1 / 37

Chapter 3 “ Scientific Measurement ”

Chapter 3 “ Scientific Measurement ”. Mt. Hebron High School Chemistry. Units of Measurement. When you measure things, what units do you typically use?: Feet, inches, miles, pounds, gallons.. This is referred to as the “English System.”

normajohnson
Download Presentation

Chapter 3 “ Scientific Measurement ”

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. Chapter 3“Scientific Measurement” Mt. Hebron High School Chemistry

  2. Units of Measurement When you measure things, what units do you typically use?: Feet, inches, miles, pounds, gallons.. This is referred to as the “English System.” In science, the metric system is used because it is a universal language that anyone can understand regardless of what country they are in.

  3. The Fundamental SI Units(Le Système International, SI)

  4. Other units and Variables used in Chemistry. Area Volume Force Pressure Energy Power Voltage Frequency Electric Charge Meter squared Cubic meter Newton Pascal Joule Watt Volt Hertz Coulomb

  5. The Metric System The units in the metric system are known as SI units (International System of Units). Sometimes there are variables that are larger or smaller than the base unit, so we use metric prefixes to make measurements more convenient. Examples of Base units… meter, liter, gram, pascal,…

  6. Scientific Notation • In science it is very common to work with really large or really small numbers. • 602,000,000,000,000,000,000,000 • 6.02 x 1023 atoms = 1 mole of a substance

  7. Scientific Notation • In science it is very common to work with really large or really small numbers. • 602,000,000,000,000,000,000,000 • 6.02 x 1023 atoms = 1 mole of a substance • _____ x 10 (# of decimal places) (+….moved left) Number between (-…..moved right) 1 and 9

  8. Rules for Scientific Notation • If the original number is larger than 1, then you move the decimal to the left and your exponent will be positive. 1. 4135 = 4.135 x 103 * Moved the decimal point 3 places to the left. • If the original number is less than 1, then you move the decimal point to the right and your exponent will be negative. 2. .00179 = 1.79 x 10-3 *Moved decimal point 3 places to the right.

  9. If you take a number out of scientific notation look at the exponent first. • Negative Exponent = move to the left • Positive Exponent = move to the right

  10. A. 43,812g = B. .000000943m = C. 5.66 x 10-3 L = D. 7.254 x 106 g = 4.3812 x 104 9.43 x 10-7 .00566 7,254,000 Try These:

  11. Uncertainity in Measurements • Measurements are an important aspect in Chemistry. • Three common measurements that we typically make are for Mass, Volume, and Temperature.

  12. When you make a measurement, write down all the numbers you can. WHY?? • Measuring instruments are never completely free from flaws. • Measuring always involves some estimation. (Digital-Balance or Scale-Graduated Cylinder)

  13. How do you know when a measurement is reliable?

  14. Accuracy, Precision, and Error • It is necessary to make good, reliable measurements in the lab • Accuracy – how close a measurement is to the true value • Precision – how close the measurements are to each other (reproducibility)

  15. Precision and Accuracy Precise, but not accurate Neither accurate nor precise Precise AND accurate

  16. Why Is there Uncertainty? • Measurements are performed with instruments, and no instrument can read to an infinite number of decimal places • Which of the balances shown has the greatest uncertainty in measurement?

  17. Percents and Percent Error • How do you calculate a percentage? • Many times in Chemistry your results will need to be written as a percent. • Percent = (Actual / Total) x 100

  18. Percent Error • Percent Error measures how accurate you data really is. • Percent Error = measured value – accepted value x 100 accepted value • This answer should be really low.

  19. Measurement Activity

  20. Length of Clothespin cm Length of Pen cm Volume of cyl. 1 Volume of cyl. 2 Volume of cyl. 3 Volume of cyl. 4 Block 1 volume Block 2 volume Volume of rock Volume of stopper Mass of cup Mass of scissors Mass of water Density of stopper Temperature of liquid 8.2cm 15.2cm 54mL 10.2mL 83mL 35.5mL 27cm3 54cm3 2.5mL 10.1mL 354.56g 32.53g 10.0g 1.37g/mL 23.5o Percent ErrorCalculate the percent error for each measurement in the activity.

  21. SI Prefixes Commonly Used

  22. 1 Kilometer equals how many meters? • 1000 meters because the Kilometer prefix represents a thousand. So since there is 1 Kilometer, one times 1000 is 1000 meters.

  23. How many meters in 1000 centimeters? • 10 meters because the centi prefix represents one hundreth of a meter. • We have to divide 1000 centimeters by 100 to get our answer. • Easy trick to remember…up the chart you divide or down the chart you multiply the number of steps you move.

  24. 550 millimeters as meters. 3.5 seconds as milliseconds 1.6 Kilograms to grams 2500 milligrams to Kilograms 4.00 centimeters to micrometers 2800 decimoles to moles 6.1 amperes to milliamperes 3 Kilograms to milligrams .550m 3500ms 1600g .0025 kg 40,000 um 280 mol 6100 ma 3,000,000mg Try These

  25. Significant Figures in Measurements • Significant figures in a measurement include all of the digits that are known, plus one more digit that is estimated. • Measurements must be reported to the correct number of significant figures.

  26. Significant Figures Which measurement is the best? What is the measured value? What is the measured value? What is the measured value?

  27. Rules for Significant Figures • Zeroes are only significant when they are not a place holder. A. 1040 Place Holder B. 0.0026701m not significant C. 19.0550 Kg All significant D. 1,809,000 L not significant

  28. Rules for Significant Figures • Multiplication or Division : The number with the fewest significant figures determines how many are in the answer. 3.05m x 2.10m x 0.75m = 4.80375 m3 = 4.8 m3

  29. Rules for Significant Figures • Addition or Subtraction : Look for the fewest decimal places to determine the number of significant figures. 6.41g + 10.2g + 11g + 3.667g = 31.277g = 31g 5.84L + 24.1L + 16.044L = 45.984L = 46.0L

  30. Density • Which is heavier: 1 lb. of lead or 1 lb. of feathers? • Most people compare masses, but in chemistry we not only compare masses, but densities as well. • Density is a mass-to-volume ratio. • Density = mass / volume • Density depends on the composition of the substance, not how much is present.

  31. Density = mass / volume

  32. Density • Density can be calculated 2 ways: • 1. Mathmatically • 2. Slope • Mathmatically---- D = m/v • Slope---- Line Graph of Mass vs. Volume slope = rise / run = mass / volume

  33. Density Problems

  34. Density Graphs

  35. Density Lab

More Related