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Introductory Chemistry:

Introductory Chemistry:

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Introductory Chemistry:

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  1. Introductory Chemistry: Chapter 1 Chemistry and You

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

  3. What is the definition of chemistry? • Study of all substances and the changes they undergo.

  4. 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

  5. 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

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

  7. 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

  8. 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 that don’t contain carbon 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 8

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

  10. 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

  11. 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)

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

  13. 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

  14. Prefix Multipliers in the SI System

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

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

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

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

  19. Accuracy vs. Precision

  20. 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 • even though every measurement has some random error, with enough measurements these errors should average out – Do multiple trials!

  21. 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 • systematic errors do not average out with repeated measurements because they consistently cause the measurement to be either too high or too low

  22. Volume • Derived unit • any length unit cubed • Measure of the amount of space occupied • SI unit = cubic meter (m3) • Commonly measure solid volume in cubic centimeters (cm3) • 1 m3 = 106 cm3 • 1 cm3 = 10-6 m3 = 0.000001 m3 • Commonly measure liquid or gas volume in milliliters (mL) • 1 L is slightly larger than 1 quart • 1 L = 1 dm3 = 1000 mL = 103 mL • 1 mL = 0.001 L = 10-3 L • 1 mL = 1 cm3

  23. 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?

  24. Mass vs. Volume of Brass

  25. 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 12.3 cm has 3 sig. figs. and its range is 12.2 to 12.4 cm 12.30 cm has 4 sig. figs. and its range is 12.29 to 12.31 cm

  26. 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 – Pacific Ocean side • 0.001050 has 4 sig. figs. • 1.050 x 10-3

  27. Counting Significant Figures • Trailing zeros may or may not be significant – Atlantic Ocean Side • Trailing zeros after a decimal point are significant • 1.050 has 4 sig. figs. • Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation • if 150 has 2 sig. figs. then 1.5 x 102 • but if 150 has 3 sig. figs. then 1.50 x 102

  28. 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.

  29. Multiplication and Division with Significant Figures • when multiplying or dividing measurements with significant figures, the result has the same number of significant figures as the measurement with the fewest number of significant figures 5.02 × 89,665 × 0.10 = 45.0118 = 45 3 sig. figs. 5 sig. figs. 2 sig. figs. 2 sig. figs. 5.892 ÷ 6.10 = 0.96590 = 0.966 4 sig. figs. 3 sig. figs. 3 sig. figs.

  30. Addition and Subtraction with Significant Figures • when adding or subtracting measurements with significant figures, the answer should reflect the largest uncertainty. 5.74 + 0.823+ 2.651 = 9.214 = 9.21 3 sf. 3 sf. 4 sf. 3 sf 4.8 - 3.965 = 0.835 = 0.8 2sf 4sf. 2sf

  31. 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 • add insignificant zeros to keep the value if necessary 5 to 9, round up • drop all digits after the last sig. fig. and increase the last sig. fig. by one • add insignificant zeros to keep the value if necessary to avoid accumulating extra error from rounding, round only at the end, keeping track of the last sig. fig. for intermediate calculations

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

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

  34. Rounding rounding to 2 significant figures • 234 rounds to 230 or 2.3 × 102 • 237 rounds to 240 or 2.4 × 102 • 234.9865 rounds to 230 or 2.3 × 102

  35. Both Multiplication/Division and Addition/Subtraction with Significant Figures • do whatever is in parentheses first, • First, evaluate the significant figures in the parentheses • Second, do the remaining steps 3.489 × (5.67 – 2.3) = 3 dp 2 dp 3.489 × 3.3 = 12 4 sf 2 sf 2 sf

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

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

  38. 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

  39. 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

  40. 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?

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

  42. Units • Always write every number with its associated unit • 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 • using units as a guide to problem solving is called dimensional analysis

  43. Problem Solving and Conversion Factors • Conversion factors are relationships between two units • May be exact or measured • Conversion factors generated from equivalence statements • e.g., 1 inch = 2.54 cm can give or METRIC TO ENGLISH CONVERSION ON PAGE 936 IN YOUR TEXT BOOK

  44. Dimensional Analysis • Using units as a guide to problem solving is called dimensional analysis • This is the technique that we have learned to convert between two different units.

  45. Problem Solving and Dimensional Analysis • Arrange conversion factors so given unit cancels • Arrange conversion factor so given unit is on the bottom of the conversion factor • May string conversion factors • So we do not need to know every relationship, as long as we can find something else the given and desired units are related to

  46. Practice – Convert 154.4 lbs to kg(1 kg = 2.20 lbs)

  47. Practice – Convert 30.0 mL to quarts(1 L = 1.057 qt)

  48. How many cubic centimeters are there in 2.11 yd3?

  49. yd3 in3 cm3 Practice 1.9 Convert 2.11 yd3 to cubic centimeters