Chapter 3Scientific Measurement

Section 3.1The Importance of

MeasurementOBJECTIVES:Distinguish between quantitative and

qualitative measurements.

Section 3.1The Importance of

MeasurementOBJECTIVES: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 3.2Uncertainty in Measurements

OBJECTIVES:Distinguish among the accuracy,

precision, and error of a measurement.

Section 3.2Uncertainty 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 valuex 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 3.3International System of

UnitsOBJECTIVES:List SI units of measurement and

common prefixes.

Section 3.3International System of

UnitsOBJECTIVES: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

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)

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

Volume Measuring Instruments

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

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

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

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

Section 3.4Density

OBJECTIVES:Calculate the density of an object

from experimental data.

Section 3.4Density

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/cm3, (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 oC = 1 g/cm3

Formula

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 =

Section 3.5Temperature

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 oC) and

boiling point (100 oC) of water as references

Divided into 100 equal intervals, or degrees Celsius

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

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