Chemistry Unit One

1. Chemistry Unit One Scientific Method, Observation, Experimentation,Theory vs. Law, Metric Measurement, Significant Figures, Scientific Notation, % Error and Dimensional Analysis!!!

2. Scientific Method • A problem solving strategy. • Used to solve scientific as well as everyday problems.

3. Steps of the Scientific Method • Define the Problem. Usually in the form of a question. • Gather Information. Become an expert in the subject. • Form a hypothesis. An educated guess. Must be testable. • Test hypothesis. With experimentation and/or observation. • State Conclusion. An interpretation of experimental results.

4. Theory vs. Law • Law • A “rule” of nature. • Describes how nature behaves. • Answers the question “what?” • Theory • Explains the behavior of a natural event. • Predicts the results of further experiments. • Answers the question “why?”

5. Similar Charges Repel Objects with mass attract one another Energy is neither created nor destroyed in a chemical reaction. Organisms change over time. Matter bends the space around it causing objects to move toward one another. Beneficial mutations cause organisms to vary over time and become more suitable for survival. Theory vs. Law LAWS THEORIES

6. Observations • Pieces of information gathered using the senses.

7. Observations • Use as many senses as is sensible! • Avoid vague terms. A ‘bad’ estimation is better than a general term. • Do not identify … describe. • Do not “liken”… describe. Note the excessive use of DESCRIBE

8. 1.The dessert is delicious. 2. The dessert is colorful. 3. The dessert is like a torte. 4. The dessert is tiny. 5. The dessert is chocolate moose cake. 1. The dessert is thick and creamy. 2. The dessert is several shades of light brown. 3. The dessert is round and flat. 4. The dessert is an 8 inch diameter. 5. The dessert is made of chocolate cream and fudge. Observations

9. SI Base Units • Systeme International d’Unites

10. Non-SI Units used in Chem.

11. Derived Units Used in Chem.

12. Metric Prefixes • Increase or decrease the size of a metric unit.

13. Metric Measurements Every measurement contains error! • Every piece of equipment has a certain range of uncertainty. (ex. +/- 0.2 g) This will be given on the piece of equipment. • Every measurement contains estimation equal to a maximum of one half of the smallest unit on the piece of equipment. (ex. If a graduated cylinder is marked off in whole mL, you must estimate volumes in between mL lines.) Your measurement should go one decimal place beyond what is known.

14. Measurement Plastic Ruler Accuracy: Object is more than 12cm and less than 12.5cm. Precision: The smallest marking represents 0.5cm, so you estimate to the nearest 0.1cm. The length of the object would be accurately and precisely recorded as 12.4cm. Uncertainty: 12.4 +/- 0.25cm.

15. Measurement Metal Ruler: Accuracy: Object is more than 12.3cm and just less than 12.4cm. Precision: The smallest marking represents 0.1cm, so you estimate to the nearest 0.01cm. The length of the object would be accurately and precisely recorded as 12.35cm. Uncertainty: 12.35 +/- 0.05cm

16. Significant Digits • The certain digits and one estimated digit of each measurement are significant. Remember! Every time you make a measurement, you record all of the certain digits and one estimated digit. 200.54 g

17. Rules for Sig. Figs • Non zeros are always significant. • Zeros between non zeros are significant. • Zeros at the end of significant digits following a decimal point are significant. *They show precision in measurement. 4) Place keeper zeros are NOT significant. • Zeros preceding significant digits. • Zeros following significant digits without a decimal point.

18. Try These Examples 7.05940 Final zero significant (follows decimal point) 6 significant digits 0.00135 Leading zeros Not significant (place keepers) 3 significant digits 20,400 Final zeros Not significant (place keepers – no decimal) 3 significant digits

19. Adding and Subtracting Round to the fewest number of decimal places given in problem. (Can only have ONE estimated digit in final answer) Multiplying and Dividing Round to the fewest number of significant digits given in the problem. Sig Figs and Calculations

20. 17.20 (.01) 4.137 (.001) + 26.6 (.1) 47.937 Least significant number is reported to the tenths, so round final answer to the tenths. 47.9 14.3 (3 sig figs) 1.0200 (5 sig figs) x 0.005 (1 sig fig) 0.07293 Fewest number of sig figs is one, so round the final answer to one sig fig. 0.07 Sample Problems

21. Scientific Notation • In chemistry, we work with very large and very small numbers. Number of particles in a mole = 602200000000000000000000 Mass of an electron = 0.000000000000000000000000000000911kg

22. Scientific Notation To make these numbers easier to work with, we put them into scientific notation. • Rewrite the significant digits as a number greater than one, but less than 10. • Count the number of places you had to move the decimal to complete step 1. • Write this number of decimal places as an exponent.

23. Sample Problems 602 200 000 000 000 000 000 000 • Rewrite as a number greater than one but less than 10. 6.022 • Count the number of places the decimal moved. (left) 23 places • Write that number as an exponent. 6.022 x 1023

24. How about another one? 0.000000000000000000000000000000911 • Rewrite as a number greater than one but less than 10. 9.11 • Count the number of places the decimal moved. (right) 31 • Write that number as an exponent. 9.11 x 10-31

25. How about the other direction? Speed of light in a vacuum is 3.00 x 108 m/s Move the decimal 8 places to the right. 300 000 000 m/s

26. One last sample Atomic Mass Unit 1.66054 x 10-27 kg Move the decimal 27 places to the left. 0.00000000000000000000000000166054 kg

27. Warm-up 1/30/12 • Why are significant figures important, especially in scientific data? Objective • TSWBAT: use various methods of recording, calculating and displaying scientific data

28. Agenda • Attendance • Warm-up • Discussion of Measurement Lab • Notes and Practice of Dimensional Analysis • Notes and Practice of Error Types • Notes and Practice of graphing/density • Wrap-up and HW

29. Dimensional AnalysisUsed to convert from one metric unit into another. • Write known value as a fraction with a one in the denominator. Include units. • Look for a conversion from the unit you begin with to the one that you wish to get to. • Write the conversion factor so that the given unit cancels out and the desired unit remains in the numerator . • Multiply straight across the top and bottom. • Divide the top product by the bottom product.

30. Sample Problem 450.0 cm = _____ m 1) 450.0 cmWrite problem as 1 a fraction. • 1 m Write conversion factor • 100 cm so units will cancel. • Multiply numerator and denominator factors, • then divide. 3) 450.0 cm x 1 m = 450.0 = 4.5 m 1 100cm 100

31. Another Sample Conversion 6.325 L = _____ mL 1) 6.325 LWrite problem as 1 a fraction. • 1 L Write conversion factor • 1000 mL so units will cancel. • Multiply numerator and denominator factors, • then divide. 3) 6.325 L x 1000 mL = 6325 = 6325 mL 1 1 L 1

32. Practice Time • Dimensional Analysis Worksheet

33. Error…

34. Accuracy How close your measurement is to an accepted value for that measurement. Depends on how carefully the measurement was made. Precision-2 Meanings The repeatability (2.2, 2.3, 2.2) of a measurement. The number of significant digits (2.2 vs. 2.20475) in the measurement. Depends on the equipment used. Accuracy vs. Precision

35. Accuracy • When throwing darts, a bulls eye is considered accurate. • In an experiment, you could also achieve accuracy with several measurements that when averaged closely match the accepted value.

36. Precision • When throwing darts, precision occurs when you consistently hit the same spot. • In lab, getting data points that are close together shows precision. • A reading with many digits beyond the decimal also shows precision.

37. Accuracy vs. Precision Who or What is to blame…. • If values are precise but not accurate? • If the values are not precise but only a few are accurate? How can precision and accuracy help in trying to solve an experimental problem?

38. Types of error • Human • Ex. Reading data before or after the time you should have recorded it • Method • Ex. Loss of product when performing the experiment because of poorly written process • Instrumental • Instrument was not calibrated correctly

39. Calculating Percent Error To determine how close your measurement or calculation is to the accepted value, we will use percent error. Why percent??? If your calculated density is “off” by 0.1 g/mL… is that a lot of error or a little error??? A percent will let you know!

40. Calculating Percent Error The formula we use to calculate % error is: measured value – accepted value x 100 accepted value

41. If you were calculating the density of gold, the accepted value is 19.2 g/ml. If you’re off by 0.5g/mL 18.7 – 19.2 x 100 = 19.2 -2.60 % error If you were calculating the density of aluminum, the accepted value is 2.7 g/mL. If you’re off by 0.5g/mL 2.2 – 2.7 x 100 = 2.7 -18 % error Comparing % Errors

42. Percent Error • So, what does a negative % error mean??? Your measured or calculated answer is lower than the accepted value. • Therefore, a positive % error means that your answer is higher than the accepted value.

43. Try this one. C’mon, try it! • You calculate the solubility of potassium chlorate at 20oC to be 19.4 g/100mL H2O. • The accepted value is 14.0 g/100mL H2O. • Calculate your % error. % error = measured – accepted x 100 accepted 19.4 – 14.0 x 100 = 38.6 % 14.0

44. Graphing Results in Science What makes a good graph?

45. Why do we use graphs in science? • Graphing is a pictorial way of representing relationships between various quantities, parameters, or measurable variables in nature. A graph basically summarizes how one quantity changes if another quantity that is related to it also changes.

46. Step 1 • Always give your graph a title in the following form: "The dependence of (your dependent variable) on (your independent variable). 

47. Step 2 • The x-axis of a graph is always your independent variable and the y-axis is the dependent variable.

48. Step 3 • Always label the x and y axes and give units.

49. Step 4 • Always make a line graph unless otherwise specified.

50. Step 5 • Do not connect the dots on your graph – use a line of best fit unless told otherwise.