1 / 29

John Dalton, 1766-1844

Antoine Lavoisier, 1743-1794. Joseph Priestly, 1766-1844. Marie Curie, 1867-1934. Dmitri Mendeleev, 1834-1907. John Dalton, 1766-1844. What is Matter? Matter : Anything that occupies space and has mass Energy: Ability to do work, accomplish a change Physical States of Matter

mary-gomez
Download Presentation

John Dalton, 1766-1844

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

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

  2. What is Matter? Matter: Anything that occupies space and has mass Energy: Ability to do work, accomplish a change 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 Each substance has a unique set of properties identifying it from other substances. Intensive Properties: Properties that do not depend on quantity of substance Examples: boiling point, density Extensive Properties: Properties that depend on or vary with the quantity of substance Examples: mass, volume

  4. 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

  5. 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

  6. 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: O2 • Compound: Pure substance consisting of 2 or more kinds of atoms (heteroatomic molecules) • Example: CO2

  7. 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)

  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. Density: mass divided by volume; d = m/v Specific gravity: density of a substance relative to the density of water

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

  11. Metric System: Meter, gram, liter

  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 cm3 • 1 ml = 1 cm3 (1 cc)

  14. Unit of Energy • Joule: Basic unit of energy • calorie: amount of heat energy needed to increase temperature of 1 g of water by 1oC • 1 cal = 4 joules • Nutritional calorie = 1000 calories = 1 kcal = 1 Calorie • Units of Temperature Fahrenheit: -459oF (absolute zero) - 212oF (water boils) Celsius: -273oC (absolute zero) - 100oC (water boils) Kelvin: 0K (absolute zero) - 373 K (water boils)

  15. Different Temperature Scales

  16. Converting Celsius and Fahrenheit: oC = 5/9 (Fo - 32) oF = 9/5 (oC) +32 Converting Celsius and Kelvin: oC = K - 273 K = oC + 273 Scientific Notation and Significant Figures Scientific notation: a shorthand way of representing very small or very large numbers Examples: 3 x 102, 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 102 450 • 3.72 x 105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 104 300 = .00045 = 4.5 x 10-4 .0005 = 3.00 x 102 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 onlyif 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. Practice: 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 20o Celsius. D. There are 6 hats in the closet. Inexact Exact 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. Practice Using and Converting Units in Calculations • Sample calculation: Convert 125m to yards. • Write down the known or given quantity (number and unit) • 125 m • Leave some blank space and set the known quantity equal to the unit of the unknown quantity • 125 m = yards • Multiply the known quantity by the factor(s) necessary to cancel out the units of the known quantity and generate the units of the unknown quantity • 125 m x 1.09 yards/1 m = yards

  28. Once the desired units have been achieved, do the necessary arithmetic to produce the final answer • 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

  29. Accuracy vs. Precision Error: difference between true value and our measurement Accuracy: degree of agreement between true value and measured value Uncertainty: degree of doubt in a measurement Precision: degree of agreement between replicated measurements

More Related