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Introductory Chemistry:

This introductory chapter explores why chemistry is central to various human activities, different learning methods in chemistry, problem-solving techniques, lab safety symbols, and branches of chemistry like organic, inorganic, biochemistry, and more.

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Introductory Chemistry:

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  1. Introductory Chemistry: Chapter 1 Chemistry and You

  2. Tuesday, 9/9/14 • Learning Target: Students must be able to explain why chemistry is central to many human endeavors. Chapter 1

  3. Learning Chemistry • Different people learn chemistry differently. • What do you see in the picture? • Some people see a vase on a dark background, some people see two faces. Chapter 1

  4. Problem Solving • Connect the 9 dots using onlyfour straight lines. • Experiment until you find a solution. • However, we have used 5 straight lines. • No matter which dot we start with, we still need 5 lines. Chapter 1

  5. Problem Solving • Are we confining the problem? • We need to go beyond the 9 dots to answer the problem. Chapter 1

  6. Lab Safety SymbolsIdentify the following symbols A. B. C. D. E. F. G. H. I.

  7. What is the definition of chemistry? • The science that studies the composition of matter and its properties.

  8. Chemistry: The Central Science • Why???? • Most other sciences demand an understanding of basic chemical principles, and Chemistry is often referred to as the Central Science Chapter 1

  9. Modern Chemistry Chemistry is a science that studies the composition of matter and its properties. Chemistry is divided into several branches: Organic chemistry is the study of substances containing carbon Inorganic chemistry is the study of all other substances Biochemistry is the study of substances derived from plants and animals Analytical is the study of matter and ways to study the properties of matter. Physical is the physics of chemistry. Thermodynamics and quantum mechanics. Chapter 1 9

  10. Wednesday, 9/10/14 Learning Target: Students must know the metric system, SI units and derived units. Learning Outcome: Measurement Pre-Lab Chapter 1

  11. The Standard Units • Scientists have agreed on a set of international standard units for comparing all our measurements called the SI units

  12. Length • SI unit = meter • About a yard • Commonly use centimeters (cm) • 1 m = 100 cm • 1 cm = 0.01 m = 10 mm • 1 inch = 2.54 cm

  13. Mass • Measure of the amount of matter present in an object • weight measures the gravitational pull on an object, which depends on its mass • SI unit = kilogram (kg) • about 2 lbs. 3 oz. • Commonly measure mass in grams (g) or milligrams (mg)

  14. Time • measure of the duration of an event • SI units = second (s)

  15. Temperature Scales • Fahrenheit Scale, °F • used in the U.S. • Celsius Scale, °C • used in all other countries • Kelvin Scale, K • The SI unit for temperature

  16. Prefix Multipliers in the SI System

  17. What Is a Measurement? • quantitative observation • every measurement has a number and a unit • every digit written is certain, the last one which is estimated

  18. Estimation in Weighing • What is the uncertainty in this reading?

  19. Thursday, 9/11/14 Learning Target: Students must be able to compare and contrast accuracy and precision in measurement. Learning Outcome: Complete “Measurement Lab” Chapter 1

  20. Uncertainty in Measured Numbers uncertainty comes from: • limitations of the instruments used for comparison, • the experimental design, • the experimenter, • nature’s random behavior

  21. Precision and Accuracy • accuracy is an indication of how close a measurement comes to the actual value of the quantity Percent error = • precision is an indication of how reproducible a measurement is

  22. Accuracy vs. Precision

  23. Precision • imprecision in measurements is caused by random errors • errors that result from random fluctuations • we determine the precision of a set of measurements by evaluating how far they are from the actual value and each other called standard deviation. • Do multiple trials to lesson error and improve precision.

  24. Accuracy • inaccuracy in measurement caused by systematic errors • errors caused by limitations in the instruments or techniques or experimental design • we determine the accuracy of a measurement by evaluating how far it is from the actual value • Use percent error to calculate how accurate you are

  25. Mass & Volume • mass and volume are extensive properties • the value depends on the quantity of matter • extensive properties cannot be used to identify what type of matter something is • if you are given a large glass containing 100 g of a clear, colorless liquid and a small glass containing 25 g of a clear, colorless liquid - are both liquids the same stuff?

  26. Mass vs. Volume of Brass

  27. Monday 9/15/14 Learning Target: Know how to use significant figures in labs and in problems. Learning Outcome: Complete significant figures problems. Chapter 1

  28. Accuracy versus Precision Chapter 1

  29. Significant Figures • the non-place-holding digits in a reported measurement are called significant figures • significant figures tell us the range of values to expect for repeated measurements • We use significant figures in science because measurement is always involved.

  30. Counting Significant Figures • All non-zero digits are significant • 1.5 has 2 sig. figs. • Interior zeros are significant • 1.05 has 3 sig. figs. • Leading zeros are NOT significant –0.001050 has 4 sig. figs.

  31. Counting Significant Figures • Trailing zeros may or may not be significant 1) If a decimal is present, trailing zeros are significant • 1.050 has 4 sig. figs. 2) If a decimal is NOT present, trailing zeros are NOT significant. • if 150 has 2 sig. figs. then 1.5 x 102 • but if 150. has 3 sig. figs. then 1.50 x 102 **These are considered ambiguous and should be avoided by using scientific notation

  32. Determining the Number of Significant Figures in a Number How many significant figures are in each of the following? 0.04450 m 5.0003 km 1.000 × 105 s 0.00002 mm 10,000 m 4 sig. figs.; the digits 4 and 5, and the trailing 0 5 sig. figs.; the digits 5 and 3, and the interior 0’s 4 sig. figs.; the digit 1, and the trailing 0’s 1 sig. figs.; the digit 2, not the leading 0’s Ambiguous, generally assume 1 sig. fig.

  33. Multiplication and Division with Significant Figures • when multiplying or dividing measurements with significant figures, the answer must reflect the fewest number of significant figures 1) 5.02 × 89,665 × 0.10 = 2) 5.892 ÷ 6.10 =

  34. Addition and Subtraction with Significant Figures • when adding or subtracting measurements with significant figures, the answer should reflect the largest uncertainty 1) 5.74 + 0.823+ 2.651 = 2) 4.8 - 3.965 =

  35. Rounding if the number after the place of the last significant figure is: 0 to 4, round down • drop all digits after the last sig. fig. and leave the last sig. fig. alone 5 to 9, round up • drop all digits after the last sig. fig. and increase the last sig. fig. by one To avoid accumulating extra error from rounding, round only at the end, keeping track of the last sig. fig. for intermediate calculations

  36. Rounding rounding to 2 significant figures • 2.34 rounds to 2.3 • 2.37 rounds to 2.4 • 2.349865 rounds to 2.3

  37. Rounding rounding to 2 significant figures • 0.0234 rounds to 0.023 • 0.0237 rounds to 0.024 • 0.02349865 rounds to 0.023

  38. Rounding rounding to 2 significant figures • 234 rounds to 230 • 237 rounds to 240 • 234.9865 rounds to 230

  39. Both Multiplication/Division and Addition/Subtraction with Significant Figures • First, evaluate the significant figures in the parentheses • Second, do the remaining steps 3.489 × (5.67 – 2.3) =

  40. Perform the following calculations to the correct number of significant figures b)

  41. Example 1.6 Perform the following calculations to the correct number of significant figures b)

  42. Tuesday 9/16/14 Learning Target: Know how to use and convert numbers into scientific notation. Learning Outcome: I will be able to use scientific notation in problems and convert standard notation into scientific notation. Chapter 1

  43. Why are significant figures not important in your math class? Chapter 1

  44. Density • Ratio of mass:volume • Solids = g/cm3 • 1 cm3 = 1 mL • Liquids = g/mL • Gases = g/L • Volume of a solid can be determined by water displacement – Archimedes Principle

  45. Density • Density : solids > liquids >>> gases • except ice is less dense than liquid water! • Heating an object generally causes it to expand, therefore the density changes with temperature

  46. Density • Iron has a density of 7.86 g/cm3. Could a block of metal with a mass of 18.2 g and a volume of 2.56 cm3be iron?

  47. Density • What volume would a 0.871 g sample of air occupy if the density of air is 1.29 g/L?

  48. Wednesday, 9/17/14 Learning Target: Be able to apply dimensional analysis to convert from one unit of measure to another. Learning Outcome: I will be able to complete single-step unit conversion problems. Chapter 1

  49. Units • Always include units in your calculations • you can do the same kind of operations on units as you can with numbers • cm × cm = cm2 • cm + cm = cm • cm ÷ cm = 1

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