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Chapter 1: Chemistry and Measurement

Chapter 1: Chemistry and Measurement. Renee Y. Becker Valencia Community College CHM 1045. Properties of Matter. Chemistry: The study of composition, properties, and transformations of matter Matter: Anything that has both mass & volume Hypothesis: Interpretation of results

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Chapter 1: Chemistry and Measurement

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  1. Chapter 1: Chemistry and Measurement Renee Y. Becker Valencia Community College CHM 1045

  2. Properties of Matter • Chemistry: The study of composition, properties, and transformations of matter • Matter: Anything that has both mass & volume • Hypothesis: Interpretation of results • Theory: Consistent explanation of observations

  3. Conservation of Mass • Law of Mass Conservation: Mass is neither created nor destroyed in chemical reactions.

  4. Example 1: Conservation of Mass C(s) + O2(g) CO2(g) • 12.3g C reacts with 32.8g O2, ?g CO2 • 0.238g C reacts with ?g O2 to make .873g CO2 • ?g C reacts with 1.63g O2 to make 2.24g CO2

  5. Dalton’s Atomic Theory • Law of Definite Proportions:Different samples of a pure chemical substance always contain the same proportion of elements by mass. • Any sample of H2O contains 2 hydrogen atoms for every oxygen atom

  6. Matter • Matter is any substance that has mass and occupies volume. • Matter exists in one of three physical states: • solid • liquid • gas

  7. Gases • In a gas, the particles of matter are far apart and uniformly distributed throughout the container. • Gases have an indefinite shape and assume the shape of their container. • Gases can be compressed and have an indefinite volume. • Gases have the most energy of the three states of matter.

  8. Liquid • In a liquid, the particles of matter are loosely packed and are free to move past one another. • Liquids have an indefinite shape and assume the shape of their container. • Liquids cannot be compressed and have a definite volume. • Liquids have less energy than gases but more energy than solids.

  9. Solid • In a solid, the particles of matter are tightly packed together. • Solids have a definite, fixed shape. • Solids cannot be compressed and have a definite volume. • Solids have the least energy of the three states of matter.

  10. Phases

  11. Changes in Physical State • Most substances can exist as either a solid, liquid, or gas. • Water exists as a solid below 0 °C; as a liquid between 0 °C and 100 °C; and as a gas above 100°C. • A substance can change physical states as the temperature changes.

  12. Solid  Liquid • When a solid changes to a liquid, the phase change is called melting. • A substance melts as the temperature increases. • When a liquid changes to a solid, the phase change is called freezing. • A substance freezes as the temperature decreases.

  13. Liquid  Gas • When a liquid changes to a gas, the phase change is called vaporization. • A substance vaporizes as the temperature increases. • When a gas changes to a liquid, the phase change is called condensation. • A substance condenses as the temperature decreases.

  14. Solid  Gas When a solid changes directly to a gas, the phase change is called sublimation. A substance sublimes as the temperature increases. When a gas changes directly to a solid, the phase change is called deposition. A substance undergoes deposition as the temperature decreases.

  15. Classifications of Matter • Matter can be divided into two classes: • mixtures • pure substances • Mixtures are composed of more than one substance and can be physically separated into its component substances. • Pure substances are composed of only one substance and cannot be physically separated.

  16. Mixtures • There are two types of mixtures: • homogeneous mixtures • heterogeneous mixtures • Homogeneous mixtures have uniform properties throughout. • Salt water is a homogeneous mixture. • Heterogeneous mixtures do not have uniform properties throughout. • Sand and water is a heterogeneous mixture.

  17. Pure Substances • There are two types of pure substances: • Compounds • Elements • A compound is a substance composed of two or more elements chemically combined • Compounds can be chemically separated into individual elements. • Water is a compound that can be separated into hydrogen and oxygen. • An element cannot be broken down further by chemical reactions.

  18. Properties of Matter • Properties: describe or identify matter • Intensive Properties do not depend on amount • temperature, boiling point, melting point • Extensive Properties do depend on amount. • length and volume

  19. Properties of Matter • Physical Properties can be determined without changing the chemical makeup of the sample. • Some typical physical properties are: • Melting Point, Boiling Point, Density, Mass, Touch, Taste, Temperature, Size, Color, Hardness, Conductivity. • Some typical physical changes are: • Melting, Freezing, Boiling, Condensation, Evaporation, Dissolving, Stretching, Bending, Breaking.

  20. Properties of Matter • Chemical Propertiesare those that do change the chemical makeup of the sample. • Some typical chemical properties are: • Burning, Cooking, Rusting, Color change, Souring of milk, Ripening of fruit, Browning of apples, Taking a photograph, Digesting food. • Note: Chemical properties are actually chemical changes

  21. PHYSICAL CHEMICAL CHANGE New form of old substance. No new substances formed. Old substance destroyed. New substance formed. PROPERTIES Description by senses – shape, color, odor, etc. Measurable properties – density, boiling point, etc. List of chemical changes possible. Properties of Matter

  22. Example 2: Matter Which of the following represents a mixture?

  23. Accuracy, Precision, and Significant Figures in Measurement • Accuracy is how close to the true value a given measurement is. • Precision is how well a number of independent measurements agree with one another.

  24. Scientific Notation • Changing numbers into scientific notation • Large # to small # • Moving decimal place to left, positive exponent 123,987 = 1.23987 x 105 • Small # to large # • Moving decimal place to right, negative exponent 0.000239 = 2.39 x 10-4 How to put into calculator

  25. Example 3: Scientific Notation Put into or take out of scientific notation • 87542 • 2.1956 x 10-3 • 0.784 • 2.78 x 106 • 92000

  26. Accuracy, Precision, and Significant Figures in Measurement • Significant Figures are the total number of digits in the measurement. • The results of calculations are only as reliable as the least precise measurement!! • Rules exist to govern the use of significant figures after the measurements have been made.

  27. Accuracy, Precision, and Significant Figures in Measurement • Rules for Significant Figures: • Zeros in the middle of a number are significant • Zeros at the beginning of a number are not significant • Zeros at the end of a number and following a period are significant • Zeros at the end of a number and before a period may or may not be significant.

  28. Example 4: Significant Figures How many Sig. Figs ? a) 0.000459 b) 12.36 c) 36,450 d) 8.005 e) 28.050

  29. Accuracy, Precision, and Significant Figures in Measurement • Rules for Calculating Numbers: • During multiplication or division, the answer can’t have more sig figs than any of the original numbers.

  30. Example 5: Significant Figures • 238.5 x 79 = • 12 / 0.1272 = • 0.2895 x 0.29 = • 32.567 / 22.98 =

  31. Accuracy, Precision, and Significant Figures in Measurement -During addition or subtraction, the answer can’t have more digits to the right of the decimal point than any of the original numbers.

  32. Example 6: Significant Figures • 238.5 + 79 = • 12.3 - 0.1272 = • 0.2895 + 0.29 = • 32.567 - 22.98 =

  33. Accuracy, Precision, and Significant Figures in Measurement • Rules for Rounding Numbers: • If the first digit removed is less than 5 • round down (leave # same) • If the first digit removed is 5 or greater • round up • Only final answers are rounded off, do not round intermediate calculations

  34. Example 7: Rounding and Significant Figures Round off each of the following measurements • 3.774499 L to four significant figures (b) 255.0974 K to three significant figures (c) 55.265 kg to four significant figures

  35. Example 8: Accuracy & Precision • Which of the following is precise but not accurate?

  36. Measurement and Units SI Units

  37. Measurement and Units Some prefixes for multiples of SI units * * * * * * * * Important

  38. Measurement and Units • Temperature Conversions: The Kelvin and Celsiusdegree are essentially the same because both are one hundredth of theinterval between freezing and boiling points of water.

  39. Measurement and Units • Temperature Conversions: • Celsius (°C) — Kelvin temperature conversion: Kelvin (K) = °C + 273.15 • Fahrenheit (°F) — Celsius (C) temperature conversions: C = 5/9 (F – 32) F = (9/5 x C) + 32

  40. Example 9: Temp. Conversions Carry out the indicated temperature conversions: (a) –78°C = ? K (b) 158°C = ? °F (c) 375 K = ? °C (d) 98.6°F = ? °C (e) 98.6°F = ? K

  41. Measurement and Units • Density: relates the mass of an object to its volume. Density = mass / VolumeD = m / V V = m / D m = V  D • Density decreases as a substance is heated because the substance’s volume increases.

  42. Example 10: Density What is the density of glass (in grams per cubic centimeter) if a sample weighing 26.43 g has a volume of 12.40 cm3?

  43. Example 11: Density What is the volume of an unknown solution if the mass is 12.567 g and the density is 14.621 g/mL ?

  44. Example 12: Density What is the mass of an unknown solution if the mass is 20.2 mL and the density is 2.613 g/mL ?

  45. Measurements and Units • Dimensional-Analysis method uses a conversion factor to express the relationship between units. Original quantity x conversion factor = equivalent quantity Example: express 2.50 kg  lb. Conversion factor: 1.00 kg = 2.205 lb 2.50 kg x 2.205 lb = 6.00 lb 1.00 kg

  46. Measurements and Units

  47. Example 13: Conversions • 1.267 km  m  cm • .784 L  mL • 3.67 x 105 cm  in

  48. Example 14: Conversions • 79 oz  g • 9.63 x 10-3 yd  ft • 23.5 cm2  m2

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