1. Chapter 1 Elements and Measurements You are responsible for all sections in this chapter

2. Chemistry and the Elements

3. Periods: 7 horizontal rows. Groups: 18 vertical columns. International standard: 1-18 US system: 1A-8A, 1B-8B

4. Elements and the Periodic Table

5. Some Chemical Properties of the Elements Physical Properties: Characteristics that do not involve a change in a sample’s chemical makeup. Chemical Properties: Characteristics that do involve a change in a sample’s chemical makeup.

6. The Metric System (SI) The metric system or SI (international system) is a decimal system based on 10. used in most of the world. used everywhere by scientists. 6

7. Experimentation and Measurement All other units are derived from these fundamental units Système Internationale d´Unités

9. Measuring Mass Mass: Amount of matter in an object. Matter: Describes anything with a physical presence—anything you can touch, taste, or smell. Weight: Measures the force with which gravity pulls on an object.

10. Measuring Temperature T F = 1.8 T C + 32 T C = (T F – 32) 1.8 K = °C + 273.15

11. Scientific Notation is used to write very large or very small numbers. for the width of a human hair of 0.000 008 m is written 8 x 10 -6 m. of a large number such as 4 500 000 s is written 4.5 x 10 6 s. 11

12. Accuracy, Precision, and Significant Figures Significant figures: The number of meaningful digits in a measured or calculated quantity. They come from uncertainty in any measurement. Generally the last digit in a reported measurement is uncertain (estimated). Exact numbers and relationships (7 days in a week, 30 students in a class, etc.) effectively have an infinite number of significant figures.

13. Accuracy, Precision, and Significant Figures length = 1.74 cm 01243 cm 1.7 cm < length < 1.8 cm

14. Accuracy, Precision, and Significant Figures Rules for counting significant figures (left-to-right): 1.Zeros in the middle of a number are like any other digit; they are always significant. 4.803 cm 4 sf 2.Rules for counting significant figures (left-to- right): Zero at the beginning of a number are not significant (placeholders). 0.00661 g 3 sfor 6.61 x 10 -3 g

15. Accuracy, Precision, and Significant Figures Rules for counting significant figures (left-to-right): 3.Zeros at the end of a number and after the decimal point are always significant. 55.220 K 5 sf 4.Zeros at the end of a number and after the decimal point may or may not be significant. 34,2000 ? SF

16. Rounding Numbers If the first digit you remove is 5 and there are more nonzero digits following, round up. 5.664 525 = 5.665 If the digit you remove is a 5 with nothing following, round down. 5.664 525 = 5.664 52

17. Multiplication and Division When multiplying or dividing the final answer must have the same number of significant figures as the measurement with the fewest significant figures. Example: 110.5 x 0.048 = 5.304 = 5.3 (rounded) 4 SF 2 SF calculator 2 SF 17

18. Addition and Subtraction When adding or subtracting the final answer must have the same number of decimal places as the measurement with the fewest decimal places. 25.2 one decimal place + 1.34 two decimal places 26.54calculated answer 26.5 final answer with one decimal place 18

19. Calculations: Converting from One Unit to Another Dimensional analysis: A method that uses a conversion factor to convert a quantity expressed in one unit to an equivalent quantity in a different unit. Conversion factor: States the relationship between two different units. original quantity x conversion factor = equivalent quantity

20. Conversion Factors A conversion factor is obtained from an equality. Equality: 1 in. = 2.54 cm written as a fraction (ratio) with a numerator and denominator. inverted to give two conversion factors for every equality. 1 in. and 2.54 cm 2.54 cm 1 in. 20

21. Conversion Factors in a Problem A conversion factor may be obtained from information in a word problem. is written for that problem only. Example : The cost of one gallon (1 gal) of gas is $4.29. 1 gallon of gasand $4.29 $4.291 gallon of gas 21

22. Example: How many ounces are in 1.0 kg? How many in 3 in 1.5 m 3 22

23. Examples If your pace on a treadmill is 65 meters per minute, how many minutes will it take for you to walk a distance of 7500 feet?

24. Derived Units: Measuring Density density = volume mass solids- cm 3 liquids- mL gases- L Typical volume units

25. Examples Osmium is a very dense metal. What is its density in g/cm 3 if 0.11 lb of osmium has a volume of 2.22 ml? The density of octane, a component of gasoline, is 0.702 g/mL. What is the mass, in kg, of 875 mL of octane?