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Unit 1: Introduction to Chemistry

Unit 1: Introduction to Chemistry. Pre- AP Chemistry Edmond NorthHigh School Chapters: 1 & 2. Scientific Method / Process. The Scientific Method is a systematic approach to problem solving. It is generally composed of the following parts: Question Hypothesis Experiment

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Unit 1: Introduction to Chemistry

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  1. Unit 1: Introduction to Chemistry Pre- AP Chemistry Edmond NorthHighSchool Chapters: 1 & 2

  2. Scientific Method / Process • The Scientific Method is a systematic approach to problem solving. • It is generally composed of the following parts: • Question • Hypothesis • Experiment • Data Analysis • Conclusion

  3. Scientific Process

  4. Measurement • Measurement: A quantitative observation consisting of a numeric value and units. • Three are 2 kinds of units: base and derived • Base Units: Units are mutually independent of each other • Ex: 7 meters, 30 kg, 6 seconds • Derived Units: Units are obtained mathematically from base units • Ex: 10 cm3, 0.9 g/mL, 22 meters/second

  5. Metric System • The Metric System is a decimalized system of measurement based on powers of 10 • Used internationally and in the scientific community • Consists of base units and prefixes The United States, Liberia, and Myanmar (Burma) are the only three countries who do not use the metric system

  6. Metric System - Base Units Volume liters (L) Mass grams (g) Distance meters (m) Temperature Time Kelvin (K) seconds (s)

  7. Accuracy & Precision Accuracy: the degree of agreement between the true value and the measured value (bullseye) Precision: the degree of agreement among several measured values (grouping)

  8. Accuracy & Precision Three different groups of students measure the mass of a medal, with a known value of 5.000 grams. Evaluate each group’s data for its accuracy and precision (low or high): high low low high low high Group 1 Group 2 Group 3 Accuracy ______ Precision ______ Accuracy ______ Precision ______ Accuracy ______ Precision ______

  9. Uncertainty in Measurement All measuring instruments have a degree of uncertainty. The more divisions a device has, the more accurate the measurement. The last digit of a measurement is always estimated (uncertain). How would you read the volume in this graduated cylinder if the markings showed every 5 mL instead of every 1 mL? read from the bottom of the meniscus

  10. Uncertainty in Measurement If we measure the length of the paw print with a decimeter ruler, we know for a certainty that it is between 0 and 1 decimeters. We estimate the next digit: 0.3 decimeters estimated digit

  11. Uncertainty in Measurement If we measure the length of the paw print with a centimeter ruler, we know for a certainty that it is between 3 and 4 centimeters. We estimate the next digit: 3.5 centimeters estimated digit

  12. Uncertainty in Measurement If we measure the length of the paw print with a millimeter ruler, we know for a certainty that it is between 34 and 35 millimeters. We estimate the next digit: 34.5 millimeters estimated digit

  13. Significant Figures Significant Figures are the digits in a measurement that are known with some degree of certainty are called significant figures. The number of significant figures in a measurement = the number of digits that are known + the estimated digit The more significant figures after the decimal, the more accurate the measurement. Ex: This triple beam balance shows a mass of 62.41 grams. There are 4 significant figures in this measurement.

  14. Significant Figure Rules There are rules for counting significant figures related to: Non-zero integers Zeroes Leading zeroes Captive zeroes Trailing zeroes

  15. Significant Figure Rules Non-zero integers are always significant 3456 cm has 4sig figs

  16. Significant Figure Rules Leading zeroes are never significant 0.0486g has 3 sig figs

  17. Significant Figure Rules Captive zeroes are always significant 16.07 mLhas 4sig figs

  18. Significant Figure Rules Trailing zeroes are only significant if the number contains a decimal 9.30 m has 3 sig figs 930 m has 2 sig figs

  19. Significant Figures Practice How many significant figures are in the following measurements? 8,675,309 km ____ 90,210 L ____ 0.07 mg ____ 0.2020 daL ____ 300.00 g ____ 7 4 1 4 5

  20. Significant Figure Operations Multiplication & Division Calculate “raw” answer Rounded answer must contain no more significant figures than the measurement with the least number of significant figures Example: What is the density of a bar of gold with a mass of 87.82 g and a volume of 4.55 cm3? (Density = mass  volume) 87.82 g  4.55 cm3 = 19.301099 g/cm3 19.3 g/cm3 4 sig figs 3 sig figs raw answer 3 sig figs

  21. Significant Figure Operations Addition & Subtraction Calculate “raw” answer Rounded answer can have no more digits to the right of the decimal point than the measurement with the least number of digits to the right of the decimal point Example: What is the difference in length between a professional shot put throw of 23.125 meters and an amateur shot put throw of 21.2 meters? 23.125 m  21.2 m = 1.925 m  1.9 m 3 digits 1 digit raw 1 digit after decimal after decimal answer after decimal

  22. Exact Numbers Exact Numbers are different from measurements because they have no uncertainty Significant figures does not apply Examples Conversions (3 feet = 1 yard) Counting Numbers (20 M&Ms)

  23. Metric System - Base Units Volume liters (L) Mass grams (g) Distance meters (m) Temperature Time Kelvin (K) seconds (s)

  24. Metric System - Prefixes

  25. Powers of Ten

  26. From Quarks to Outer Space

  27. Metric Conversions Example: How many kilograms are in 75 decigrams? Example: How centiliters are in 12 dekaliters? k h da d c m  base x x x x k h da d c m  base Move decimal 4 places to the left: 75 dg = 0.0075 kg (start) Move decimal 3 places to the right: 12 daL = 12,000 cL (start)

  28. Metric Conversions Practice Practice metric conversions: _______ dg = 4.2 hg _______ dam = 6,055 mm _______ L = 1 cL _______ cK = 0.003 kK _______ s = 11,700,000 s k h da d c m  base x x 4,200 0.6055 0.01 300 11.7

  29. Remember King Henry died by drinking chocolate milk. Kilo HectoDekaBase Deci, Centi, and Milli.

  30. Metric System: Pressure • Pressure is the force over a given area • If someone stepped on your foot, which shoe would you prefer they wore? • Pressure is measured in units of: • 1 Atmospheres (atm) • 101.3 kiloPascals (N/m2) • 760 mmHg (mm of Mercury) • 760 torrs • Pressure is measured by 2 instruments: • Barometer • Manometer

  31. Pressure Conversions 400 torr = __________ kPa 400 torr 101.3kPa = 53.32kPa 760 torr  328 mmHg = __________ atm 328mmHg 1atm = 0.432atm 760 mmHg  1 atm = 760 torr = 760 mmHg = 101.3 kPa

  32. Density • Density is the ratio of an object’s mass and volume • The formula for density is D = m / v • In chemistry, the 2 most common units of density will be g/mL and g/cm3 The four cubes to the right have the same volume (1 cm3), but different masses. How does this effect their densities?

  33. Density • The density of an object is an intensive physical property, meaning it cannot be changed no matter the quantity • A property that changes with the amount or quantity of the substance is an extensive property. • The density of water is 1 g/mL • Objects that float < 1 g/mL • Objects that sink > 1 g/mL • What happens to the density of an object when it is sawed I half?

  34. Counting Matter • What are some common ways we count matter? • Dozen = 12 • Ream = 500 • Gross = 144 • Mole = 6.02 x 1023

  35. The Mole • In the same way a dozen is worth 12, a mole is worth 6.02 x 1023 • This number is called Avogadro’s number • Mole is abbreviated as “mol” • Written in expanded form, that number is: 602,000,000,000,000,000,000,000 • The mole is a large number because particles are so small, it takes many of them make up an amount we can see and understand • We can use it to count anything!

  36. The Mole • 1 dozen cookies = 12 cookies • 1 mole of cookies = 6.02 X 1023 cookies • 1 gross cars = 144 cars • 1 mole of cars = 6.02 X 1023 cars • 1 ream Al particles = 500 Al particles • 1 mole of Al particles = 6.02 X 1023 particles

  37. Mole Calculations 6.02 x 1023 particles = 1 mole Ex: How many particles are in 3.00 moles of N2? 3.00 mol N2 6.02 x 1023 particles = 1 mol N2 Ex: How many moles of Na are in 1.10 x 1023 particles? 1.10 x 1023 particles Na 1 mol Na = 6.02 x 1023 particles Na 1.81 x 1024 particles 1.83 x 1022 moles

  38. 4.56 x 10-7 Exponent Coefficient Base Scientific Notation • Scientific Notation is mathematical shorthand that makes large and small numbers manageable • It is composed of three parts: • Coefficient • Base • Exponent

  39. Scientific Notation Rules • Coefficients • Coefficients must be greater than or equal to 1 and less than 10 • Coefficients can be positive or negative • All numbers in the coefficient are counted as significant • Which of the following numbers are written incorrectly? 22 x 105 9.5 x 102 10 x 108 7 x 10-3 0.3 x 10-9 -1.00 x 106

  40. Scientific Notation Rules • Bases • Base is always a 10 • Bases are never counted as significant • Exponents • Exponents are always integers • Exponents can be positive (big number) or negative (small number) • Ex: 1 x 103 = 1000 and 1 x 10-3 = 0.001 • Exponents are never counted as significant • Which of the following numbers are written incorrectly? 4.1 x 10-5 2.2 x 620 1.0 x 10-1 7 x 100 0.15 x 10-9 -7 x 106.3

  41. Scientific Notation • How to enter the number 2.5 x 10-8 into the calculator: • Enter the coefficient 2.5 • Press 2nd, then EE • Enter the exponent -8 It should appear on your screen as 2.5E-8

  42. Two Unit Calculations 40 Rods = X in 1 Hogshead Gallon

  43. 1st Length Conversion • 40 rods = 1 furlong • 1 furlong = 10 chains • 1 chain = 66 feet

  44. 2nd Volume Conversion • 2 Mouthfuls = 1 Jigger • 2 Jiggers = 1 Jack • 2 Jacks = 1 Gill • 2 Gills = 1 Cup • 2 Cups = 1 Pint • 2 Pints = 1 Quart • 2 Quarts = 1 Pottle • 2 Pottles = 1 Gallon • 2 Gallons = 1 Peck • 2 Pecks = 1 Doublepeck • 2 Doublepecks = 1 Bushel • 2 Bushels = 1 Cask • 2 Casks = 1 Barrel • 2 Barrels = 1 Hogshead • 2 Hogsheads = 1 Pipe • 2 Pipes = 1 Tun

  45. 40 Rods = 7920 inches 1 Hogshead = 64 gallons

  46. 123.75 inches gallon

  47. 2 Unit Conversions • Convert the density of titanium (4.54g/mL) to kg/L. • Convert 13.2 mg/mL to g/cm3.

  48. Classifying Matter

  49. carbon (C) sulfur (S) copper (Cu) mercury (Hg) sugar (C6H12O6) salt (NaCl) water (H2O) rust (Fe2O3) Pure Substances • Pure substances cannot be separated by physical means • Elements: cannot be chemically separated, listed on the periodic table • Compounds: can be chemically separated, made up of elements

  50. Pure Substances • Particle representations of… • Elements • Compounds

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