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**Chapter 3 Scientific Measurement**

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**Section 3.1 The Importance of Measurement**

OBJECTIVES: Distinguish between quantitative and qualitative measurements.

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**Section 3.1 The Importance of Measurement**

OBJECTIVES: Convert measurements to scientific notation.

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**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

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**Working with Scientific Notation**

Multiplication Multiply the coefficients, add the exponents Division Divide the coefficients, subtract the denominator exponent from numerator exponent

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**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

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**Working with Scientific Notation**

Subtraction Line up decimal; subtract coefficients as usual; exponent remains the same

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**Section 3.2 Uncertainty in Measurements**

OBJECTIVES: Distinguish among the accuracy, precision, and error of a measurement.

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**Section 3.2 Uncertainty in Measurements**

OBJECTIVES: Identify the number of significant figures in a measurement, and in the result of a calculation.

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**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

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**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

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**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 % error = x 100%

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**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

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**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

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**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

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**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

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**Section 3.3 International System of Units**

OBJECTIVES: List SI units of measurement and common prefixes.

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**Section 3.3 International System of Units**

OBJECTIVES: Distinguish between the mass and weight of an object.

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**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

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**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

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**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)

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**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

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**Common prefixes Kilo (k) = 1000 (one thousand)**

Deci (d) = 1/10 (one tenth) Centi (c) = 1/100 (one hundredth) Milli (m) = 1/1000 (one thousandth) Micro () = (one millionth) Nano (n) = (one billionth)

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**Volume The space occupied by any sample of matter**

Calculated for a solid by multiplying the length x width x height SI unit = cubic meter (m3) Everyday unit = Liter (L), which is non-SI

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**Volume Measuring Instruments**

Graduated cylinders Pipet Buret Volumetric Flask Syringe Fig. 3.12, page 66

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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 oC, which is about normal room temperature

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**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

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Working with Mass The SI unit of mass is the kilogram (kg), even though a more convenient unit is the gram Measuring instrument is the balance scale

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**Section 3.4 Density OBJECTIVES:**

Calculate the density of an object from experimental data.

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**Section 3.4 Density OBJECTIVES:**

List some useful application of the measurement of specific gravity.

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**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

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**Density The formula for density is: mass volume**

Common units are g/mL, or possibly g/cm3, (or g/L for gas) Density is a physical property, and does not depend upon sample size Density =

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**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?

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**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

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**Density and water Sample 3-5, page 71 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

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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 oC = 1 g/cm3

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**Formula Note there are no units left, since they cancel each other**

D of substance (g/cm3) D of water (g/cm3) Note there are no units left, since they cancel each other Measured with a hydrometer – p.72 Uses? Tests urine, antifreeze, battery Specific gravity =

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**Section 3.5 Temperature OBJECTIVES:**

Convert between the Celsius and Kelvin temperature scales.

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**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

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**Temperature scales Celsius scale- named after a Swedish astronomer**

Uses the freezing point(0 oC) and boiling point (100 oC) of water as references Divided into 100 equal intervals, or degrees Celsius

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**Temperature scales Kelvin scale (or absolute scale)**

Named after Lord Kelvin K = oC + 273 A change of one degree Kelvin is the same as a change of one degree Celsius No degree sign is used

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**Temperature scales Water freezes at 273 K Water boils at 373 K**

0 K is called absolute zero, and equals –273 oC Fig. 3.19, page 75 Sample 3-6, page 75

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