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Introduction to Chemistry: Matter, Changes, and Scientific Method

This chapter provides an introduction to the study of chemistry, covering topics such as matter, changes it undergoes, scientific method, branches of chemistry, and units of measurement. It also explores topics like density, temperature, scientific notation, and conversion factors.

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Introduction to Chemistry: Matter, Changes, and Scientific Method

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  1. Chapter 1 and 2 Introduction to Chemistry

  2. Chemistry • Chemistry is the study of matter and the changes it undergoes • Matter is anything that has mass AND takes up space. • Mass measures the amount of matter an object takes up.

  3. Mass vs. Weight • Which one depends on the gravitational pull on matter? • Weight • Which one is independent of gravitational pull? • Mass *Which would scientists all around the world prefer?

  4. Chemistry • We see all matter on a macroscopic level. Chemistry tries to explain how matter behaves on a submicroscopic level.

  5. 5 branches of Chemistry • Organic Chemistry-study of carbon containing cmpds. • Inorganic Chemistry-matter without carbon • Physical Chemistry-energy changes of matter • Biochemistry-matter of living things • Analytical Chemistry-identification of matter

  6. Scientific Method • Scientific method-systematic approach to study the sciences. • Why is it needed?

  7. 5 Steps of the Scientific Method 1. Observation-gathering information • Quantitative data-numerical info ex. 4.56 g of sugar • Qualitative data-info using 5 senses ex. white crystals

  8. Scientific Method 2. Hypothesis-tentative explanation of observations. 3. Experiments-a set of controlled observations testing hypothesis. • Independent variable-the variable you plan to change • Dependent variable-the variable that depends on the independent variable

  9. Variables Ex. Comparing the time you study to the grades you make in chemistry. • The independent variable is the time you study. • The dependent variable is the grades. Ex. Comparing the rate of dissolving to the size of substance. • The independent variable is the size of the subst. • The dependent variable is the rate of dissolving.

  10. Scientific Method 4. Data Analysis-usually with graphs 5. Conclusion-judgment based on the info obtained. • Theory-an explanation supported by repeated experiments • Ex. Atomic theory that explains how the atom behaves • Law-statement of truth; no exceptions • Ex. Law of Gravity: What goes up, must come down.

  11. Scientific Research • Pure research-for the sake of gaining knowledge. • Ex. What are protons made of? • Applied research-to solve a problem. • Ex. Determine how to cure those with SARS. • Technology-practical use of scientific knowledge. • Ex. Catalytic converters developed for cars to remove air pollutants

  12. Units of Measurements • The international system developed by the French in 1960 called the SI base units is used so scientists can communicate around the world.

  13. SI Base Units (Memorize) QuantityBase unitAbbrev Time Seconds s Length Meter m Mass Kilogram kg Temperature Kelvin K Amount of substance Mole mol

  14. SI Prefixes (Memorize) PrefixSymbolFactor Mega M 1 000 000 Kilo k 1 000 Deci d 1/10 Centi c 1/100 Milli m 1/1000 Micro u 1/1000 000

  15. Units Derived Units – combination of base units Ex. Volume = l x w x h = m3 Density = mass/volume = g/cm3 Volume: SI unit is m3. Also used is mL.

  16. Density • Density = Mass D = M Units: g/cm3 Volume V g/ mL Ex. Suppose a sample of aluminum is placed in a 25 mL graduated cylinder containing 10.5 mL of water. The level of water rises to 13.5 mL. What is the mass of the aluminum sample, if the density is 1.09 g/mL ?

  17. Density Ex. What is the volume of an object that has a mass of 3.43 grams and a density of 1.45 g/mL?

  18. Temperature Scientists use 2 scales: degree Celsius AND Kelvin K = 273 + ºC Ex. Convert the following temp to K. a. 20 ºC b. 100 ºC c. 39 ºC Ex. Convert the following temp to ºC. a. 332K b. 264K c. 273K

  19. Scientific Notation • Scientific notation-expressing a number in this form: _._ _ _ x 10x • When the number is larger than one, the exponent is positive. • When the number is smaller than one, the exponent is negative. • Ex. 1 392 000 km = 0.000 000 028 g/cm3 =

  20. Using the scientific calculator • Enter 1392000, 2nd , sci Display: 1.392 06 On some TI calculators, it is the 3rd function. • For graphing calculators: Go into sci mode first (mode, sci) Display: 1.392 E6

  21. Addition and Subtraction of Exponential Numbers • You can add or subtract numbers by hand only if they have the same power of 10. Example: 7.11 X 104 + 4.0 X 103 = Calculator: enter: 7.11 EE 4 + 4.0 EE 3 = 2nd SCI

  22. Multiplication of Exp. Numbers • 10A X 10B = 10A+B • Ex. • 102 X 105 = • (4.0 X 103) (2.0 X 10-5) = • By calculator: enter: 4.0 EE 3 X 2.0 EE +/- 5 = 2nd SCI

  23. Division of Exponential Numbers • 10A = 10A-B107 = 107-4 10B 104 Ex. 1.06 X 106 = 2.0 X 101 Calculator: 1.06 EE 4 divide 2.0 EE 1 = 2nd SCI

  24. Conversion Factors • Conversion Factors-a ratio of equivalents of different units Ex. 3 teaspoons = 1 tablespoon is a conversion factor in these two forms: 3 teaspoons or 1 tablespoon 1 tablespoon 3 teaspoons

  25. Conversion Factor • 1 km = 1000 m is written as a conversion factor as: 1 km or 1000 m 1000 m 1 km *** The larger unit gets the 1 ***

  26. Writing Conversion Factors Practice: Write a conversion factor for the following: a. ms, s b. g, kg c. dm, m d. g, mg e. cm, m

  27. Dimensional Analysis • Dimensional Analysis- a method of problem solving that focuses on the units used to describe matter.

  28. Steps to Solving Dimensional Analysis Problems: 1. Write given on left. 2. Set conversion factor so given unit will cancel out. 3. Give answer with units.

  29. Practice Examples: 1. Convert 48 km to m. 2. Convert 2500 mL to L. 3. Convert 73 mg to g. 4. Convert 2.3 Mg to g. 5. Convert 563 dL to L.

  30. Practice 6. A football player’s weight is 245 lb. Convert to grams. 1 lb = .454 kg 7. Convert 4.5 years to seconds. 8. A man has 141 tuz and wants to trade them in for vums. How many will he get? 1 fot = 5 vum 4 bef=3 tuz 9 fot= 2 bef

  31. 2 Step Dimensional Analysis • Practice: 1. 3256 mL to kL 2. 9.3 km to cm 3. 0.04 Mg to dg 4. 53 674 um to cm

  32. Accuracy vs. Precision • Accuracy- how close a measured value is to an accepted value • Precision-how close a series of measurements are to each other • Illustration:

  33. Accepted value: 1.59 g/cm3 Questions: 1. Who collected the most accurate density? 2. Who collected the most precise density? Density Data Table (g/cm3)

  34. Percent Error • Percent Error = Accepted –Experimental * 100 Accepted Calculate the percent error for Sue on Trial 1,2,3.

  35. Significant Figures Scientists indicate the precision of measurements by the number of digits that they report. Significant figures indicate the precision of the measuring instruments. Today better measuring devices show more precise measurements. *Sig. fig. are numbers that are not estimated.*

  36. Precision Ex. Two timing instruments recorded the following: 36 s and 35.95 s Which instrument is more precise?

  37. Significant Figures • Significant Figures-include all known digits and one estimated digit

  38. Rules for Significant Figures 1. Non-zero numbers are always significant. Ex. 2. Zeros between non-zero numbers are always significant. Ex. 3. All final zeros to the right of the decimal place are significant. Ex. 4. Zeros that act as place holder are not significant. Convert quantities to scientific notation to remove placeholder zeros. Ex.

  39. Rounding Off Numbers Use math rules: if number to be considered is 5 or greater, round up. if number to be considered is less than 5, keep the same. Ex. Round this number to 5 sign. fig., then to 3 s.f. and finally to one s.f. 3.515014

  40. *Rounding • Round the following to 4 significant figures. 1. 84791 kg 2. 38.5432 3. 256.75 cm 4. 0.000548 18 g 5. 308 659 000 mm 6. 136 758 g

  41. Rounding 1. Which digit in 45.67 is the estimated digit? Round the following to 3 significant figures: 2. 0.0023665 3. 362 230 4. 0.002365 5. 123 236 6. 2.0236

  42. Addition and Subtraction of S.F. Keep the fewest number of decimal places that is in any of the entries. Ex. 28.0 cm + 23.538 cm + 25.68 cm = 77.218 cm • Round answer to one decimal place because the first entry has only one decimal place.

  43. Practice Round the following to the correct number of significant figures: 1. 0.0487 mg + 0.05834 mg + 0.00483 mg = 2. 6.23 cm + 5.2 cm = 3. 124.36 kg – 100 kg =

  44. Multiplication and Division of Significant Figures Keep the fewest significant figures that is any of the entries. Ex. 3.20 cm X 3.651 cm X 2.2 cm = 25.70304 cm3 • Round answer to two significant figures because the last entry has the fewest s.f.

  45. Practice Round the following to the correct number of significant figures: 1. 3.65 cm X 3.20 cm X 2.05 cm = 2. 4.62 m X 5.2 m = 3. 5.23 g/ 1.3 mL = 4. 620 m X 1.24 m

  46. Graphing Data • Graphing Data-In chemistry, most graphs will be line graphs. Review: 1. Determine which is the independent variable and which is the dependent variable. Independent variable goes on the x axis. Dependent variable goes on the y axis.

  47. Graphing 2. Label what is measured, units, and set up scale. 3. Plot data 4. Draw a best fit line = the line drawn has as many points above the line as below line. 5. If a straight line = variables are directly proportional = both variables increase or decrease at the same time. *Variables are inversely proportional if one variable increases and another decreases.

  48. Graphing 6. Determine the slope if it is a straight line. Slope = Rise Run 7. Extrapolation-extend line beyond plotted points to get more information. Interpolation-read data between measured values.

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