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Chapter 1

Chapter 1. The Study of Chemistry. Topics. Introduction Scientific Method Classifications of Matter Properties of Matter Units of Measurement – Metric system Temperature Conversion Metric Conversion (Prefixes) Accuracy vs. Precision Significant Figures Density. States of Matter.

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Chapter 1

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  1. Chapter 1 The Study of Chemistry

  2. Topics • Introduction • Scientific Method • Classifications of Matter • Properties of Matter • Units of Measurement – Metric system • Temperature Conversion • Metric Conversion (Prefixes) • Accuracy vs. Precision • Significant Figures • Density

  3. States of Matter • Solid • Liquid • Gas • Plasma

  4. ATOM • Is the simplest unit of matter.

  5. Definitions • Elements – can’t be decomposed further into simpler substances - 118 elements presently - Lv (element 116 - livermorium) • - Fl (element 114 - flerovium) • Compound – combination of 2 or more elements • Pure substance – has distinct properties and composition; does not vary from sample to sample(ex. Water, NaCl)

  6. Definitions • Mixtures – combinations of 2 or more substances (ex. sugar in water) • 2 Types of Mixtures • 1. Homogenous Mixtures (solutions) = 1 phase • 2. Heterogeneous Mixtures = > 2 phases

  7. SOLUTIONS • Homogeneous mixtures are called SOLUTIONS.

  8. Solution • Solution – homogenous mixture • A solution is not necessarily a liquid. Can be gas or solid.

  9. Physical vs. Chemical Properties • Physical properties – can be measured w/o changing identity and composition of substance (ex. Boiling pt.,freezing pt., color, odor, density, hardness) • Chemical properties – describe how substance reacts or changes to form other compounds (ex. Flammability, toxicity)

  10. Changes of State and Properties • Physical changes – does not change composition of compound • Chemical changes – converts to a different chemical substance • Intensive Properties – independent of amt. (ex. Density, Temperature, Melting Pt) • Extensive Properties – dependent on amt. (ex. Mass, Volume)

  11. Units of Measurement • Mass – grams; kilogram • Length – centimeter; meter • Volume – milliliter or cubic centimeter (cm3) • Temperature – Celcius; Kelvin

  12. Precision vs Accuracy • Accuracy – when acquired value agrees with true value • Precision – when acquired values exhibit reproducibility

  13. Significant Figures • More significant figures = more certainty • Helps in determining how to round measured values and still precise

  14. SIGNIFICANT FIGURES • In counting and definitions, there are an infinite number of sig figs • In measurements, the number of sig figs consists of all certain and the first uncertain digits • Unit conversions do not determine # of sig. figs.

  15. Rules of Significant Figures • 1. Non-zero integers always count. • Ex. 1234.5 grams = 5 Sig. Figs. • 2. Captive zeros are always significant. • Ex. 100.3 grams = 4 Sig. Figs.

  16. Rules of Significant Figures • 3. Leading zeros are NEVER significant. • Ex. 0.6780 grams = 4 Sig. Figs. • 4. Trailing zeroes are significant ONLY if there is a decimal point • Ex. 12.0 grams = 3 Sig. Figs • 120 grams = 2 Sig. Figs

  17. Rules of Significant Figures • 5. Exact numbers (obtained by counting) are infinite and do not determine the number of significant figures. • Example: 4 cows = ?

  18. Determine the # of Sig. Fig. • 200.0 • 1050 • 3003 • 0.0006 • 10,000 • 0.5

  19. Rules of Significant Figures • Multiplication/Division • Answer will have the same # of sig figs as the value with the least # of sig figs • Ex: 3.8 x 200.0 = 2 Sig. Figs.

  20. Rules of Significant Figures • Addition/Subtraction • Answer has the same # of decimal places as the number with the least # of decimal places • Ex. 3.1 + 2.500 + 5.76 = 11.4

  21. Order of Operations • Parenthesis • Multiplication/division • Addition/subtraction

  22. Rounding • Look only to the right of the number you are rounding to: • - If 5 or more, round up • - If less than 5, round down

  23. General Rule • Carry ALL figures through to the end of a problem. Round the final answer to the correct number of significant figures

  24. Problem • Indicate the number of sig. figs. in each of the following measured quantities: • A. 358 kg • B. 0.054 s • C. 6.3050 cm • D. 0.0105 L • E. 7.0500 x 10-3 m3

  25. Problem • Carry out the following operations and express the answer with the appropriate number of sig. figs. • A. 12.0550 + 9.05 • B. 257.2 – 19.789 • C. (6.21 x 103)(1.1050) • D. 0.0577 / 0.753

  26. Prefixes in Metric System • Mega - million • Kilo - 1,000 • Hecto - 100 • Deka - 10 • ----- - 1 (liter, gram, meter) • Deci - 1/10 or 0.1 • Centi - 1/100 or 0.01 • Milli - 1/1000 or 0.001

  27. Temperature Conversions • 0 oC = 273.15 K • oF = 1.8 oC + 32

  28. Things to Remember! • 1 milliliter = 1 cc • 1000 milliliter = 1 liter • 0 oC = 32 oF = 273.15 K

  29. Density • Is the amount of mass in a unit volume of the substance • Is affected by Temperature. • The higher the temp., the lower the density. D = mass of substance = grams volume of substance mL or cm3

  30. Density • Density = mass volume = gram mL

  31. Different ways of calculating volume • I. For solids with regular shapes: • A. For a cube: Vcube = s3 • B. For a rectangular solid, V = L x W x H • C. For a cylinder: V= pr2h • D. For a sphere: V = 4/3 pr3

  32. Different ways of calculating volume • II. For an Irregular Solid • Water displacement

  33. Different ways of calculating volume • III. For a liquid • Use of graduated cylinder, beaker, pipet or buret.

  34. Problem • A cube of osmium metal 1.500 cm on a side has a mass of 76.31 grams at 25 oC. What is its density in g/cm3 at this temperature?

  35. Problem • The density of titanium metal is 4.51 g/cm3 at 25 oC. What mass of titanium displaces 65.8 mL of water at 25 oC?

  36. Problem • The density of benzene at 15 oC is 0.8787 g/mL. Calculate the mass of 0.1500 L of benzene at this temperature.

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