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Chemistry: The Study of Change

Explore the fundamentals of chemistry, including matter, mixtures, elements, compounds, states of matter, and scientific notation in this introductory chapter.

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Chemistry: The Study of Change

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  1. Welcome to CHM 1045 Chemistry: The Study of Change Chapter 1

  2. Assignment Read pages 1-10 of chapter 1. Homework: pages 23 - 26 2, 3, 4, 7, 8, 10, 14, 15, 17, 18, 22, 24, 26, 28, 30, 32, 34, 36, 40, 44, 48, 50, 60, 62 Note: problems with red numbers are answered at the back of the text.

  3. Chemistry is the study of matter and the changes it undergoes • Matter is anything that occupies space and has mass. • A (pure) substance is a form of matter that has a definite composition and distinct properties. water, ammonia, sucrose, gold, oxygen 1.4

  4. soft drink, milk, solder Oil and water, wood, iron filings in sand A mixture is a combination of two or more substances in which the substances retain their distinct identities. • Homogenous mixture – composition of the mixture is the same throughout. • Heterogeneous mixture – composition is not uniform throughout. 1.4

  5. magnet distillation Physical means can be used to separate a mixture into its pure components. 1.4

  6. An element is a substance that cannot be separated into simpler substances by chemicalmeans. • 115 elements have been identified • 83 elements occur naturally on Earth • gold, aluminum, lead, oxygen, carbon • 32 elements have been created by scientists • technetium, americium, seaborgium 1.4

  7. Fig. 1.4a,b Earth Earth’s crust

  8. Water (H2O) Glucose (C6H12O6) Ammonia (NH3) A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions. Compounds can only be separated into their pure components (elements) by chemical means. 1.4

  9. 1.4

  10. Three States of Matter 1.5

  11. sugar dissolving in water ice melting hydrogen gas burns in oxygen gas to form water Physical or Chemical? A physical change does not alter the composition or identity of a substance. A chemical change alters the composition or identity of the substance(s) involved. 1.6

  12. TA p9

  13. A 1 kg bar will weigh 2.2 lb on earth 0.4 lb on moon Matter - anything that occupies space and has mass. • mass – measure of the quantity of matter • SI unit of mass is the kilogram (kg) • 1 kg = 1000 g = 1 x 103 g weight – force that gravity exerts on an object 1.7

  14. Measurements • All measured quantities make known three pieces of information. • The quantity or number • The unit • The uncertainty in the measurement.

  15. 1.7

  16. Derived SI Units Quantity Definition of Quantity SI unit Area Length squared m2 Volume Length cubed m3 Density Mass per unit volume kg/m3 Speed Distance traveled per unit time m/s Acceleration Speed changed per unit time m/s2 Force Mass times acceleration of object kg * m/s2 ( =newton, N) Pressure Force per unit area kg/(ms2) ( = pascal, Pa) Energy Force times distance traveled kg * m2/s2 ( = joule, J)

  17. 1.7

  18. Volume – SI derived unit for volume is cubic meter (m3) 1 L = 1 dm3 (Definition) 1 mL = 1 x 10-3 L or 1L = 1 x 103 mL 1 dm3 = (10 cm)3 = 1000 cm3 1 L = 1000 mL = 1000 cm3 = 1 dm3 1 mL = 1 cm3 1.7

  19. mass density = volume A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass? m m d = d = V V Density – SI derived unit for density is kg/m3 1 g/cm3 = 1 g/mL = 1000 kg/m3 = 21.5 g/cm3 x 4.49 cm3 = 96.5 g m = d x V 1.7

  20. 0F = x 0C + 32 9 5 K = 0C + 273.15 273 K = 0 0C 373 K = 100 0C 32 0F = 0 0C 212 0F = 100 0C 1.7

  21. 0F – 32 = x 0C 0F = x 0C + 32 x (0F – 32) = 0C 5 5 5 0C = x (0F – 32) 0C = x (172.9 – 32) = 78.3 9 9 9 9 9 5 5 Convert 172.9 0F to degrees Celsius. 1.7

  22. The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 The mass of a single carbon atom in grams: 0.0000000000000000000000199 Scientific Notation 6.022 x 1023 1.99 x 10-23 N x 10n N is a number between 1 and 10 n is a positive or negative integer 1.8

  23. move decimal left move decimal right Scientific Notation 568.762 0.00000772 n > 0 n < 0 568.762 = 5.68762 x 102 0.00000772 = 7.72 x 10-6 Addition or Subtraction • Write each quantity with the same exponent n • Combine N1 and N2 • The exponent, n, remains the same 4.31 x 104 + 3.9 x 103 = 4.31 x 104 + 0.39 x 104 = 4.70 x 104 1.8

  24. Scientific Notation Multiplication (4.0 x 10-5) x (7.0 x 103) = (4.0 x 7.0) x (10-5+3) = 28 x 10-2 = 2.8 x 10-1 • Multiply N1 and N2 • Add exponents n1and n2 Division 8.5 x 104÷ 5.0 x 109 = (8.5 ÷ 5.0) x 104-9 = 1.7 x 10-5 • Divide N1 and N2 • Subtract exponents n1and n2 1.8

  25. SIGNIFICANT FIGURES Except when all numbers are integers it is impossible to measure the exact value of a quantity. The uncertainty in a measurement is indicated by the number of significant figures, which are the meaningful digits in a measured or calculated quantity.The last digit is understood to be uncertain by + or - 1.

  26. Uncertainty(per cent) 6 mL 17% 6.0 mL 1.7% 6.00 mL 0.17% As a rule of thumb estimate 1 figure beyond the smallest subdivision on the scale.

  27. The Number of Significant Figures in a Measurement Depends Upon the Measuring Device Fig 1.15A

  28. Fig. 1.6

  29. Significant Figures • Any digit that is not zero is significant • 1.234 kg 4 significant figures • Zeros between nonzero digits are significant • 606 m 3 significant figures • Zeros to the left of the first nonzero digit are not significant • 0.08 L 1 significant figure • If a number is greater than 1, then all zeros to the right of the decimal point are significant • 2.0 mg 2 significant figures • If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant • 0.00420 g 3 significant figures (4.20 mg) 1.8

  30. For numbers that do not contain decimal points, the trailing zeros may or may not be significant. Use scientific notation to avoid ambiguity. • Example 1.3 • 478 cm b. 6.01 cm c. 0.825 m • d. 0.043 kg e. 1.310 x 1022 atoms f. 7000 mL Do Practice Exercise

  31. How many significant figures are in each of the following measurements? 24 mL 2 significant figures 4 significant figures 3001 g 0.0320 m3 3 significant figures 6.4 x 104 molecules 2 significant figures 560 kg 2 significant figures 1.8

  32. 89.332 + 1.1 one significant figure after decimal point two significant figures after decimal point 90.432 round off to 90.4 round off to 0.79 3.70 -2.9133 0.7867 Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers. 1.8

  33. 3 sig figs round to 3 sig figs 2 sig figs round to 2 sig figs Significant Figures Multiplication or Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x 3.6666 = 16.536366 = 16.5 6.8 ÷ 112.04 = 0.0606926 = 0.061 1.8

  34. 6.64 + 6.68 + 6.70 = 6.67333 = 6.67 = 7 3 Significant Figures Exact Numbers Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures The average of three measured lengths; 6.64, 6.68 and 6.70? Because 3 is an exact number 1.8

  35. Example 1.4 Practice Exercise • 26.5862 L + 0.17 L • 9.1 g – 4.682 g • 7.1 x 104 dm x 2.2654 x 102 • 6.54 g / 86.5542 mL • 7.55 x 104 m – 8.62 x 103 m

  36. Accuracy – how close a measurement is to the true value Precision – how close a set of measurements are to each other accurate & precise precise but not accurate not accurate & not precise 1.8

  37. 1000 mL 1L L2 1.63 L x = 1630 mL mL 1L 1.63 L x = 0.001630 1000 mL Factor-Label Method of Solving Problems • Determine which unit conversion factor(s) are needed • Carry units through calculation • If all units cancel except for the desired unit(s), then the problem was solved correctly. How many mL are in 1.63 L? 1 L = 1000 mL 1.9

  38. 60 min m x x x 343 60 s 1 mi s 1 hour = 767 1 min 1609 m mi hour The speed of sound in air is about 343 m/s. What is this speed in miles per hour? meters to miles seconds to hours 1 mi = 1609 m 1 min = 60 s 1 hour = 60 min 1.9

  39. Common SI-English Equivalent Quantities Quantity English to SI Equivalent Length 1 mile = 1.61 km 1 yard = 0.9144 m 1 foot (ft) = 0.3048 m 1 inch = 2.54 cm (exactly!) Volume 1 cubic foot = 0.0283 m3 1 gallon = 3.785 dm3 1 quart = 0.9464 dm3 1 quart = 946.4 cm3 1 fluid ounce = 29.6 cm3 Mass 1 pound (lb) = 0.4536 kg 1 pound (lb) = 453.6 g 1 ounce = 28.35 g

  40. Example 1.5 Practice exercise A roll of Al foil has a mass of 1.07 kg. What is the mass in pounds. Example 1.6 Practice exercise The density of silver is 10.5 g/cm3. Convert this to kg/m3.

  41. Sample Problem The volume of an irregularly shaped solid can be determined from the volume of water it displaces. A graduated cylinder contains 245.0 mL water. When a small piece of Pyrite, an ore of Iron, is submerged in the water, the volume increases to 315.8 mL. What is the volume of the piece of galena in cm3 and in liters. Vol (mL) = 315.8 mL - 245.0 mL = 70.8 mL Vol (cm3) = 70.8 mL x 1 cm3/ 1 mL = 70.8 cm3 Vol (liters) = 70.8 mL x 10 –3 L / mL = 7.08 x 10 -2 L

  42. Density Density = mass/volume (an intensive quantity) A small rectangular slab of lithium has a mass of 1.49 g and measures 2.09 cm by 1.11 cm by 1.19 cm. Find its density in g/mL and lb/in3. What is the volume of 15.3 g of Li? What is the mass of 134 mL of Li?

  43. Densities of Some Common Substances Substance Physical State Density (g/cm3) Hydrogen Gas 0.000089 Oxygen Gas 0.0014 Grain alcohol Liquid 0.789 Water Liquid 1.0 Table salt Solid 2.16 Aluminum Solid 2.70 Lead Solid 11.3 Gold Solid 19.3

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