# Edward Wen, PhD Matters and Measurement. Chemistry is about Everyday experience Why Cookies tastes different from Cookie Dough? Why Baking Powder or Baking.

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Edward Wen, PhDMatters and Measurement

Chemistry is about Everyday experienceWhy Cookies tastes different from Cookie Dough?Why Baking Powder or Baking Soda?Why using Aluminum Foil, not Paper Towel?What if the Temperature is set too high?Photo credit: itsnicethat.com*

Chapter OutlineClassification of mattersMeasurement, Metric system (SI)Scientific NotationSignificant figuresConversion factorDensity

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*In Your RoomEverything you can see, touch, smell or taste in your room is made of matter.Chemists study the differences in matter and how that relates to the structure of matter.

*What is Matter?Matter: anything that occupies space and has massMatter is actually composed of a lot of tiny little pieces: Atoms and Molecules

*Atoms and MoleculesAtoms: the tiny particles that make up all matter. Helium gas (for blimp) is made up of Helium atoms.

Molecules: In most substances, the atoms are joined together in units. Liquid water is made up of water molecules (2 Hydrogen atoms + 1 Oxygen atoms)

*Physical States of MattersMatter can be classified as solid, liquid or gas based on what properties it exhibits

State

Shape

Volume

Compress

Flow

Solid

Fixed

Fixed

No

No

Liquid

Indef.

Fixed

No

Yes

Gas

Indef.

Indef.

Yes

Yes

*Why different States of a Matter? Structure Determines Propertiesthe atoms or molecules have different structures in solids, liquid and gases

*SolidsParticles in a solid: packed close together and are fixed in positionthough they may vibrateIncompressibleretaining their shape and volumeUnable to flow

*LiquidsParticles are closely packed, but they have some ability to move around IncompressibleAble to flow, yet not to escape and expand to fill the container (not antigravity)

*GasesThe particles have complete freedom from each other (not sticky to each other)The particles are constantly flying around, bumping into each other and the containerThere is a lot of empty space between the particles (low density) Compressible Able to flow and Fill space (antigravity)

*Classifying Matter:

Sugar, Copper, Coke, Gasoline/Water

*Classification of Matter

*Pure substanceMatter that is composed of only one kind of piece.Solid: Salt, Sugar, Dry ice, Copper, Diamond

Liquid: Propane, distilled water (or Deionized water, DI water)

Gas: Helium gas (GOODYEAR blimp)

*Classifying Pure Substances: Elements and CompoundsElements: Substances which can not be broken down into simpler substances by chemical reactions. (A,B)

Compounds: Most substances are chemical combinations of elements. (C)Examples: Pure sugar, pure watercan be broken down into elementsProperties of the compound not related to the properties of the elements that compose it

*ElementsExample: Diamond (pure carbon), helium gas.116 known, 91 are found in natureothers are man-madeAbundance = percentage found in natureHydrogen: most abundant in the universe Oxygen: most abundant element (by mass) on earth and in the human bodySilicon: abundant on earth surface every sample of an element is made up of lots of identical atoms

*CompoundsComposed of elements in fixed percentageswater is 89% O & 11% Hbillions of known compoundsOrganic (sugar, glycerol) or inorganic (table salt)same elements can form more than one different compoundwater and hydrogen peroxide contain just hydrogen and oxygencarbohydrates all contain just C, H & O (sugar, starch, glucose)

*MixtureMatter that is composed of different kinds of pieces. Different samples may have the same pieces in different percentages. (D)

Examples:Solid: Flour, Brass (Copper and Zinc), RockLiquid: Salt water, soda, GasolineGas: air

*Classification of MixturesHomogeneous = composition is uniform throughout appears to be one thingevery piece of a sample has identical properties, though another sample with the same components may have different propertiessolutions (homogeneous mixtures): Air; Tap waterHeterogeneous = matter that is non-uniform throughout contains regions with different properties than other regions: gasoline mixed with water; Italian salad dressing

*What is a Measurement?Quantitative observationcomparison to an agreed upon standard

Every measurement has a number and a unit:77 Fahrenheit: Room temperature7.5 pounds: Average newborn body weight in the US:55 0.5 grams: amount of sugar in one can of Coca Cola

UNIT: what standard you are comparing your object to the number tells youwhat multiple of the standard the object measuresthe uncertainty in the measurement ()

*Some Standard Units in the Metric System

Quantity MeasuredName of UnitAbbreviationMassgramgLengthmetermVolumeliterLTimesecondssTemperatureKelvinK

*Related Units in the SI SystemAll units in the SI system are related to the standard unit by a power of 10 (exactly!)1 kg = 103 g 1 km = 103 m1 m = 102 cm

The power of 10 is indicated by a prefixThe prefixes are always the same, regardless of the standard unit

*Prefixes Used to Modify Standard Unitkilo = 1000 times base unit = 1031 kg = 1000 g = 103 gdeci = 0.1 times the base unit = 10-11 dL = 0.1 L = 10-1 L; 1 L = 10 dLcenti = 0.01 times the base unit = 10-21 cm = 0.01 m = 10-2 m; 1 m = 100 cmmilli = 0.001 times the base unit = 10-31 mg = 0.001 g = 10-3 g; 1 g = 1000 mgmicro = 10-6 times the base unit1 m = 10-6 m; 106 m = 1 mnano = 10-9 times the base unit1 nL = 10-9L; 109 nL = 1 L

*Common Prefixes in the SI System

PrefixSymbolDecimalEquivalentPower of 10mega-M1,000,000Base 106kilo-k 1,000Base 103deci-d 0.1Base 10-1centi-c 0.01Base 10-2milli-m 0.001Base 10-3micro-m or mc 0.000 001Base 10-6nano-n 0.000 000 001Base 10-9

*Standard Unit vs. PrefixesUsing meter as example:1 km = 1000 m = 103 m

1 g = 10 dm= 100 cm = 102 cm= 1000 mm= 103 mm= 1,000,000 m= 106 m= 1,000,000,000 nm= 109 nm

*LengthTwo-dimensional distance an object covers

SI unit: METER (abbreviation as m)About 3 inches longer than a yard1 m = 10-7 the distance from the North Pole to the Equator Commonly use centimeters (cm)1 m = 100 cm = 1.094 yard1 cm = 0.01 m = 10 mm1 inch = 2.54 cm (exactly)

*MassAmount of matter present in an object

SI unit: kilogram (kg)about 2 lbs. 3 oz.

Commonly measure mass in grams (g) or milligrams (mg)1 kg = 2.2046 pounds (1 lbs. = 0.45359)1 g = 1000 mg = 103 mg1 g = 0.001 kg = 10-3 kg

*VolumeAmount of three-dimensional space occupiedSI unit = cubic meter (m3)

Commonly measure solid volume in cubic centimeters (cm3)1 m3 = 106 cm3 1 cm3 = 10-6 m3 = 0.000001 m3

Commonly measure liquid or gas volume in milliliters (mL)1 gallon (gal) = 3.78 L = 3.78 103 mL1 L = 1 dm3 = 1000 mL = 103 mL 1 mL = 1 cm3 = 1 cc (cubic centimeter)

*Common Everyday Units and Their EXACT Conversions

11 cm1 inch (in)=2.54 cm1 mile=5280 feet (ft)1 foot= 12 in1 yard=3 ft

*Common Units and Their Equivalents

Volume1 liter (L)=1.057 quarts (qt)1 U.S. gallon (gal)=3.785 liters (L)

Mass1 kilogram (km)=2.205 pounds (lb)1 pound (lb)=453.59 grams (g)1 ounce (oz)=28.35 (g)

*UnitsAlways write every number with its associated unitAlways include units in your calculationsyou can do the same kind of operations on units as you can with numberscm cm = cm2cm + cm = cmcm cm = 1using units as a guide to problem solving

*Conversion FactorRelationships to Convert one unit of measurement to another: US dollar Canadian dollar, dollar cent

Conversion Factors: Relationships between two unitsBoth parts of the conversion factor have the same number of significant figuresConversion factors generated from equivalence statementse.g. 1 inch = 2.54 cm can giveor *

*How to Use Conversion FactorArrange conversion factors so starting unit cancelsArrange conversion factor so starting unit is on the bottom of the conversion factorunit 1unit 2unit 1unit 2x=Conversion Factor*

*We have been using the Conversion Factor ALL THE TIME! How are we converting #cents into #dollars? Why?From 1 dollar = 100 cents45,000 cents dollarcents450 dollarsx= 1 dollar 100 cents*

*Given:325 mgFind:? gConv. Fact.1 mg = 10-3 g Soln. Map:mg gConvert 325 mg to grams0.325 g*

Practice: How to set up Conversion?To convert 5.00 inches to cm, from 1 in = 2.54 cm (exact), which one of the two conversion factors should be used?

or*

Practice: Conversion among Units500 mg = ? g

3.78 L = ? mL

1.2 nm = ? m

* 8.0 in = ? m*

Scientific NotationVery Large vs. Very Small numbers:

The suns diameter is 1,392,000,000 m; An atoms diameter is 0.000 000 000 3 m

Scientific Notation: 1.392 x 109 m & 3 x 10-10 mthe sunsdiameter is1,392,000,000 m

*Scientific Notation (SN)Power of 10 (Math language): 10 x 10 = 100 100 = 102 (2nd power of 10)10 x 10 x 10 = 1,000 1,000 = 103 (3rd power of 10)

each Decimal Place in our number system represents a different power of 1024 = 2.4 x 101 = 2.4 x 101,000,000,000 (1 billion) = 1090.0000000001 (1/10 billionth ) = 10-10

Easily comparable by looking at the power of 10

*Exponents 10Ywhen the exponent on 10 (Y) is positive, the number is that many powers of 10 larger suns diameter = 1.392 x 109 m = 1,392,000,000 m

when Y is negative, the number is that many powers of 10 smalleravg. atoms diameter = 3 x 10-10 m = 0.0000000003 m1.23 x 105 > 4.56 x 1024.56 x 10-2 > 7.89 x 10-57.89 x 1010 > 1.23 x 1010

*Writing Numbers in SNBig numbers:

12,340,000

Small numbers:

0.00002341.234 x 1072.34 x 10-5

*Writing a Number in Standard Form1.234 x 10-6since exponent is -6, move the decimal point to the left 6 placesif you run out of digits, add zeros000 001.234

If the exponent > 1, add trailing zeros:1.234 x 10101.2340000000

0.000 001 23412,340,000,000

*Scientific calculators

*Inputting Scientific Notation into a Calculatorinput decimal part of the numberif negative press +/- key() on somepress EXP keyEE on some (maybe 2nd function)input exponent on 10press +/- key to change exponent to negative-1.23 x 10-3

*Significant Figures (Sig. Fig.)Definition: The non-place-holding digits in a reported measurementsome zeros in a written number are only there to help you locate the decimal point

What is Sig. Fig. for? the range of values to expect for repeated measurementsthe more significant figures there are in a measurement, the smaller the range of values is12.3 cmhas 3 sig. figs. and its range is12.2 to 12.4 cm0.1230 cmhas 4 sig. figs. and its range is0.1229 to 0.1231 cm

*Counting Significant FiguresAll non-zero digits are significant1.5 : 2 Sig. Fig.s

Interior zeros are significant1.05 : 3 Sig. Fig.s

Trailing zeros after a decimal point are significant1.050 : 4 Sig. Fig.s. Leading zeros are NOT significant0.001050 : 4 Sig. Fig.s Place-holding zeros= SN : 1.050 x 10-3

*Counting Significant Figures (Contd)4. Exact numbers has infinite () number of significant figures:example: 1 pound = 16 ounces1 kilogram = 1,000 grams = 1,000,000 milligrams1 water molecule contains 2 hydrogen atoms

5. Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation. Example: 150. has 3 sig. fig150 is ambiguous number1.50 x 102 has 3 sig. fig.

*ExampleCounting Sig. Fig. in a NumberHow many significant figures are in each of the following numbers?

0.00351.080

272.97 105

1 m = 1000 mm2 Sig. Fig. leading zeros not sig.4 Sig. Figs trailing & interior zeros sig.2 sig. Figs, all digits sig.3 Sig. Figs only decimal parts count sig.both 1 and 1000 are exact numbers. unlimited sig. figs.

Practice: How many Significant figures vs. Decimal places?2.2 cm

2.50 cm

2 sig. Figs; 1 decimal place

3 sig. Figs; 2 decimal places*

*Sig. Fig. in Multiplication/Division; Rounding vs. ZeroingWhen multiplying or dividing measurements with Sig. Fig., the result has the same number of significant figures as the measurement with the fewest number of significant figures Rounding 5.02 89,665 0.10 = 45.0118

5.892 6.10= 0.96590

3 SF5 SF2 SF2 SF4 SF3 SF3 SF= 45= 0.966

*Sig. Fig. in Multiplication/Division: Scientific notationOccasionally, scientific notation is needed to present results with proper significant figures.

5.89 6,103 = 35946.67 = 3.59 104

*Example: Determine the Correct Number of Sig. Fig.1.01 0.12 53.51 96 = 0.06755656.55 0.920 34.684 = 1.5 3 SF2 SF4 SF2 SFresult should have 2 Sig. Fig. 4 SF.3 SF.6 SF.result should have 3 Sig. Fig. = 0.068 = 1.50

*Sig. Fig. in Addition/Subtractionwhen adding or subtracting measurements with significant figures, the result has the same number of decimal places as the measurement with the fewest number of decimal places

5.74 + 0.823 +2.651= 9.214

4.865 - 3.965= 0.92 dp3 dp3 dp2 dp3 dp3 dp3 dp= 9.21= 0.900

*Example: Determine the Correct Number of Significant Figures0.987 + 125.1 1.22 = 124.8670.764 3.449 5.315 = -83 dp1 dp2 dpresult should have 1 dp 3 dp3 dp3 dpresult should have 3 dp = 124.9 = -8.000

*Sig. Fig. in Combined CalculationsDo and/or , then + and/or -3.489 5 .67 2.3 3 dp 3 Sig. Fig. 2 Sig. Fig. = 3.489 13 = -9.511 = -10 3 dp 0 dp 0 dp (2 sig. fig.) Parentheses (): Do calculation in () first, then the rest3.489 (5.67 2.3) 2 dp 1 dp

= 3.489 3.4 = 11.8628= 124 Sig. Fig. 2 Sig. Fig. 2 Sig. Fig.

Practice: Calculation with Proper Significant Figuresa. 12.99 + 2.09 x 1.921 =

b. 2.00 x 3.5 - 1.000 =

12.99 + 4.01 = 17.00

7.0 1.000 = 6.0

15.00/3.75 = 4.00

6.7 8.8 = 59*

*How to solve Unit Conversion ProblemsWrite down Given Amount and UnitWrite down what you want to Find and UnitWrite down needed Conversion Factors or EquationsDesign a Solution Map for the Problemorder Conversions to cancel previous units orarrange Equation so Find amount is isolated. Example: from Equation A = b c to solve for b

*Solution Map for Unit Conversion

5) Check the Answer to see if its Reasonablecorrect size and unitApply the Steps in the Solution Map check that units cancel properlymultiply terms across the top and divide by each bottom termExample:

*

Example: Unit Conversion

*Given:7.8 kmFind:? miConv. Fact.1 mi = 5280 ft1 foot = 12 in1 in = 2.54 cm (exact)Soln. Map:km miAlternative Route:Convert 7.8 km to mileskmmcminftmi*

*Apply the Solution Map:Given:7.8 kmFind:? miConv. Fact.1 mi = 5280 ft1 foot = 12 in1 in = 2.54 cm (exact)Soln. Map:km miAlternative Route:Convert 7.8 km to miles= 4.84692 mi = 4.8 mi Sig. Figs. & Round:*

*TemperatureTemperature is a measure of the average kinetic energy of the molecules in a sampleNot all molecules have in a sample the same amount of kinetic energya higher temperature means a larger average kinetic energy

*Fahrenheit Temperature ScaleTwo reference points:Freezing point of concentrated saltwater (0F)Average body temperature (100F)more accurate measure now set average body temperature at 98.6FRoom temperature is about 75F

*Celsius Temperature ScaleTwo reference points:Freezing point of distilled water (0C)Boiling point of distilled water (100C)more reproducible standardsmost commonly used in scienceRoom temperature is about 25C

*Fahrenheit vs. Celsiusa...