2. Bell Work Prepare for Quiz Hey Kim….

3. Write these in your Bell Work Composition Book 1.Kr 2.Ar 3.Neon 4.Hellium 5.Astatine 6.F 7.Chlorine 8.Bromine 9.Iodine 10.Xenon 11.Rn 12.Mn 13.Iron 14.Co 15.Nickel 16.Cu 17.Zinc 18.Ag 19.Au 20.Hg

4. Answers: 7A, 8A and transition elements 1.Kr - Krypton 2.Ar - Argon 3.Neon - Ne 4.Helium - He 5.Astatine - At 6.F - Fluorine 7.Chlorine - Cl 8.Bromine - Br 9.Iodine - I 10.Xenon - Xe 11.Rn - Radon 12.Mn - Manganese 13.Iron - Fe 14.Co - Cobalt 15.Nickel - Ni 16.Cu - Copper 17.Zinc - Zn 18.Ag - Silver 19.Au - Gold 20.Hg - Mercury

5. Today’s Element Zn Zinc Zn Periodic Table Placement Group 2B or 12 Period 4 Video: Is Zinc Boring?

6. Zinc Chemical PropertiesPhysical Properties Importance Periodic Table Atomic # Atomic mass # of Protons # of Electrons # of Neutrons Period Group

7. What’s in a Battery?Battery Modern batteries use a variety of chemicals to power their reactions. Common battery chemistries include: Zinc-carbon battery: The zinc-carbon chemistry is common in many inexpensive AAA, AA, C and D dry cell batteries. The anode is zinc, the cathode is manganese dioxide, and the electrolyte is ammonium chloride or zinc chloride.zinc Alkaline battery: This chemistry is also common in AA, C and D dry cell batteries. The cathode is composed of a manganese dioxide mixture, while the anode is a zinc powder. It gets its name from the potassium hydroxide electrolyte, which is an alkaline substance. Lithium-ion battery (rechargeable): Lithium chemistry is often used in high- performance devices, such as cell phones, digital cameras and even electric cars. A variety of substances are used in lithium batteries, but a common combination is a lithium cobalt oxide cathode and a carbon anode. Lead-acid battery (rechargeable): This is the chemistry used in a typical car battery. The electrodes are usually made of lead dioxide and metallic lead, while the electrolyte is a sulfuric acid solution.

8. Water: Separation by Electrolysis Video of Electrolysis: Water to Hydrogen and Oxygen

10. Atomic Mass of a Compound H 2 O Try these CO 2 C 6 H 12 O 6 1.01 + 1.01 + 16 = 18.02 H 2 1.01 x 2 = 2.02 O 16 x 1 = 16.00 Add totals 2.02 + 16.00 18.02

11. Practice – Finding Atomic Mass CO 2 C 12.01 x 1 = 12.01 O 2 16 x 2 = 32.00 Add totals 12.01 + 32.00 44.01

12. Practice – Finding Atomic Mass C 6 H 12 O 6 C 6 12.01 x 6 = 72.06 H 12 1.01 x 12 = 12.12 O 2 16 x 6 = 96.00 Add totals 72.06 12.12 +96.00 108.18

13. Percent Composition of Mass for Mixtures A 6g mixture of sulfur and iron is separated using a magnet. Data Sulfur (S) Iron (Fe) 5g 1g Calculate the percent composition of S and Fe.

14. Percent Composition of Mass for Mixtures A 6g mixture of sulfur and iron is separated using a magnet. Data Sulfur (S) Iron (Fe) 5g 1g Calculate the percent composition of S and Fe. Part / Whole x 100 = % composition Sulfur: 5g/6g x 100 = Iron : 1g/6g x 100 =

15. Percent Composition of Mass for Mixtures A 6g mixture of sulfur and iron is separated using a magnet. Data Sulfur (S) Iron (Fe) 5g 1g Calculate the percent composition of S and Fe. Part / Whole x 100 = % composition Sulfur: 5g/6g x 100 = 83.33% S Iron : 1g/6g x 100 = 16.66% Fe

16. Use Percent Composition to find the composition of a compound Use the periodic table to find the compound’s percent composition of each element. List the atomic weight of each element in the compound Note how many of each type of atom is in the compound Add it all up to get the atomic weight of the whole compound

17. Atomic Mass of a Compound H 2 O Try these CO 2 C 6 H 12 O 6 1.01 + 1.01 + 16 = 18.02 H 2 1.01 x 2 = 2.02 O 16 x 1 = 16.00 Add totals 2.02 + 16.00 18.02

18. Practice – Percent Composition H 2 O part / whole x 100 = % composition % composition of H % composition of O

19. Practice – Percent Composition CO 2 part / whole x 100 = % composition % composition of C % composition of O

20. Practice – Percent Composition C 6 H 12 O 6 part / whole x 100 = % composition % composition of C % composition of H % composition of O

21. Law of Conservation of Mass Mass is neither created nor destroyed in any process. It is conserved. Mass reactants = Mass products 2H 2 O + electricity yields 2H 2 + O 2

22. Isotopes The atomic weight found on the periodic table is based on the average weight of all the isotopes of the element Isotope – atoms of the same element with the same number of protons but different numbers of neutrons M&M activity

23. Writing Isotopes

24. Reading Isotopes Mass number - the sum of the protons and neutrons

25. Isotopes of Hydrogen

26. Write Isotopes for Iron

27. More isotopes Argon 36, Argon 37…

28. M&Mium Isotope Activity

29. http://www.chem.memphis.edu/bridson/Fund Chem/T07a1100.htm

30. S.I.Units http://2012books.lardbucket.org/books/gener al-chemistry-principles-patterns-and- applications-v1.0/section_05.html#averill_1.0- ch01_s09_s01_s02_t02 http://2012books.lardbucket.org/books/gener al-chemistry-principles-patterns-and- applications-v1.0/section_05.html#averill_1.0- ch01_s09_s01_s02_t02

31. http://chemwiki.ucdavis.edu/Physical_Chemist ry/Atomic_Theory/The_Mole_and_Avogadro's _Constant

32. MeasurementsandCalculations Where to Round Song

33. Steps in the Scientific Method 1.Observations 1.Observations - quantitative - quantitative - qualitative - qualitative 2.Formulating hypotheses 2.Formulating hypotheses - possible explanation for the observation - possible explanation for the observation 3.Performing experiments 3.Performing experiments - gathering new information to decide - gathering new information to decide whether the hypothesis is valid whether the hypothesis is valid

34. Outcomes Over the Long-Term Theory (Model) Theory (Model) - A set of tested hypotheses that give an - A set of tested hypotheses that give an overall explanation of some natural phenomenon. overall explanation of some natural phenomenon. Natural Law Natural Law - The same observation applies to many - The same observation applies to many different systems different systems - Example - Law of Conservation of Mass - Example - Law of Conservation of Mass

35. Law vs. Theory A law summarizes what happens  A law summarizes what happens  A theory (model) is an attempt to explain why it happens.

36. Nature of Measurement Part 1 - number Part 1 - number Part 2 - scale (unit) Part 2 - scale (unit) Examples: Examples: 20 grams 20 grams 6.63 x 10 -34 Joule seconds 6.63 x 10 -34 Joule seconds Measurement - quantitative observation consisting of 2 parts consisting of 2 parts

37. The Fundamental SI Units (le Système International, SI) International System of Units a system of measurement units agreed on by scientists to aid in the comparison of results worldwide.International System of Units

38. SI Units

39. Metric Prefixes Metric Prefixes Common to Chemistry PrefixUnit Abbr.Exponent Kilok10 3 Decid10 -1 Centic10 -2 Millim10 -3 Micro  10 -6 Nanon10 -9

40. Metric Prefixes and Conversion Examples

41. Uncertainty in Measurement A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

42. Why Is there Uncertainty?  Measurements are performed with instruments  No instrument can read to an infinite number of decimal places Which of these balances has the greatest uncertainty in measurement?

43. Precision and Accuracy Accuracy refers to the agreement of a particular value with the true value. Precision refers to the degree of agreement among several measurements made in the same manner. Precision refers to the degree of agreement among several measurements made in the same manner. Neither accurate nor precise Precise but not accurate Precise AND accurate

44. Types of Error Random Error (Indeterminate Error) - measurement has an equal probability of being high or low. Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration. Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration.

45. Rules for Counting Significant Figures - Details Nonzero integers always count as significant figures. 3456 has 3456 has 4 sig figs. 4 sig figs.

46. Rules for Counting Significant Figures - Details Zeros Zeros - Leading zeros do not count as - Leading zeros do not count as significant figures. 0.0486 has 0.0486 has 3 sig figs. 3 sig figs.

47. Rules for Counting Significant Figures - Details Zeros Zeros - Captive zeros always count as - Captive zeros always count as significant figures. 16.07 has 16.07 has 4 sig figs. 4 sig figs.

48. Rules for Counting Significant Figures - Details Zeros Zeros Trailing zeros are significant only if the number contains a decimal point. Trailing zeros are significant only if the number contains a decimal point. 9.300 has 9.300 has 4 sig figs. 4 sig figs.

49. Rules for Counting Significant Figures - Details Exact numbers have an infinite number of significant figures. 1 inch = 2.54 cm, exactly 1 inch = 2.54 cm, exactly

50. Sig Fig Practice #1 How many significant figures in each of the following? 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs 3.29 x 10 3 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000  2 sig figs

51. Rules for Significant Figures in Mathematical Operations Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in the calculation. 6.38 x 2.0 = 6.38 x 2.0 = 12.76  13 (2 sig figs) 12.76  13 (2 sig figs)

52. Sig Fig Practice #2 3.24 m x 7.0 m CalculationCalculator says:Answer 22.68 m 2 23 m 2 100.0 g ÷ 23.7 cm 3 4.219409283 g/cm 3 4.22 g/cm 3 0.02 cm x 2.371 cm 0.04742 cm 2 0.05 cm 2 710 m ÷ 3.0 s 236.6666667 m/s240 m/s 1818.2 lb x 3.23 ft5872.786 lb·ft 5870 lb·ft 1.030 g ÷ 2.87 mL 2.9561 g/mL2.96 g/mL

53. Rules for Significant Figures in Mathematical Operations Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement. 6.8 + 11.934 = 6.8 + 11.934 = 18.734  18.7 (3 sig figs) 18.734  18.7 (3 sig figs)

54. Sig Fig Practice #3 3.24 m + 7.0 m CalculationCalculator says:Answer 10.24 m 10.2 m 100.0 g - 23.73 g 76.27 g 76.3 g 0.02 cm + 2.371 cm 2.391 cm 2.39 cm 713.1 L - 3.872 L 709.228 L709.2 L 1818.2 lb + 3.37 lb1821.57 lb 1821.6 lb 2.030 mL - 1.870 mL 0.16 mL 0.160 mL

55. In science, we deal with some very LARGE numbers: 1 mole = 602000000000000000000000 In science, we deal with some very SMALL numbers: Mass of an electron = 0.000000000000000000000000000000091 kg Scientific Notation

56. Imagine the difficulty of calculating the mass of 1 mole of electrons! 0.000000000000000000000000000000091 kg x 602000000000000000000000 x 602000000000000000000000 ???????????????????????????????????

57. Scientific Notation: A method of representing very large or very small numbers in the form: M x 10n M x 10n  M is a number between 1 and 10  n is an integer

58. 2 500 000 000 Step #1: Insert an understood decimal point. Step #2: Decide where the decimal must end up so that one number is to its left up so that one number is to its left Step #3: Count how many places you bounce the decimal point the decimal point 1234567 8 9 Step #4: Re-write in the form M x 10 n

59. 2.5 x 10 9 The exponent is the number of places we moved the decimal.

60. 0.0000579 Step #2: Decide where the decimal must end up so that one number is to its left up so that one number is to its left Step #3: Count how many places you bounce the decimal point the decimal point Step #4: Re-write in the form M x 10 n 12345

61. 5.79 x 10 -5 The exponent is negative because the number we started with was less than 1.

62. PERFORMING CALCULATIONS IN SCIENTIFIC NOTATION ADDITION AND SUBTRACTION

63. Review: Scientific notation expresses a number in the form: M x 10 n 1  M  10 n is an integer

64. 4 x 10 6 + 3 x 10 6 IF the exponents are the same, we simply add or subtract the numbers in front and bring the exponent down unchanged. 7 x 10 6

65. 4 x 10 6 - 3 x 10 6 The same holds true for subtraction in scientific notation. 1 x 10 6

66. 4 x 10 6 + 3 x 10 5 If the exponents are NOT the same, we must move a decimal to make them the same.

67. 4.00 x 10 6 + 3.00 x 10 5 Student A 40.0 x 10 5 43.00 x 10 5  Is this good scientific notation? NO! = 4.300 x 10 6 To avoid this problem, move the decimal on the smaller number!

68. 4.00 x 10 6 + 3.00 x 10 5 Student B.30 x 10 6 4.30 x 10 6  Is this good scientific notation? YES!

69. A Problem for you… 2.37 x 10 -6 + 3.48 x 10 -4

70. 2.37 x 10 -6 + 3.48 x 10 -4 Solution… 002.37 x 10 -6 0.0237 x 10 -4 3.5037 x 10 -4

71. Direct Proportions  The quotient of two variables is a constant  As the value of one variable increases, the other must also increase  As the value of one variable decreases, the other must also decrease  The graph of a direct proportion is a straight line

72. Inverse Proportions  The product of two variables is a constant  As the value of one variable increases, the other must decrease  As the value of one variable decreases, the other must increase  The graph of an inverse proportion is a hyperbola