements Ex1.9 A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is 2.70 g/mL. What is the percent error for the student’s results? Equation: % Error = | Accepted Value – Exp. Value | x 100 Accepted Value

45. Scientific Notation Exponential (Scientific) Notation—See Worksheet

46. Significant Figures: Why are they Important? Numbers in math: no units, abstract, no context, can read calculator output exactly for answer. vs. Numbers in chemistry: measurements – include units. SIG FIGS WILL BE IMPORTANT THROUGHOUT THIS COURSE!

47. Graduated Cylinder Example http://learningchemistryeasily.blogspot.com/2013/07/precision-of-measurement-and.html

48. What are significant figures?(aka sig figs) • Significant figures are all the digits in a measurement that are known with certainty plus a last digit that must be estimated. • With experimental values your answer can have too few or too many sig figs, depending on how you round.

49. How Rounding Influences Sig Figs • 1.024 x 1.2 = 1.2288Too many numerals(sig figs) Too precise • 1.024 x 1.2 = 1Too few numerals(sig figs) Not precise enough

50. Why This Concept is Important • We will be adding, subtracting, multiplying and dividing numbers throughout this course. • You MUST learn how many sig figs to report each answer in or the answer is meaningless. • You must report answers on lab reports & tests/quizzes with the correct number of sig figs (+/- 1) or else you will lose points!!