Presentation on theme: "Chapter 3 Scientific Measurement Pioneer High School Mr. David Norton."— Presentation transcript:
Chapter 3 Scientific Measurement Pioneer High School Mr. David Norton
Section 3.1 The Importance of Measurement OBJECTIVES: Distinguish between quantitative and qualitative measurements.
Section 3.1 The Importance of Measurement OBJECTIVES: Convert measurements to scientific notation.
Measurements Qualitative measurements - words Quantitative measurements – involves numbers (quantities) Depends on reliability of instrument Depends on care with which it is read Scientific Notation Coefficient raised to power of 10
Working with Scientific Notation Multiplication Multiply the coefficients, add the exponents Division Divide the coefficients, subtract the denominator exponent from numerator exponent
Working with Scientific Notation Before adding or subtracting in scientific notation, the exponents must be the same Calculators will take care of this Addition Line up decimal; add as usual the coefficients; exponent stays the same
Working with Scientific Notation Subtraction Line up decimal; subtract coefficients as usual; exponent remains the same
Section 3.2 Uncertainty in Measurements OBJECTIVES: Distinguish among the accuracy, precision, and error of a measurement.
Section 3.2 Uncertainty in Measurements OBJECTIVES: Identify the number of significant figures in a measurement, and in the result of a calculation.
Uncertainty in Measurements Need to make 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) Fig. 3.4, page 54
Uncertainty in Measurements Accepted value – correct value based on reliable references Experimental value – the value measured in the lab Error – the difference between the accepted and experimental values
Uncertainty in Measurements Error = accepted – experimental Can be positive or negative Percent error = the absolute value of the error divided by the accepted value, times 100% | error | accepted value x 100% error =
Significant Figures (sig. figs.) Significant figures in a measurement include all of the digits that are known, plus a last digit that is estimated. Note Fig. 3.6, page 56 Rules for counting sig. figs.? Zeroes are the problem East Coast / West Coast method
Counting Significant Fig. Sample 3-1, page 58 Rounding Decide how many sig. figs. Needed Round, counting from the left Less than 5? Drop it. 5 or greater? Increase by 1 Sample 3-2, page 59
Sig. fig. calculations Addition and Subtraction The answer should be rounded to the same number of decimal places as the least number in the problem Sample 3-3, page 60
Sig. Fig. calculations Multiplication and Division Round the answer to the same number of significant figures as the least number in the measurement Sample 3-4, page 61
Section 3.3 International System of Units OBJECTIVES: List SI units of measurement and common prefixes.
Section 3.3 International System of Units OBJECTIVES: Distinguish between the mass and weight of an object.
International System of Units The number is only part of the answer; it also need UNITS Depends upon units that serve as a reference standard The standards of measurement used in science are those of the Metric System
International System of Units Metric system is now revised as the International System of Units (SI), as of 1960 Simplicity and based on 10 or multiples of 10 7 base units Table 3.1, page 63
International System of Units Sometimes, non-SI units are used Liter, Celsius, calorie Some are derived units Made by joining other units Speed (miles/hour) Density (grams/mL)
Length In SI, the basic unit of length is the meter (m) Length is the distance between two objects – measured with ruler We make use of prefixes for units larger or smaller Table 3.2, page 64
Volume changes? Volume of any solid, liquid, or gas will change with temperature Much more prominent for GASES Therefore, measuring instruments are calibrated for a specific temperature, usually 20 o C, which is about normal room temperature
Units of Mass Mass is a measure of the quantity of matter Weight is a force that measures the pull by gravity- it changes with location Mass is constant, regardless of location
Section 3.4 Density OBJECTIVES: Calculate the density of an object from experimental data.
Section 3.4 Density OBJECTIVES: List some useful application of the measurement of specific gravity.
Density Which is heavier- lead or feathers? It depends upon the amount of the material A truckload of feathers is heavier than a small pellet of lead The relationship here is between mass and volume- called Density
Density The formula for density is: mass volume Common units are g/mL, or possibly g/cm 3, (or g/L for gas) Density is a physical property, and does not depend upon sample size Density =
Things related to density Note Table 3.7, page 69 for the density of corn oil and water What happens when corn oil and water are mixed? Why? Will lead float?
Density and Temperature What happens to density as the temperature increases? Mass remains the same Most substances increase in volume as temperature increases Thus, density generally decreases as the temperature increases
Density and water Water is an important exception Over certain temperatures, the volume of water increases as the temperature decreases Does ice float in liquid water? Why? Sample 3-5, page 71
Specific Gravity A comparison of the density of an object to a reference standard (which is usually water) at the same temperature Water density at 4 o C = 1 g/cm 3
Formula D of substance (g/cm 3 ) D of water (g/cm 3 ) Note there are no units left, since they cancel each other Measured with a hydrometer – p.72 Uses? Tests urine, antifreeze, battery Specific gravity =
Section 3.5 Temperature OBJECTIVES: Convert between the Celsius and Kelvin temperature scales.
Temperature Heat moves from warmer object to the cooler object Glass of iced tea gets colder? Remember that most substances expand with a temp. increase? Basis for thermometers
Temperature scales Celsius scale- named after a Swedish astronomer Uses the freezing point(0 o C) and boiling point (100 o C) of water as references Divided into 100 equal intervals, or degrees Celsius
Temperature scales Kelvin scale (or absolute scale) Named after Lord Kelvin K = o C + 273 A change of one degree Kelvin is the same as a change of one degree Celsius No degree sign is used
Temperature scales Water freezes at 273 K Water boils at 373 K 0 K is called absolute zero, and equals –273 o C Fig. 3.19, page 75 Sample 3-6, page 75