Matter and Energy: The Origin of the Universe

1. 1 Matter and Energy: The Origin of the Universe

2. Chapter Outline • 1.1 Classes of Matter • Classification of Matter: Pure Substances vs. Mixtures • Elements vs. Compounds • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

3. Chemistry: Definitions Chemistry – the study of the composition, structure, and properties of matter and of the energy consumed or given off when matter undergoes a change. Matter– anything that occupies space and has mass. Mass– defines the quantity of matter in an object. Energy– the capacity to do work.

4. Classes of Matter • Types of Matter: • Pure substance • Same physical and chemical properties throughout • Cannot be separated into simpler substances by a physical process • Mixture • Combination of two or more pure substances • Can be separated by physical processes

5. Mixtures • Homogeneous • Components are distributed uniformly throughout the sample and have no visible boundaries or regions. • Heterogeneous • Components are not distributed uniformly, and may have distinct regions of different composition.

6. Pure Substances • Element • A pure substance that cannot be separated into simpler substances by any chemical process. • Compound • Pure substance composed of two or more elements bonded together in fixed proportions. • It can be broken down into individual elements by chemical means.

8. Compounds • Law of Constant Composition • All samples of a particular compound contain the same elements combined in the same proportions. • Water (H2O) • Consists of two units of hydrogen (H) combined with one unit of oxygen (O) • Elements and proportions are represented by chemical formulas (Section 1.2).

9. Figure 1.4 Pure Substances Elements: hydrogen (H2) and oxygen (O2) Compound: water (H2O)

10. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • Atoms, Molecules, and Chemical Formulas • Chemical Reactions and Chemical Equations • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

11. An Atomic View • Atom • The smallest particle of an element that retains the chemical characteristics of that element. • Molecule • A collection of atoms chemically bonded together in characteristic proportions. • Chemical formula • Notation for representing elements and compounds • Consists of symbols of constituent elements, and subscripts identifying the number of atoms of each element in one molecule.

12. Chemical reaction The transformation of one or more substances into different substances. Chemical Reactions • Chemical equation • - Use chemical formulas to express the identities of substances involved in a reaction. • - Use coefficients to indicate quantities of substances involved in a reaction.

13. Chemical bonds link atoms together to make molecules. Arrangement of atoms can be represented in four ways: Chemical Formulas • - Chemical formulas • - Structural formulas • - Ball-and-stick models • - Space-filling models

14. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • Methods of Separation: Filtration and Distillation • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

15. Separating Mixtures • Constituents in a mixture can be isolated by physical means (i.e., no chemical reactions are needed). • Filtration – a process for separating particles from a liquid or gas by passing the mixture through a medium that retains the particles. • Distillation – a separation technique in which the more volatile components of a mixture are vaporized and then condensed, thereby separating them from less volatile components.

16. Distillation Figure 1.12 Collection of pure water from seawater using a solar still.

17. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • COAST: A Strategy for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

18. COAST • Collect and Organize • Identify key concepts and skills required to solve problem, and assemble information needed. • Analyze • Evaluate information and relationships or connections; sometimes units will help identify steps needed to solve the problem. • Solve • Perform calculations, check units, etc. • Think About It • Is the answer reasonable? Are the units correct?

19. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • Intensive Properties vs. Extensive Properties • Chemical Properties vs. Physical Properties • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Result • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

20. Properties of Matter • Intensive property: • A property that is independent of the amount of substance present. • Examples: color, melting point • Extensive property: • A property that varies with the quantity of the substance present. • Examples: volume, mass

21. Properties of Matter (cont.) • Physical property: • A property of a substance that can be observed without changing it into another substance. • Examples: luster, hardness, color, etc. • Chemical property: • A property of a substance that can be observed only by reacting it to form another substance. • Examples: flammability

22. Density = ratio of the mass of an object/substance to the volume of the object/substance: Density (Density: Intensive or extensive property?)

23. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • Solids, Liquids, and Gases • Transformations: Changes in States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

24. States of Matter • Solids • Definite shape and volume • Liquids • Occupies definite volume, but flows to assume the shape of its container • Gases (vapors) • Neither definite volume nor shape; expands to fill its container • Changes of state • Transformation from one state to another due to addition or removal of heat

25. States of Matter (cont.)

26. States of Matter (cont.)

27. States of Matter (cont.)

28. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • Hypotheses, Theories, and Models • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

29. The Scientific Method • Scientific Method • Approach to acquire knowledge through the observation of phenomena • Hypothesis • A tentative and testable explanation for an observation or a series of observations • Scientific Theory • A general explanation of widely observed phenomena • Example: Big Bang Theory in Section 1.10

30. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • SI Units • Significant Figures in Measurements and Calculations • Precision and Accuracy • 1.9 Unit Conversion and Dimensional Analysis • 1.10 Testing a Theory: The Big Bang Revisited

31. Making Measurements Number Unit • Measurements • Essential for characterizing physical and chemical properties of matter • Two parts of every measurement: • 5,280 feet • Standardization of the units of measurement is essential.

35. Uncertainty in Measurements ± 0.01 g ± 0.0001 g • All measurements contain uncertainty. • Depends on instruments used to make measurement • A digit that must be estimated is called uncertain (last recorded digit).

36. Significant Figures • Includes all digits known with certainty plus one digit that is uncertain • Rules for counting significant figures: • All nonzero integers are significant. • Zeros (depends on location): • Leading zeros • Trailing zeros • “Captive” zeros • Exact numbers

37. Rules for Counting Significant Figures • Nonzero integers are always significant. • 3456 → 4 sig. figs. • 7.35 → 3 sig. figs. • Zeros • Leading zeros are not significant. • 0.0392 → 3 sig. figs.

38. Counting Significant Figures (cont.) • Zeros (cont.) • Trailing zeros are not significant unless they come after a decimal point. • 3700 → 2 sig. figs. • 140.00 → 5 sig. figs. • Captive zeros are always significant. • 16.07 → 4 sig. figs. • Exact numbers have an infinite # of sig. figs.

39. Practice: Counting Sig. Figs. 0.04550 g = ?? 100 lb = ?? 101.05 mL = ?? 350.0 g = ?? How many significant figures are in the following numbers?

40. Practice: Counting Sig. Figs. How many significant figures are in the following numbers? 0.04550 g The leading zeros are not significant, but the trailing zero is significant. 100 lb Without a decimal point, these trailing zeros are not significant. 101.05 mL All nonzero integers, plus the captured zeros 350.0 g The presence of the decimal point makes the trailing zeros significant. (4 sig figs.) (1 sig. fig.) (5 sig. figs.) (4 sig figs.)

41. Significant Figures in Mathematical Operations • Rounding off • Drop “insignificant” digits. • Only at the end of calculations! • “Weakest link” Principle • The number of significant figures in the final result cannot be greater than the “weakest link” used in the calculation. • The actual rule depends on the mathematical operation.

42. Significant Figures in Mathematical Operations (cont.) (2 sig. figs.) (3 sig.figs.) (3 sig. figs.) • Multiplication/Division: • # of sig figs in the result = # in the least precise measurement used in the calculation • 6.38  2.0 = 12.76  13 • 16.84/2.54 = 6.6299  6.63 • Addition/Subtraction: • # of sig figs in the result depends on the number of decimal places in the least accurate measurement • 6.8 + 11.934 = 18.734  18.7

43. Practice: Rounding (Answer: 1.9 g/cm3) • Round the answer for the mathematical operation below to the appropriate number of significant figures. 1.23 g - 0.567 g = 1.924975 ? 0.34442 cm3

44. 1.23 g - 0.567 g = 1.924975 0.34442 cm3 This calculation involves multiple steps. To determine the number of significant figures in the numerator, we must first perform addition/subtraction to obtain a value of 0.663. 1.23 g – 0.567 g / 0.34442 cm3 => Practice: Rounding .0663 g / 0.34442 cm3

45. 1.23 g - 0.567 g = 1.924975 0.34442 cm3 Since the first number in the calculation is only known to the second decimal place, then the calculated result in the numerator can only be known to the second decimal place (0.66) and so has only 2 sig, figs. .0663 g => Practice: Rounding 0.66g (2 sig. figs)

46. 1.23 g - 0.567 g = 1.924975 0.34442 cm3 Therefore, the final calculated answer can only have two significant figures (truncate to 1.9). Take notice that we do not round 0.663 before calculating the final result. Simply use the rule to determine how many sig figs would be appropriate in the numerator. 1.924975 g/cm3 => Practice: Rounding 1.9 g/cm3 (2 sig. figs.)

47. Precision and Accuracy Accuracy – agreement between a measured value and the accepted or true value Precision – agreement among repeated measurements Which is accurate? Which is precise?

48. Chapter Outline • 1.1 Classes of Matter • 1.2 Matter: An Atomic View • 1.3 Mixtures and How to Separate Them • 1.4 A Framework for Solving Problems • 1.5 Properties of Matter • 1.6 States of Matter • 1.7 The Scientific Method: Starting Off with a Bang • 1.8 Making Measurements and Expressing the Results • 1.9 Unit Conversion and Dimensional Analysis • Conversion Factors for Converting between Units • 1.10 Testing a Theory: The Big Bang Revisited

49. Changing Units: Conversion Factors • Unit conversion factor • A fraction in which the numerator and denominator represent equivalent quantities, but are expressed in different units. 1 km = 0.6214 mi  • Converting a value from one unit to another:

50. Example: Using Densities as Conversion Factors 8.533 g of Iron 1 mL = 7.87 g 1.08 mL of Fe - To convert volumes into masses. 8.53 g of Fe 1.08 mL of Fe 7.87 g = 1 mL • Use density as a conversion factor: - To convert masses into volumes.