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Chapter 3 Scientific Measurement

Chapter 3 Scientific Measurement. Ms. Riggins Lawndale High School. Chapter 3.1 – Measurements and Their Uncertainty. Measurement - a quantity that has both a number and a unit. It is important to be able to make measurements and decide whether a measurement is correct. Scientific Notation.

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Chapter 3 Scientific Measurement

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  1. Chapter 3Scientific Measurement Ms. Riggins Lawndale High School

  2. Chapter 3.1 – Measurements and Their Uncertainty • Measurement - a quantity that has both a number and a unit • It is important to be able to make measurements and decide whether a measurement is correct

  3. Scientific Notation We will be working with rather small and large numbers in our science class. For example, Avogadro's number is 602,214,000,000,000,000,000,000 and a human hair is .0002 meters in diameter. It will be much easier writing values in scientific notation, rather than standard form.

  4. Scientific Notation • Scientific notation – a given number written as the product of two numbers • a coefficient and 10 raised to a power

  5. Rules for Scientific Notation • The first number is always between 1 and 9.9999... • 2. Multiply the first number by 10 raised to an exponent. • Positive exponents are large numbers • Negative exponents are small numbers (decimals)

  6. Practice • 66 • 222 • 0.00046 • .08 • 602,214,000,000,000,000,000,000 • 0.0000000546 • 56,938 • 0.000000000000144

  7. Significant Figures Significant Figures are VERY important in chemistry. Each recorded measurement has a certain number of significant digits. Any digit that is actually measured or estimated will be considered significant. Placeholders, or digits that have not been measured or estimated, are not considered significant. (Example: 0.00009)

  8. Rules for Significant Figures • Digits from 1 to 9 are always significant. (Example: 458kg has 3 sig. fig.) 2. Zeros between two other significant digits are always significant. (Example: 5057L has 4 sig. fig.) 3. Zeros to the left of the decimal point are not significant. (Example: 500 has 1 sig. fig.) 4. Zeros to the right of the decimal point are significant, unless they are placeholders. (Example: 0.007 has 1 sig. fig., but 0.070 has 2 sig. fig.)

  9. Practice • 123 • 9.8000 x 104 • 40,506 • 22 • 0.07080 • 98,000 • 0.00700 • 0.000000000000144

  10. More Practice • 0.05730 • 0.00073 • 8765 • 40.007 • 143 • 0.074 • 1.072 • 8.750 x 10-2

  11. Significant Figures in Calculations • When adding and subtracting, your answers should be rounded to the same number of decimal places as the measurement with the least number of decimal places (line up your decimal places) Practice • 12.52 + 349.0 + 8.24 = • 61.2 + 9.35 + 8.6 = • 9.44 – 2.11 = • 34.61 – 17.3 =

  12. Significant Figures in Calculations • When multiplying and dividing, your answers should be rounded to the same number of significant figures as the measurement with the least number of significant figures Practice • 7.55 x 0.34 = • 2.10 x 0.70 = • 2.4526 / 8.4 = • 8432 / 12.5 =

  13. Chapter 3.2 - International System of Units • The five basic SI units used by chemists are… Meter (length) Kilogram (mass) Kelvin (temperature) Second (time) Mole (amount of substance)

  14. Units of Length Common metric units of length include the centimeter, meter, and kilometer

  15. Units of Volume Common metric units of volume include the liter, milliliter, and cubic centimeter

  16. Units of Mass • Common metric units of mass include the kilogram and gram • Weight is a force that measures the pull on a given mass by gravity

  17. Units of Temperature • Measures how hot or cold an object is • Measures the kinetic energy of particles • The two commonly used units for temperature are Celsius and Kelvin • The zero point on the Kelvin scale is called Absolute Zero (-273ºC) K = °C +273

  18. Units of Energy Joule and Calorie are two common units of energy 1 calorie = 4.184 Joules

  19. Chapter 3.4 - Density Density = mass/volume The density of a substance generally decreases as its temperature increases Density is an intensive property that depends only on the composition of a substance, not on the size of the sample.

  20. Calculations with Density 1. What is the volume of a coin with a mass of 14grams and a density of 10.5g/cm3? 2. What is the volume of 14.8grams of boron if the density of the sample is 2.34g/cm3? 3. What is the volume of a 4.62gram sample of mercury if it’s density is 13.5g/cm3?

  21. Chapter 3.3 – Unit ConversionCommonly Used Metric PrefixesMUST MEMORIZE (PAGE 74)

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