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Chem 121.08 Fall 2013

Chem 121.08 Fall 2013. Chemistry 121.08 Fall 2013. Tuesdays: Lecture 6:00-8:20 pm with breaks Worksheet: 8:20-9:20pm Thursdays: Lecture 6:00-8:00 pm with a break Lab: 8:00-9:20pm. Reminders. First worksheet on Monday Worksheet: group quiz Closed book (allowed a cheat sheet)

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Chem 121.08 Fall 2013

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  1. Chem 121.08 Fall 2013

  2. Chemistry 121.08 Fall 2013 Tuesdays: Lecture 6:00-8:20 pm with breaks Worksheet: 8:20-9:20pm Thursdays: Lecture 6:00-8:00 pm with a break Lab: 8:00-9:20pm

  3. Reminders • First worksheet on Monday • Worksheet: • group quiz • Closed book (allowed a cheat sheet) • Allowed to talk to anyone in the classroom • On Monday I will lecture on lab1 (Accuracy and Density) • Labs posted on facweb.

  4. Chemistry, matter & mass • 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. Chemistry is the study of matter

  5. States of matter

  6. Physical vs chemical properties! Physical property is a property that can be observed or measured without changing the identity of that matter Examples: color, odor, shape, mass ect • Chemical property is a property that can be observed or measured only by changing the identity of matter into a new substance. • Examples: flammability and the ability to react

  7. Physical changes vs chemical change • Chemical change are always accompanied by a change in the identity of matter (composition). Physical change take place without changing the identity of matter. Examples: freezing, melting, or evaporation of a substance • Examples: burning of paper and the fizzing of a mixture of vinegar and baking soda

  8. Classification of matter Most pure Most impure

  9. What is a significant figure? Exact number : results from counting objects or is part of a definition Approximate: results from a measurement or observation and contains some uncertainty. 65.2g 29.04g 25.70 0.0254 2,570 3 significant figure 4 significant figures 4 significant figures 3 significant figures 3 significant figures

  10. Multiplication and division • Rule: In multiplication and division the answer has the same significant figures as the original number with the fewest significant figures. • Example • 56.78 cm x 2.45cm = 139.111 cm2 • Round to  139cm2 • if last number is 4 or less then round down and if last number 5 or more then round up.

  11. Rounding off numbers

  12. Adding and subtracting • Rule: In addition and subtraction the answer has the same number of decimal places as the original number with the fewest decimal places. • Example: • 2.45cm + 1.2cm = 3.65cm, • Round off to = 3.7cm • 7.432cm + 2cm = 9.432 round to  9cm

  13. Question 1.55 • Carry out each calculation and report the answer using the proper number of significant figures. • a. 53.6x0.41 = 21.976 round to 22 3 2 2 • b. 25.825-3.86 = 21.965 round to 21.97 5 3 two digits after the decimal c. 65.2/12 = 5.433 round to 54 3 2 2

  14. Question 1.55 • d. 41.0+9.135 = 50.135 round to 50.1 3 4 on place after decimal • e. 694.2x0.2 = 138.82 round to 100 4 1 1 • f. 1,045- 1.26 = 1043.74 round to 1044 4 3 0 place after decimal

  15. Scientific notation • Scientific notation: provides a convenient way to express very large or very small numbers. In scientific notation, a number is written as: Exponent: Any positive or negative whole number. y x 10x Coefficient: A number between 1 and 10.

  16. Scientific notation HOW TO Convert a Standard Number to Scientific Notation Convert these numbers to scientific notation. Example 2,500 0.036 Move the decimal point to give a number between 1 and 10. Step [1] 2500 0.036 Multiply the result by 10x, where x = number of places the decimal was moved. Step [2] • move decimal left, • x is positive • move decimal right, • x is negative 2.5 x 103 3.6 x 10−2

  17. Scientific notation Converting a Number in Scientific Notation to a Standard Number • When the exponent x is positive, move the decimal point x places to the right. 2.800 x 102 = 280.0 • When the exponent x is negative, move the decimal point x places to the left. 2.80 x 10–2 = 0.0280

  18. Question 1.57 • Write each quantity in scientific notations • a. 1,234 g • 1.234 x103 g • b. 0.0000162m • 1.62 x 10-5 m • c. 5,244,000 L • 5.244 x 106 L • d. 0.00562 g • 5.62 x 10-3 g

  19. Question 1.59 • Convert each number to its standard form. • a. 3.4 x108 • 340,000,000 • b. 5.822 x 10-5 • 0.00005822 • c. 3 x 102 • 300 • d. 6.86 x 10-8 g • 0.0000000682

  20. Question 1.61 • Which number in each pair is larger? • a. 4.44 x103 or 4.8 x102 • 4440 or 480 • 4.44 x103 • b. 5.6 x 10-6 or 5.6 x10-5 • 0.0000056 or 0.000056 • 5.6 x10-5 • c. 1.3 x 108 or 52,300,000 • 130,000,000 or 52,300,000 • 1.3 x 108

  21. Conversion factors Using the Factor-Label Method • Conversion factor: A term that converts a quantity in • one unit to a quantity in another unit. desired quantity original quantity x conversion factor = • Conversion factors are usually written as equalities. 2.21 lb = 1 kg • To use them, they must be written as fractions. 2.21 lb 1 kg 1 kg 2.21 lb or

  22. Conversion factors Factor-label method: Using conversion factors to convert a quantity in one unit to a quantity in another unit. • units are treated like numbers • make sure all unwanted units cancel To convert 130 lb into kilograms: 130 lb x ?kg conversion factor = original quantity desired quantity

  23. Using the Factor-Label Method 2.21 lb 1 kg Answer 2 sig. figures 130 lb x or 1 kg 2.21 lb = 59 kg • The bottom conversion factor has the original unit in the denominator. • The unwanted unit lb cancels. • And you get your desired unit of Kg

  24. Using the factor-label method HOW TO Solve a Problem Using Conversion Factors How many grams of aspirin are in a 325-mg tablet? Example Identify the original quantity and the desired quantity, including units. Step [1] original quantity desired quantity 325 mg ? g

  25. Using the factor-label method HOW TO Solve a Problem Using Conversion Factors Step [2] Write out the conversion factor(s) needed to solve the problem. 1 g = 1000 mg This can be written as two possible fractions: or 1000 mg 1g 1 g 1000 mg Choose this factor to cancel the unwanted unit, mg.

  26. Using the factor-label method HOW TO Solve a Problem Using Conversion Factors Step [3] Set up and solve the problem. 1 g 1000 mg 0.325 g 0.325 g 325 mg x = 325 mg 3 sig. figures 3 sig. figures Unwanted unit cancels Step [4] Write the answer with the correct number of significant figures.

  27. Using the factor-label method Solving a Problem Using Two or More Conversion Factors Always arrange the factors so that the denominator in one term cancels the numerator in the preceding term. How many liters is in 1.0 pint? 1.0 pint original quantity ? L desired quantity • Two conversion factors are needed: 2 pints = 1 quart 1.06 quarts = 1 liter 2 pt 1 qt 1 qt 2 pt 1 L 1.06 qt 1.06 qt 1 L or or First, cancel pt. Then, cancel qt.

  28. Solving a problem using two or more conversion factors • Set up the problem and solve: 1 L 1.06 qt 1 qt 2 pt 0.47 L 0.471698113 L 1.0 pt 1.0pt x x = 2 sig. figures 2 sig. figures • Write the answer with the correct number of significant figures.

  29. Question 1.65 • The Average mass of human liver is 1.5kg. Convert that to (a) grams; (b) pounds; (c) ounces • a. 1. Identify the original quantity and the desired quantity, including units. original quantity desired quantity 1.5 Kg ? g 2.Write out the conversion factor(s) needed to solve the problem. 1 kg = 1000 g This can be written as two possible fractions: 1 kg 1000 g 1000 g 1kg or Choose this factor to cancel the unwanted unit, mg.

  30. 3. Set up and solve the problem. 1000 g 1 kg 1500 g 1500 g 1.5 kg x = 1.5 kg 2sig. figures 2sig. figures Unwanted unit cancels 4. Write the answer with the correct number of significant figures.

  31. Question 1.65 • The Average mass of human liver is 1.5kg. Convert that to (a) grams; (b) pounds; (c) ounces • b. 1. Identify the original quantity and the desired quantity, including units. original quantity desired quantity 1.5 Kg ? Ib 2.Write out the conversion factor(s) needed to solve the problem. 1 kg = 2.20 Ib This can be written as two possible fractions: 1 kg 2.20 Ib 2.20 Ib 1kg or Choose this factor to cancel the unwanted unit, kg.

  32. 3. Set up and solve the problem. 2.20 Ib 1 kg 3.30 Ib 3.3Ib 1.5 kg x = 1.5 kg 2sig. figures 2sig. figures Unwanted unit cancels 4. Write the answer with the correct number of significant figures.

  33. Question 1.65 • The Average mass of human liver is 1.5kg. Convert that to (a) grams; (b) pounds; (c) ounces • c. 1. Identify the original quantity and the desired quantity, including units. original quantity desired quantity 1.5 Kg ? oz 2. Two conversion factor(s) needed to solve the problem. 1 kg = 2.20 Ib 1Ib = 16 oz This can be written as two possible fractions: 2.20 Ib 1kg 16 oz 1 Ib And

  34. 3. Set up and solve the problem. 16 oz 1 Ib 2.20 Ib 1 kg 53oz 52.8 oz 1.5 kg 1.5 kg x x = 2sig. figures 2sig. figures Unwanted unit cancels 4. Write the answer with the correct number of significant figures.

  35. Temperature • Temperature is a measure of how hot or cold an object is. • Three temperature scales are used: Degrees Fahrenheit (oF) Degrees Celsius (oC) Kelvin (K) To convert from oF to oC: To convert from oC to oF: oC = oF− 32 1.8 oF = 1.8(oC) + 32 To convert from oC to K: To convert from K to oC: K = oC + 273 oC = K − 273

  36. Temperature Comparing the Three Temperature Scales

  37. Question 1.75 • Which temperature in each pair is higher? • a. -10 oC or 10 oF • oF = 1.8(-10) + 32 • oF = 14 • So -10 oC is higher temp than 10 oF • b. -50oC or -50 oF • -50oC is higher than -50 oF oF = 1.8(oC) + 32 oC = oF− 32 1.8

  38. Density and Specific Gravity Density: A physical property that relates the mass of a substance to its volume. mass (g) density = volume (mL or cc) To convert volume (mL) to mass (g): To convert mass (g) to volume (mL): g mL mL x = g g x = mL mL g inverse of density density

  39. Density and Specific Gravity Example: • If the density of acetic acid is 1.05 g/mL, what is the volume of 5.0 grams of acetic acid? 5.0 g ? mL original quantity desired quantity • Density is the conversion factor, and can be written two ways: 1.05 g 1 mL 1 mL 1.05 g Choose the inverse density to cancel the unwanted unit, g.

  40. Density and Specific Gravity • Set up and solve the problem: 1 mL 1.05 g 5.0 g x 4.8 mL 5.0 g = 4.761904762 mL 2 sig. figures 2 sig. figures Unwanted unit cancels • Write the final answer with the correct number of significant figures.

  41. Question 1.81 • If milk has a density of 1.03g/ml, what is the mass of one quart, reported in Kg? • Given quantity • One quart (qt) of milk and the density of milk 1.03g/ml • Needed conversions • 1qt=946ml, 1.03g/ml and 1kg = 1000g 1.03g 1ml 1kg 1000g 946 mL 1qt 1qt x x x 0.97438kg = 3sig. figures Unwanted unit cancels 3sig. figures 0.974 kg

  42. Density and Specific Gravity Specific gravity: A quantity that compares the density of a substance with the density of water at the same temperature. density of a substance (g/mL) density of water (g/mL) specific gravity = • The units of the numerator (g/mL) cancel the units of the denominator (g/mL). • The specific gravity of a substance is equal to its density, but contains no units.

  43. Question 1.83 • Which is the upper layer when each of the following liquids is added to water? • a. heptane (density = 0.684g/ml) • Less dense than water so upper layer is heptane • b. olive oil (density = 0.92g/ml) • Less dense than water so upper layer is olive oil. • c. chloroform (density = 1.49 g/ml) • More dense than water so upper layer is water • d. carbon tetrachloride (density = 1.59g/ml) • More dense than water so upper layer is water

  44. Question 1.85 • (a). What is the specific gravity of mercury, the liquid used in thermometers, if it has a density of 13.6g/ml? (b) what is the density of ethanol if it has a specific gravity of 0.789? • a. • Specific gravity = 13.6g/ml = 13.6 1 g/ml (b) Density of ethanol = 0.789 g/ml specific gravity density of a substance (g/mL) density of water (g/mL) =

  45. Lab Check in • Please choose your lab partner • See me for assigned drawer number • Each drawer has a lock and on the back of the lock there is a number that you will use to check the combination and open the lock • Before you leave fill and sign your check in pink information sheet • Also sign the yellow safety sheet

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