AP Chemistry

1. AP Chemistry Unit 1: Chemical Foundations Chapter 1 Zumdahl

2. Measurement andSignificant Figures www.lab-initio.com

3. Steps in the Scientific Method 1. Observations - quantitative - qualitative 2. Formulating hypotheses - possible explanation for the observation 3. Performing experiments - gathering new information to decide whether the hypothesis is valid

4. Outcomes Over the Long-Term Theory (Model) = WHY? - A set of tested hypotheses that give an overall explanation of some natural phenomenon. Natural Law = HOW? - The same observation applies to many different systems

5. Law vs. Theory A law summarizes what happens A theory (model) is an attempt to explain why it happens. Einstein's theory of gravity describes gravitational forces in terms of the curvature of spacetime caused by the presence of mass

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

7. Scientific Quantities Convert: 2 X 10-3 4 X 102 6050000 0.0009 0.002 400 6.05 X 106 9 X 10-4 Amounts in Chemistry are often very large or very small. Therefore we use scientific notation.

8. Nature of Measurement A measurement is a quantitative observation consisting of 2 parts: • Part 1 - number • Part 2 - scale (unit) Examples: • 20 grams • 6.63 x 10-34 Joule·seconds

9. The Fundamental SI Units(le Système International, SI)

10. SI Units

11. SI Prefixes Common to Chemistry What does the exponent mean?

12. Uncertainty in Measurement • A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty. • Measurements are performed with instruments • No instrument can read to an infinite number of decimal places

13. Precision and Accuracy • Accuracy refers to how close the experimental value is to the true value. (Bull’s-eye) • Precision refers to the degree of agreement among several measurements made in the same manner. Precise but not accurate Precise AND accurate Neither accurate nor precise

14. Types of Error • Random Error(Indeterminate Error) - measurement has an equal probability of being high or low. • Systematic Error(Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique or incorrect calibration. This can result in measurements that are precise, but not accurate. WANT TO AVOID!

15. Because of Uncertainty we have… Significant Figures!

16. Rules for Counting Significant Figures - Details • Nonzero integersalways count as significant figures. • 3456has • 4 sig figs.

17. Rules for Counting Significant Figures - Details • Zeros • - Leading zeros do not count as significant figures. • 0.0486has • 3sig figs.

18. Rules for Counting Significant Figures - Details • Zeros • - Captive zerosalways count as significant figures. • 16.07has • 4sig figs.

19. Rules for Counting Significant Figures - Details • Zeros • Trailing zeros are significant only if the number contains a decimal point. • 9.300has • 4sig figs.

20. Rules for Counting Significant Figures - Details • Exact numbershave an infinite number of significant figures. • 1inch=2.54cm, exactly

21. Pacific Atlantic

22. Sig Fig Practice #1 How many significant figures in each of the following? 1.0070 m  5 sig figs 17.10 kg  4 sig figs 100,890 L  5 sig figs 3.29 x 103 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000  2 sig figs

23. Dimensional Analysis Using the units to solve problems

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

25. 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?

26. 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?

27. 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 made her an apple pie. How many times would she have to clean her room in order to gain acceptance to 10 colleges?

28. Can you do these? • How many miligrams are in 424 kilograms? • How many meters are in 75 picometers?

29. Multiple units 65 mi hr • 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

30. 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

31. Temperature and Density

32. Temperature • A measure of the average kinetic energy • Different temperature scales, all are talking about the same height of mercury.

33. Celsius & Kelvin

34. Equations for Temperature • K=C + 273 • C=(9/5)F

35. Density • Ratio of mass to volume • D = m/V • Useful for identifying a compound • Useful for predicting weight

36. Density Problem • An empty container weighs 121.3 g. Filled with carbon tetrachloride (density 1.53 g/cm3) the container weighs 283.2 g. What is the volume of the container?

37. Density Problem • A 55.0 gal drum weighs 75.0 lbs. when empty. What will the total mass be when filled with ethanol? density 0.789 g/cm3 1 gal = 3.78 L 1 lb = 454 g

38. Organization of Matter Unit 1

39. The Organization of Matter Matter Mixtures: a) Homogeneous (Solutions) b) Heterogeneous Pure Substances Elements Compounds Atoms Nucleus Electrons Protons Neutrons Quarks Quarks

40. Phase Differences Solid– definite volume and shape; particles packed in fixed positions. Liquid– definite volume but indefinite shape; particles close together but not in fixed positions Gas– neither definite volume nor definite shape; particles are at great distances from one another Plasma– high temperature, ionized phase of matter as found on the sun.

41. Changes & Properties Chemical Vs Physical

42. Physical Changes vs. Chemical Changes • Physical changes - A change that changes appearances, without changing the composition. • Chemical changes - a change where a new form of matter is formed. • Also called chemical reaction. • Not phase changes • Ice is still water.

44. Physical Properties vs. Chemical Properties • Physical properties – properties that do not change the chemical nature of the matter. • Chemical properties – properties that do change the chemical nature of the matter.