1. John Dalton, 1766-1844 Marie Curie, 1867-1934 Antoine Lavoisier, 1743-1794 Joseph Priestly, 1733-1804 Dmitri Mendeleev, 1834-1907

2. What is Matter? Matter: Anything that occupies space and has mass Physical States of Matter Gas: Indefinite volume, indefinite shape, particles far away from each other Liquid: Definite volume, indefinite shape, particles closer together than in gas Solid: Definite volume, definite shape, particles close to each other

3. Properties of Matter Property: Characteristic of a substance Physical Properties: Properties of matter that can be observed without changing the composition or identity of a substance Example: Size, physical state Chemical Properties: Properties that matter demonstrates when attempts are made to change it into new substances, as a result of chemical reactions Example: Burning, rusting

4. Changes in Matter Physical Changes: Changes matter undergoes without changing composition Example: Melting ice; crushing rock Chemical Changes: Changes matter undergoes that involve changes in composition; a conversion of reactants to products Example: Burning match; fruit ripening

5. Classifying Matter Pure substance: Matter that has only 1 component; constant composition and fixed properties Example: water, sugar Element: Pure substance consisting of only 1 kind of atom (homoatomic molecule) Example: O 2 Compound: Pure substance consisting of 2 or more kinds of atoms (heteroatomic molecules) Example: CO 2

6. Mixture: A combination of 2 or more pure substances, with each retaining its own identity; variable composition and variable properties Example: sugar-water Homogenous matter: Matter that has the same properties throughout the sample Heterogenous matter: Matter with properties that differ throughout the sample Solution: A homogenous mixture of 2 or more substances (sugar- water, air)

7. Classification of Matter

8. Measurement Systems Measurement: Determination of dimensions, capacity, quantity or extent of something; represented by both a number and a unit Examples: Mass, length, volume, energy, density, specific gravity, temperature Mass vs. Weight Mass: A measurement of the amount of matter in an object Weight: A measurement of the gravitational force acting on an object

9. English System Units: Inch, foot, pound, quart

10. Each type of measurement has a base unit. The Metric System of Units

11. Other units are related to the base unit by a power of 10. The prefix of the unit name indicates if the unit is larger or smaller than the base unit.

12. Unit of Length Meter = basic unit of length, approximately 1 yard 1 meter = 1.09 yards Kilometer = 1000 larger than a meter Centimeter = 1/100 of a meter 100 cm = 1 meter Millimeter = 1/1000 of a meter 1000 mm = 1 meter

13. Unit of Mass Gram: basic unit of mass 454 grams = 1 pound Kilogram: 1000 times larger than a gram 1 Kg = 2.2 pounds Milligram: 1/1000 of a gram Unit of Volume Liter: basic unit of volume 1 Liter = 1.06 quarts 1 Liter = 10 cm x 10 cm x 10 cm 1 liter = 1000 cm 3 1 ml = 1 cm 3 (1 cc)

14. Units of Temperature Fahrenheit: -459 o F (absolute zero) - 212 o F (water boils) Celsius: -273 o C (absolute zero) - 100 o C (water boils) Kelvin: 0K (absolute zero) - 373 K (water boils)

15. Different Temperature Scales

16. Converting Celsius and Fahrenheit: o C = 5/9 (F o - 32) o F = 9/5 ( o C) +32 Converting Celsius and Kelvin: o C = K - 273K = o C + 273 Scientific Notation and Significant Figures Scientific notation: a shorthand way of representing very small or very large numbers Examples: 3 x 10 2, 2.5 x 10 -4

17. The exponent is the number of places the decimal must be moved from its original position in the number to its position when the number is written in scientific notation If the exponent is positive, move the decimal to the right of the standard position Example: 4.50 x 10 2  450 3.72 x 10 5  372,000 If the exponent is negative, move the decimal to the left of the standard position Example: 9.2 x 10 -3 .0092

18. Practice with Scientific Notation 50,000 = 5.0 x 10 4 300 =.00045 = 4.5 x 10 -4.0005 = 3.00 x 10 2 5 x 10 -4

19. Significant Figures Significant Figures: Numbers in a measurement that reflect the certainty of the measurement, plus one number representing an estimate Example: 3.27cm Rules for Determining Significance:  All nonzero digits are significant  Zeroes between significant digits are significant Example: 205 has 3 significant digits 1,006 has 10,004 has 4 sig. figs. 5 sig. figs.

20.  Leading zeroes are not significant Example: 0.025 has 2 significant digits 0.000459 has 3 significant digits 0.0000003645  Trailing zeroes are significant only if there is a decimal point in the number Examples: 1.00 has 3 significant figures 2.0 has 2 significant digits 20 has 1500 1.500 4 sig. figs. 1 sig. fig. 2 sig. figs. 4 sig. figs.

21. Calculations and Significant Figures Answers obtained by calculations cannot contain more certainty (significant figures) than the least certain measurement used in the calculation Multiplication/Division: The answers from these calculations must contain the same number of significant figures as the quantity with the fewest significant figures used in the calculation Example: 4.95 x 12.10 = 59.895 Round to how many sig. figs.? Final answer: 3 59.9

22. Addition/Subtraction: The answers from these calculations must contain the same number of places to the right of the decimal point as the quantity in the calculation that has the fewest number of places to the right of the decimal Example: 1.9 + 18.65 = 20.55 How many sig. figs. required? Final answer: Rounding Off Rounding off: a way reducing the number of significant digits to follow the above rules 1 20.6

23. Rules of Rounding Off: Determine the appropriate number of significant figures; any and all digits after this one will be dropped. If the number to be dropped is 5 or greater, all the nonsignificant figures are dropped and the last significant figure is increased by 1 If the number to be dropped is less than 5, all nonsignificant figures are dropped and the last significant figure remains unchanged Example: 4.287 (with the appropriate number of sig. figs. determined to be 2) 4.287  4.3

24. We only use significant figures when dealing with inexact numbers Exact (counted) numbers: numbers determined by definition or counting Example: 60 minutes per hour, 12 items = 1 dozen Inexact (measured) numbers: numbers determined by measurement, by using a measuring device Example: height = 1.5 meters, time elapsed = 2 minutes

25. Classify each of the following as an exact or a inexact number. A. A field is 100 meters long. B. There are 12 inches in 1 foot. C. The current temperature is 20 o Celsius. D. There are 6 hats in the closet. Practice: Inexact Exact

26. Calculating Percentages percent = “per hundred” % = (part/total) x 100 Example: 50 students in a class, 10 are left-handed. What percentage of students are lefties? % lefties = (# lefties/total students) x 100 = 10/50 x 100 =.2 x 100 = 20%

27. Conversion factor: A term that converts a quantity in one unit to a quantity in another unit. Conversion factors are usually written as equalities. 2.21 lb = 1 kg 2.21 lb 1 kg or 1 kg 2.21 lb To use them, they must be written as fractions. original quantity original quantity conversion factor desired quantity desired quantity x = Problem Solving Using Conversion Factors

28. Make sure all unwanted units cancel Units are treated like numbers Example: Convert 130 lb into kilograms 130 lbx conversion factor = ? kg original quantity desired quantity

29. or 2.21 lb 1 kg 2.21 lb 1 kg 2.21 lb 1 kg 2.21 lb 130 lbx =59 kg The bottom conversion factor has the original unit in the denominator and the desired unit in the numerator. The unwanted unit lb cancels. The desired unit kg does not cancel.

30. Example How many grams of aspirin are in a 325-mg tablet? Step [1] Identify the original quantity and the desired quantity, including units. original quantity desired quantity 325 mg? g

31. Step [2] Write out the conversion factor(s) needed to solve the problem. 1 g = 1000 mg This can be written as two possible fractions: 1000 mg 1g 1000 mg or Choose this factor to cancel the unwanted unit, mg.

32. Step [3] Set up and solve the problem. 325 mgx 1 g 1000 mg = 0.325 g Step [4] Write the answer with the correct number of significant figures. Unwanted unit cancels. 325 mg 0.325 g 3 sig. figures

33. Practice Using and Converting Units in Calculations Practice: Convert 125m to yards. 125 x 1.09 yards /1 = 136.25 yards Determine appropriate amount of sig. figs. and round accordingly Fewest sig. figs. in original problem is 3 (from 125), so final answer is 136 yards

34. Example: How many liters is in 1.0 pint? 2 pints = 1 quart1.06 quarts = 1 liter Two conversion factors are needed: 1.0 pint original quantity ? L desired quantity Solving a Problem Using Two or More Conversion Factors 2 pt 1 qt 2 pt or 1.06 qt 1 L 1.06 qt or First, cancel pt. Then, cancel qt. x conversion factor =

35. Set up the problem and solve: 1.0 pt x 1 qt 2 pt x 1 L 1.06 qt = 0.47 L 1.0 pt 2 sig. figures 0.47 L 2 sig. figures

36. Density and Specific Gravity density = mass (g) volume (mL) Density: A physical property that relates the mass of a substance to its volume. To convert volume (mL) to mass (g): To convert mass (g) to volume (mL): mLx g =g gx g = density inverse of density

37. Solving Problems with Density Example: If the density of acetic acid is 1.05 g/mL, what is the volume of 5.0 grams of acetic acid? 5.0 g original quantity ? mL desired quantity Density is the conversion factor, and can be written two ways: 1.05 g 1 mL 1.05 g Choose the inverse density to cancel the unwanted unit, g. x conversion factor =

38. Set up and solve the problem: 5.0 gx 1 mL 1.05 g =4.8 mL Write the final answer with the correct number of significant figures. 5.0 g 4.8 mL Unwanted unit cancels.

39. Specific Gravity Specific gravity: A quantity that compares the density of a substance with the density of water at the same temperature. specific gravity= density of a substance (g/mL) density of water (g/mL) The units of the numerator (g/mL) cancel the units of the denominator (g/mL). The specific gravity of a substance is equal to its density, but contains no units.