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LEARNING OBJECTIVES/ASSESSMENT 

LEARNING OBJECTIVES/ASSESSMENT  When you have completed your study of this chapter, you should be able to:   1. Explain what matter is.  (Section 1.1; Exercise 1.2)   2. Explain differences between the terms physical and chemical  as applied to:  

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LEARNING OBJECTIVES/ASSESSMENT 

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  1. LEARNING OBJECTIVES/ASSESSMENT  When you have completed your study of this chapter, you should be able to:   1. Explain what matter is.  (Section 1.1; Exercise 1.2)   2. Explain differences between the terms physical and chemical  as applied to:       a.  Properties of matter (Section 1.2; Exercises 1.10 b & c)      b.  Changes in matter (Section 1.2; Exercises 1.8 a & b)   3. Describe matter in terms of the accepted scientific model.   (Section 1.3; Exercise 1.12)   4. On the basis of observation or information given to you, classify  matter into the correct category of each of the following pairs:      a.  Heterogeneous or homogeneous (Section 1.4; Exercise 1.22)      b.  Solution or pure substance (Section 1.4; Exercise 1.24)      c.  Element or compound (Section 1.4; Exercise 1.18) 

  2. 5. Recognize the use of measurement units in everyday activities.  5. Recognize the use of measurement units in everyday activities.   (Section 1.5; Exercise 1.28)   6. Recognize units of the metric system, and convert measurements done  using the metric system into related units.  (Section 1.6; Exercises 1.30 and 1.40)   7. Express numbers using scientific notation, and do calculations with  numbers expressed in scientific notation.  (Section 1.7; Exercises 1.48 and 1.60)   8. Express the results of measurements and calculations using the  correct number of significant figures.   (Section 1.8; Exercises 1.64 and 1.66)    9. Use the factor‐unit method to solve numerical problems.   (Section 1.9; Exercise 1.82)   10. Do calculations involving percentages.  (Section 1.10; Exercise 1.92)   11. Do calculations involving densities.  (Section 1.11; Exercise 1.98)  

  3. Topics covered in Chapter One All sections of Chapter One will be covered. Example exercises include: 1.7 – 1.17 1.22 – 1.25 1.29 – 1.38 1.46 – 1.75 1.77 – 1.80 1.87 – 1.99

  4. MATTER & MASS • Matter is anything that has mass and occupies space. • Mass is a measurement of the amount of matter in an object. • Mass is independent of the location of an object. • An object on the earth has the same mass as the same object on the moon.

  5. WEIGHT • Weight is a measurement of the gravitational force acting on an object. • Weight depends on the location of an object. • An object weighing 1.0 lb on earth weighs about 0.17 lb on the moon.

  6. PHYSICAL & CHEMICAL PROPERTIES • PHYSICAL PROPERTIESOF MATTER • Physical properties can be observed or measured without attempting to change the composition of the matter being observed. • Examples: physical state (duh), color, shape and mass • CHEMICAL PROPERTIESOF MATTER • Chemical properties can be observed or measured only by attempting to change the matter into new substances. • Examples: flammability and the ability to react (e.g. when vinegar and baking soda are mixed)

  7. PHYSICAL & CHEMICAL CHANGES • PHYSICAL CHANGES OF MATTER • Physical changes take place without a change in composition. • Examples: changes of state, like freezing, melting, or evaporation of a substance (e.g. water) • CHEMICAL CHANGES OF MATTER • Chemical changes are always accompanied by a change in composition. Most color changes. • Examples: burning of paper and the fizzing of a mixture of vinegar and baking soda

  8. Physical or Chemical Properties? • Color of gold • Tendency of silver to tarnish • Flammability of gasoline • Boiling point of alcohol • Smell of perfume • Different properties of ozone • Has the formula of O3 • Bluish color • A gas a room temperature • Decomposes on exposure to sunlight

  9. Physical or Chemical Changes? • Sugar dissolves in hot water • Copper metal turns green over time with exposure to air • Compressed liquid propane evaporates when the valve is opened on the container • Sugar burns in a pot when heated to a high temperature

  10. PARTICULATE MODEL OF MATTER • All matter is made up of tiny particles called molecules andatoms. • MOLECULES • A molecule (or atom, depending on the substance) is the smallest particle of a pure substance that is capable of a stable independent existence. • ATOMS • Atoms are the particles that make up molecules.

  11. MOLECULE CLASSIFICATION • Diatomic molecules contain two atoms. • Triatomic molecules contain three atoms. • Polyatomic molecules contain many atoms.

  12. MOLECULE CLASSIFICATION (continued) • HOMOATOMIC MOLECULES • The atoms contained inhomoatomic moleculesare of the same kind. • HETEROATOMIC MOLECULES • The atoms contained inheteroatomic moleculesare of two or more kinds. homoatomic heteroatomic

  13. MOLECULE CLASSIFICATION EXAMPLE • Classify the molecules in these diagrams using the terms diatomic, triatomic, or polyatomic molecules. • Classify the molecules using the terms homoatomic or heteroatomic molecules.

  14. CLASSIFICATION OFMATTER • Matter can be classified into several categories based on chemical and physical properties. • PURE SUBSTANCES • Pure substances have a constant composition and a fixed set of otherphysicaland chemical properties. • Example: pure water (always contains the same proportions of hydrogen and oxygen and freezes at a specific temperature)

  15. CLASSIFICATION OF MATTER(continued) • MIXTURES • Mixturescan vary in composition and properties. • Example: mixture of table sugar and water (can have different proportions of sugar and water) • A glass of water could contain one, two, three, etc. spoons of sugar. • Properties such as sweetness would be different for the mixtures with different proportions.

  16. HETEROGENEOUS MIXTURES • The properties of a sample of a heterogeneous mixture depends on the location from which the sample was taken. • A pizza pie is a heterogeneous mixture. A piece of crust has different properties than a piece of pepperoni taken from the same pie.

  17. HOMOGENEOUS MIXTURES • Homogeneous mixtures are also called solutions. The properties of a sample of a homogeneous mixture are the same regardless of where the sample was obtained from the mixture. • Samples taken from any part of a mixture made up of one spoon of sugar mixed with a glass of water will have the same properties, such as the same taste.

  18. ELEMENTS • Elementsare pure substances that are made up of homoatomic molecules or individual atoms of the same kind. • Examples: oxygen gas made up of homoatomic molecules and copper metal made up of individual copper atoms

  19. COMPOUNDS • Compoundsare pure substances that are made up of heteroatomic molecules or individual atoms (ions) of two or more different kinds. • Examples: pure water made up of heteroatomic molecules and table salt made up of sodium atoms (ions) and chlorine atoms (ions)

  20. MATTER CLASSIFICATION SUMMARY

  21. MATTER CLASSIFICATION EXAMPLE • Classify H2, F2, and HF using the classification scheme from the previous slide. • Solution: • H2, F2, and HF are all pure substances because they have a constant composition and a fixed set of physicaland chemical properties. • H2 and F2 are elements because they are pure substances composed of homoatomic molecules. • HF is a compound because it is a pure substance composed of heteroatomic molecules.

  22. Matter Classification • Coffee • Apple juice • Gold • Table salt • Methane (CH4) • Urine • Clean air • Pure water • Nitrogen gas • Concrete • Bronze (composed of tin and copper) • water and gasoline mixed together

  23. MEASUREMENTS & UNITS • Measurements consist of two parts, a number and a unit or label such as feet, pounds, or gallons. • Measurement units are agreed upon by those making and using the measurements. • Measurements are made using measuring devices (e.g. rulers, balances, graduated cylinders, etc.). (Name other common measuring devices)

  24. METRIC SYSTEM • The metric system is a decimal system in which larger and smaller units are related by factors of 10. • TYPES OF METRIC SYSTEM UNITS • Basic or defined units [e.g. 1 meter (1 m)] are used to calculate derived units [e.g. 1 square meter (1 m2)]. • Which of the above are basic units?

  25. THE USE OF PREFIXES • Prefixes are used to relate basic and derived units. • The common prefixes are given in the following table:

  26. TEMPERATURE SCALES • The three most commonly-used temperature scales are the Fahrenheit, Celsius and Kelvin scales. • The Celsius and Kelvin scales are used in scientific work.

  27. TEMPERATURE CONVERSIONS • Readings on one temperature scale can be converted to the other scales by using mathematical equations. • Converting Fahrenheit to Celsius. • Converting Celsius to Fahrenheit. • Converting Kelvin to Celsius. • Converting Celsius to Kelvin.

  28. TEMPERATURE CONVERSION PRACTICE • Covert 22°C and 54°C to Fahrenheit and Kelvin.

  29. COMMONLY USED METRIC UNITS

  30. SCIENTIFIC NOTATION • Scientific notation provides a convenient way to express very large or very small numbers. • Numbers written in scientific notation consist of a product of two parts in the form M x 10n, where M is a number between 1 and 10 (but not equal to 10) and n is a positive or negative whole number. • The number M is written with the decimal in the standard position.

  31. SCIENTIFIC NOTATION (continued) • STANDARD DECIMAL POSITION • The standard position for a decimal is to the right of the first nonzero digit in the number M. • SIGNIFICANCE OF THE EXPONENT n • A positive n value indicates the number of places to the right of the standard position that the original decimal position is located. • A negative n value indicates the number of places to the left of the standard position that the original decimal position is located.

  32. Convert to scientific notation • 1,000 • 0.1 • 10 • 1 • 1,016,200 • 0.0310 • 0.0000009203 Convert to regular numbers • 1.034 X 100 • 9.02 X 101 • 2.003 X 10-5 • 6.023 X 107

  33. More conversions Put final answer into scientific notation. • 12.5 mm into km • 3.650 X 108ng into g • 125 uL into mL Make these changes in exponential numbers • 670 X 10-8 = ? X 10-11 • 0.0475 X 106 = ? X 105 • 1.25 X 10-3 = ? X 102

  34. SCIENTIFIC NOTATION MULTIPLICATION • Multiply the M values (a and b) of each number to give a product represented by M'. • Add together the n values (y and z) of each number to give a sum represented by n'. • Write the final product as M' x 10n'. • Move decimal in M' to the standard position and adjust n' as necessary.

  35. SCIENTIFIC NOTATION DIVISION • Divide the M values (a and b) of each number to give a quotient represented by M'. • Subtract the denominator (bottom) n value (z) from the numerator (top) n value (y) to give a difference represented by n'. • Write the final quotient as M' x 10n'. • Move decimal in M' to the standard position and adjust n' as necessary.

  36. SCI NOTATION ADD/SUBTRACT • Change the exponential values of numbers involved to be the same for all numbers. • Add or subtract (a and b values) in the usual way, the exponential value doesn’t change • Write the answer as M' x 10n'. • Move decimal in M' to the standard position and adjust n' as necessary. 1.25 X 105 + 4.62 X 104 = 1.25 X 105 + 0.462 X 105 = 1.712 X 105 or 1.71 X 105

  37. SIGNIFICANT FIGURES • Significant figures are the numbers in a measurement that represent the certainty of the measurement, plus one number representing an estimate. • COUNTING ZEROS AS SIGNIFICANT FIGURES • Leading zeros are never significant figures. • Buried zeros are always significant figures. • Trailing zeros are generally significant figures.

  38. SIGNIFICANT FIGURES (continued) Using the apparatus supplied at each station, determine the measurement with the correct number of significant figures.

  39. SIGNIFICANT FIGURES (continued) • The answer obtained by multiplication or division must contain the same number of significant figures (SF) as the quantity with the fewest number of significant figures used in the calculation.

  40. SIGNIFICANT FIGURES (continued) • The answer obtained by addition or subtraction must contain the same number of places to the right of the decimal (prd) as the quantity in the calculation with the fewest number of places to the right of the decimal.

  41. ROUNDING RULES FOR NUMBERS • If the first of the nonsignificant figures to be dropped from an answer is 5 or greater, all the nonsignificant figures are dropped and the last remaining significant figure is increased by one. • If the first of the nonsignificant figures to be dropped from an answer is less than 5, all nonsignificant figures are dropped and the last remaining significant figure is left unchanged. Round 10.825 to 1place to the right of the decimal. ⇒10.8 Round −0.175 to 1 place to the right of the decimal. ⇒ −0.2

  42. EXACT NUMBERS • Exact numbersare numbers that have no uncertainty (they do not affect significant figures). • A number used as part of a defined relationship between quantities is an exact number (e.g. 100 cm = 1 m). • A counting number obtained by counting individual objects is an exact number (e.g. 1 dozen eggs = 12 eggs). • A reduced simple fraction is an exact number (e.g. 5/9 in equation to convert ºF to ºC).

  43. EXACT NUMBERS Some exact numbers are defined. 2.54 cm = 1 in (exact numbers) 1 mL = 1 cm3

  44. Calculate the following Report answers to the correct number of significant figures and round as appropriate. • 89.3 X 77.0 X 0.08 = • (7.890 X 1012) / (6.7 X 104) = • 67.8 X 9.8 / 100.04 = • 89.6 + 98.33 – 4.674 = • 0.004 + 0.09879 = • (2.45 X 108) + (1.225 X 10-3) = • (568.99 – 232.1) X 5.3 =

  45. USING UNITS IN CALCULATIONS • The factor-unit method for solving numerical problems is a four-step systematic approach to problem solving. • Step 1: Write down the known or given quantity. Include both the numerical value and units of the quantity. • Step 2: Leave some working space and set the known quantity equal to the units of the unknown quantity. • Step 3: Multiply the known quantity by one or more factors, such that the units of the factor cancel the units of the known quantity and generate the units of the unknown quantity. • Step 4: After you generate the desired units of the unknown quantity, do the necessary arithmetic to produce the final numerical answer.

  46. SOURCES OF FACTORS • The factors used in thefactor-unit method are fractions derived from fixed relationships between quantities. These relationships can be definitions or experimentally measured quantities. • An example of a definition that provides factors is the relationship between meters and centimeters: 1m = 100cm. This relationship yields two factors: and

  47. FACTOR UNIT METHOD EXAMPLES • A length of rope is measured to be 1834 cm. How many meters is this? • Solution: Write down the known quantity (1834 cm). Set the known quantity equal to the units of the unknown quantity (meters). Use the relationship between cm and m to write a factor (100 cm = 1 m), such that the units of the factor cancel the units of the known quantity (cm) and generate the units of the unknown quantity (m). Do the arithmetic to produce the final numerical answer.

  48. PERCENTAGE • The word percentage means per one hundred. It is the number of items in a group of 100 such items. • PERCENTAGE CALCULATIONS • Percentages are calculated using the equation: • In this equation, part represents the number of specific items included in the total number of items.

  49. EXAMPLE PERCENTAGE CALCULATION • A student counts the money she has left until pay day and finds she has $36.48. Before payday, she has to pay an outstanding bill of $15.67. What percentage of her money must be used to pay the bill? • Solution: Her total amount of money is $36.48, and the part is what she has to pay or $15.67. The percentage of her total is calculated as follows:

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