CHAPTER 1* Introduction: Matter and Measurement

1. CHAPTER 1*Introduction: Matter and Measurement Suh Kwon

2. 1.1 – The Study of Chemistry • The Molecular Perspective of Chemistry • Matter = physical material of the universe that has mass and occupies space • Element = substance that cannot be separated into simpler substances by chemical means • Atom = almost infinitesimally small building blocks of matter • Molecules = chemical combination of two ore more atoms

3. 1.2 – Classifications of Matter • States of Matter

4. Pure Substances and Mixtures • Pure Substance(referred to as a substance)= matter that has a fixed composition and distinct properties • Classified as either .. • Elements (composed of only one kind of atom) • Compounds (composed of two ore more elements) • Mixture = combinations of two or more substances in which each substance retains its own chemical identity and its own properties; compositions vary • Heterogeneous = do not have the same composition, properties, and appearance throughout the mixtures (Ex: sand, rocks, and wood) • Homogeneous = uniform throughout the mixtures(Ex: air, such as nitrogen, oxygen, etc)

5. Separation of Mixtures • Filtration (for heterogeneous mixtures) • Ex: to separate iron filings from gold ones, use a magnet to attract the iron • Distillation (for homogeneous mixtures) • Ex: to separate salt from water, boil the solution; thus, water will evaporate while the salt is left behind because water has a much lower boiling point than table salt • Chromatography  https://www.crimescene.com/store/bmz_cache/b/b3ac10ad9721f4d3399d6f4b82111806.image.220x223.jpg*

6. 1.3 – Properties of Matter • Physical properties = w/o changing identity and composition; (color, odor, density, melting point, boiling point, and hardness) • Ex: When water evaporates, it changes from liquid to gas; however, its composition does not change; it is still water • Chemical properties = substance that may change or react to form other substances • Ex: Flammability, which is the ability to burn a substance in the presence of oxygen, transforms one substance into a chemically different one • Intensive properties = does not depend on the amount of the sample being examined; (temperature, melting point, and density) • Extensive properties = depend on the quantity and amount of the sample; (measurements of mass and volume)

7. 1.4 – Units of Measurement • SI Units

8. Derived SI Units • Density = mass volume  Practice Problem!!: ** Calculate the density of mercury if 1.00 x 102 g occupies a volume of 7.36 cm3. (Answer) Density = mass = 1.00 x 102 g = 13.6 g/cm3 volume 7.36 cm3

9. 1.5 – Uncertainty in Measurement • Precision and Accuracy • Precision = a measure of how closely individual measurements agree with one another • Accuracy = how closely individual measurements agree with the correct, or “true” value http://celebrating200years.noaa.gov/magazine/tct/accuracy_vs_precision_556.jpg*

10. Significant Figures • Guidelines to determine the number of sig. figures: • Nonzero digits are always significant (214= THREE significant figures) • Zeros between nonzero digits are always significant (1004 = FOUR significant figures) • Zeros at the beginning of a number are never significant (0.01 = ONE significant figure) • Zeros that fall both at the end of a number and after the decimal point are always significant (4.0 = TWO significant figures)

11. Significant Figures in Calculations • Multiplication and Division: the result must be reported with the same number of significant figures as the measurement with the fewest significant figures • Addition and Subtraction: the result cannot have more digits to the right of the decimal point than any of the original numbers

12. 1.6 – Dimensional Analysis • Conversion factor • given unit x desired unit = desired unit given unit  Example: *Converting 8.00 meters to inches 8.00 m x 100 cm x 1 in. = 315 inches 1 m 2.54 cm