2. INTRODUCTION Matter And Measurement

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) - A set of tested hypotheses that give an overall explanation of some natural phenomenon Natural Law - The same observation applies to many different systems - Example: Law of Conservation of Mass

5. Law vs. Theory A law summarizes what happens A theory (model) is an attempt to explain why it happens.

6. Part 1 - number Part 2 - scale (unit) Examples: 20 grams 20 grams 6.63 x 10 -34 Joule seconds Measurement - quantitative observation consisting of 2 parts consisting of 2 parts Nature of Measurement

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

8. SI Units

9. SI Prefixes Common to Chemistry PrefixUnit Abbr.Exponent MegaM 10 6 Kilok 10 3 Decid 10 -1 Centic 10 -2 Millim 10 -3 Micro  10 -6 Nanon 10 -9 Picop 10 -12

10. 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

11. Accuracy refers to the agreement of a particular value with the true (known) value. Precision refers to the degree of agreement among several measurements made in the same manner. (aka – reproducibility) Neither accurate nor precise Precise but not accurate Precise AND accurate Precision and Accuracy

12. 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.

13. Rules for Counting Significant Figures 1.If the number contains a decimal, count from right to left until only zeros or no digits remain. Examples: 20.05 grams 4 sig figs 7.2000 meters 5 sig figs 0.0017 grams 2 sig figs

14. 2. If the number does not contain a decimal, count from left to right until only zeros or no digits remain. Examples: 255 meters 3 sig figs 1,000 kilograms 1 sig fig

15. 3.Exact numbers have an infinite number of significant figures. 1 inch = 2.54 cm, exactly

16. 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 10 3 s  3 sig figs 0.0054 cm  2 sig figs 3,200,000  2 sig figs Sig Fig Practice #1

17. Rules for Significnt Figures in Mathematical Operations Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement.Addition and Subtraction: The number of decimal places in the result equals the number of decimal places in the least precise measurement. 6.8 + 11.934 = 18.734  18.7 (3 sig figs)

18. Sig Fig Practice #2 3.24 m + 7.0 m CalculationCalculator says:Answer 10.24 m 10.2 m 100.0 g - 23.73 g 76.27 g 76.3 g 0.02 cm + 2.371 cm 2.391 cm 2.39 cm 713.1 L - 3.872 L 709.228 L709.2 L 1818.2 lb + 3.37 lb1821.57 lb 1821.6 lb 2.030 mL - 1.870 mL 0.16 mL 0.160 mL

19. Rules for Significant Figures in Mathematical Operations Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in the calculation.Multiplication and Division: # sig figs in the result equals the number in the least precise measurement used in the calculation. 6.38 x 2.0 = 12.76  13 (2 sig figs)

20. Sig Fig Practice #3 3.24 m x 7.0 m CalculationCalculator says:Answer 22.68 m 2 23 m 2 100.0 g ÷ 23.7 cm 3 4.219409283 g/cm 3 4.22 g/cm 3 0.02 cm x 2.371 cm 0.04742 cm 2 0.05 cm 2 710 m ÷ 3.0 s 236.6666667 m/s240 m/s 1818.2 lb x 3.23 ft5872.786 lb·ft 5870 lb·ft 1.030 g ÷ 2.87 mL 2.9561 g/mL2.96 g/mL

21. Converting Celsius to Kelvin Kelvin =  C + 273°C = Kelvin - 273

22. Extensive properties Intensive properties Volume Mass Energy Content (think Calories!) depend on the amount of matter that is present. do not depend on the amount of matter present. Melting point Boiling point Density Properties of Matter

23. Three Phases

24. Solid Solid – definite volume and shape; particles packed in fixed positions. Liquid Liquid – definite volume but indefinite shape; particles close together but not in fixed positions Gas 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. Phase Differences

25. Classification of Matter

26. Separation of a Mixture The constituents of the mixture retain their identity and may be separated by physical means.

27. The components of dyes such as ink may be separated by paper chromatography. Separation of a Mixture The components of dyes such as ink may be separated by paper chromatography.

28. Separation of a Mixture Distillation

29. Matter Matter Mixtures: a) Homogeneous (Solutions) b) Heterogeneous Pure Substances Compounds Elements Elements Atoms NucleusElectrons ProtonsNeutrons Quarks Quarks Organization of Matter

30. Water  Hydrogen + Oxygen H 2 O  H 2 + O 2 Reactant  Products Compounds must be separated by chemical means. With the application of electricity, water can be separated into its elements Separation of a Compound Separation of a Compound The Electrolysis of water

31. Dimensional Analysis -aka: factor label unit cancellation fence-post -provides a systematic way of solving numerical problems Set-up: Given Desired Units___ 1 Units to Eliminate

32. Dimensional Analysis Examples 115 lbs = ______ g 115 lbs 453.6 g  5.22 x 10 4 g 1 1 lb

33. Useful Conversions 1 mi = 1.6093 km 1 lb = 453.59 g 1 in = 2.54 cm 0 F = (9/5) 0 C + 32 1 L = 1.0567 qt 1 mL = 1cm 1 kg = 2.2046 lb 0 C = (5/9)( 0 F – 32 ) 3