Chemistry II: Preliminaries (Stuff you should already know!)

1. Chemistry II: Preliminaries(Stuff you should already know!) Scientific Method Metric System Uncertainty Sig Figs Dimensional Analysis

2. ACT I The Scientific Method

3. Scientific method • A way of solving problems • Observation- what is seen or measured • Hypothesis- educated guess of why things behave the way they do. (possible explanation) • Experiment- designed to test hypothesis • leads to new observations, • and the cycle goes on

4. Scientific method • After many cycles, a broad, generalizable explanation is developed for why things behave the way they do • Theory • Also regular patterns of how things behave the same in different systems emerges • Law • Laws are summaries of observations • Often mathematical relationship

5. Scientific method • Theories have predictive value. • The true test of a theory is if it can predict new behaviors. • If the prediction is wrong, the theory must be changed. • Theory- why • Law – how • Law – equation of how things change

6. Theory (Model) Law Observations Hypothesis Modify Experiment Prediction Experiment

7. ACT II: The Metric System

8. Metric System • Every measurement has two parts • Number • Scale (unit) • SI system (le Systeme International) based on the metric system • Prefix + base unit • Prefix tells you the power of 10 to multiply by - decimal system -easy conversions

9. Metric System • Base Units • Mass - kilogram (kg) • Length- meter (m) • Time - second (s) • Temperature- Kelvin (K) • Electric current- ampere (amp, A) • Amount of substance- mole (mol)

10. Prefixes • giga- G 1,000,000,000 109 • mega - M 1,000,000 106 • kilo - k 1,000 103 • deci- d 0.1 10-1 • centi- c 0.01 10-2 • milli- m 0.001 10-3 • micro- m 0.000001 10-6 • nano- n 0.000000001 10-9

11. Give it some thought! • Which of the following quantities is the smallest? • 1 mg • 1 mg • 1 pg

12. Deriving the Liter • Liter is defined as the volume of 1 dm3 • gram is the mass of 1 cm3

13. Mass and Weight • Mass is measure of resistance to change in motion • Weight is force of gravity. • Sometimes used interchangeably • Mass can’t change, weight can

14. Give it some thought! • Which of the following quantities represents a volume measurement? • 15 m2 • 2.5 x 102 m3 • 5.77 L/s

15. ACT III: Uncertainty

16. Uncertainty • Basis for significant figures • All measurements are uncertain to some degree • Precision- how repeatable (consistent) • Accuracy- how correct - closeness to true value. (correct) • Random error -Random errors in experimental measurements are caused by unknown and unpredictable changes in the experiment. These changes may occur in the measuring instruments or in the environmental conditions.

17. Uncertainty • Systematic error-Systematic errors in experimental observations usually come from the measuring instruments. These may occur because: there is something wrong with the instrument or its data handling system, or because the instrument is wrongly used by the experimenter. • Want to avoid this • Bad equipment or bad technique. • Better precision implies better accuracy • You can have precision without accuracy • You can’t have accuracy without precision

18. ACT IV: SIG FIG’S

19. SIG FIG RULES • Numbers other than zero are always significant. • One or more final zeros used after the decimal point are always significant. • Zeros between two other numbers are always significant. • Zeros used for spacing the decimal point are not significant.

20. 3 SF SIG FIG RULES • Multiply/Divide - The # with the fewest sig figs determines the # of sig figs in the answer. (13.91g/cm3)(23.3cm3) = 324.103g 4 SF 3 SF 324g

21. SIG FIG RULES • Add/Subtract - The # with the lowest decimal value determines the place of the last sig fig in the answer. 3.75 mL + 4.1 mL 7.85 mL  7.9 mL

22. Sig Fig Examples • How many sig figs are in each of the following? • 8.01 • 80.1 • 80 • 6.2  104 • 0.0083 • 0.1040900300

23. Sig Fig Examples • Round the following to 4 sig figs: • 3.682417 • 375.6523 • 45467.3 • Round the following to 2 decimal places: • 0.03062 • 41.86632 • 22.494

24. (15.30 g) ÷ (6.4 mL)  2.4 g/mL 2 SF Sig Fig Examples Practice Problems 4 SF 2 SF = 2.390625 g/mL 18.9 g - 0.84 g  18.1 g 18.06 g

25. ACT V: Dimensional Analysis

26. Dimensional Analysis • Use conversion factors to change the units • From an equivalence statement you can make 2 conversion factors • 1 foot = 12 inches (equivalence statement) • 12 in =or= 1 ft 1 ft 12 in • multiply by the one that will give you the correct units in your answer.

27. Examples • 11 yards = 2 rod • 40 rods = 1 furlong • 8 furlongs = 1 mile • The Kentucky Derby race is 1.25 miles. How long is the race in rods, furlongs, meters, and kilometers? • A marathon race is 26 miles, 385 yards. What is this distance in rods and kilometers?

28. Examples • Because you never learned dimensional analysis, you have been working at a fast food restaurant for the past 35 years wrapping hamburgers. Each hour you wrap 184 hamburgers. You work 8 hours per day. You work 5 days a week. you get paid every 2 weeks with a salary of $840.34. How many hamburgers will you have to wrap to make your first one million dollars?

29. A senior was applying to college and wondered how many applications she needed to send. Her counselor explained that with the excellent grade she received in chemistry she would probably be accepted to one school out of every three to which she applied. She immediately realized that for each application she would have to write 3 essays, and each essay would require 2 hours work. Of course writing essays is no simple matter. For each hour of serious essay writing, she would need to expend 500 calories which she could derive from her mother's apple pies. Every three times she cleaned her bedroom, her mother would make her an apple pie. How many times would she have to clean her room in order to gain acceptance to 10 colleges?

30. 65 mi hr Multiple units • The speed limit is 65 mi/hr. What is this in m/s? • 1 mile = 1760 yds • 1 meter = 1.094 yds 1760 yd 1 m 1 hr 1 min 1 mi 1.094 yd 60 min 60 s

31. Multiple units • Lead has a density of 11.4 g/cm3. What is this in pounds per quart? • 454 g = 1 lb • 1 L = 1.094 qt

32. 1 m 1 m 1 m 1 m 100 cm 100 cm 100 cm 100 cm 3 Units to a Power • How many m3 is 1500 cm3? 1500 cm3 1500 cm3

33. Temperature and Density

34. Temperature • A measure of the average kinetic energy • Different temperature scales, all are talking about the same height of mercury. • Derive a equation for converting ºF toºC

35. ºF -459 32 212 ºC -273 0 100 K 0 273 373 Temperature Conversions • Always use absolute temperature (Kelvin) when working with gases.

36. C. Temperature Conversions • Conversion formulas for Kelvin, Celsius and Fahrenheit K = ºC + 273 ºC= K - 273

37. Density • Ratio of mass to volume • D = m/V • Useful for identifying a compound • Useful for predicting weight • An intrinsic property- does depend on what the material is

38. Density Problem • What is the mass in grams of a cube of gold if the length of the cube is 2.00 cm? density 19.32 g/cm3