1 Introduction: Matter and Measurement Chapter 1.

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Introduction: Matter and Introduction: Matter and MeasurementMeasurement

Chapter 1Chapter 1

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The Study of ChemistryThe Study of Chemistry

Chemistry is the study of the properties and behavior of matter.

Matter – anything that occupies space and has mass.

What is chemistry?What is chemistry?

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Classification of MatterClassification of Matter

States of Matter

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States of Matter

Gas

Liquid

Solid

5


States of Matter

Shape Volume

Gas

Liquid

Solid

6


States of Matter

Shape Volume

Gas indefinite

Liquid

Solid

7


States of Matter

Shape Volume

Gas indefinite indefinite

Liquid

Solid

8


States of Matter

Shape Volume


Liquid indefinite

Solid

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States of Matter

Shape Volume


Liquid indefinite definite

Solid

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States of Matter

Shape Volume



Solid definite

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States of Matter

Shape Volume



Solid definite definite

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The basic difference between these states is the distance between the “bodies.”

• Gas – bodies are far apart and in rapid motion.• Liquid – bodies closer together, but still able to

move past each other.• Solid – bodies are closer still and are now held

in place in a definite arrangement.

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Mixture – combination of two or more substances in which each substance retains its own chemical identity.• Homogeneous mixture – composition of this mixture

is consistent throughout.

• Heterogeneous mixture – composition of this mixture varies throughout the mixture.

Pure Substances and Mixtures

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It is also possible for a homogeneous substance to It is also possible for a homogeneous substance to be composed of a single substance – be composed of a single substance – pure pure substance.substance.

• Element – A substance that can not be separated into simpler substances by chemical means.

• Compound – A substance composed of two or more elements united chemically in definite proportions.


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The smallest unit of an element is an atom.

Atom – the smallest unit of an element that retains a substances chemical activity.


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• Mixtures can be separated by physical means.– Filtration.

– Chromatography.

– Distillation.

Separation of Mixtures

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Classification of MatterClassification of MatterSeparation of Mixtures

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• There are 114 elements known.• Each element is given a unique chemical

symbol (one or two letters).– Carbon C, nitrogen N, titanium Ti.

– Notice that the two letter symbols are always capital letter then lower case letter because:

• CO – carbon and oxygen.

• Co – element cobalt.

Elements

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• Formed by combining elements.• The proportions of elements in compounds are

the same irrespective of how the compound was formed.

Law of Constant Composition (or Law of Definite Proportions):– The composition of a pure compound is always the

same, regardless of its source.

Compounds

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Properties of MatterProperties of Matter

Physical Property (Change) – A property that can be measured without changing the identity of the substance.

Example: melting point, boiling point, color, odor, density

Physical changes do not result in a change of composition.

Physical and Chemical Changes

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Intensive properties – independent of sample size.

Extensive properties - depends on the quantity of the sample.


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Chemical change (chemical reaction) – the transformation of a substance into a chemically different substance.– When pure hydrogen and pure oxygen react

completely, they form pure water.


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Units of MeasurementUnits of Measurement

• There are two types of units:– fundamental (or base) units;

– derived units.

• There are 7 base units in the SI system.• Derived units are obtained from the 7 base SI units.

SI Units

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m/ssecondsmeters

time of unitsdistance of units

velocity of Units

• There are two types of units:– fundamental (or base) units;

– derived units.

• There are 7 base units in the SI system.• Derived units are obtained from the 7 base SI units.• Example:

SI Units

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SI Units


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SI Units

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Mass is the measure of the amount of material in an Mass is the measure of the amount of material in an object.object.

– This is not the same as weight which is dependant on This is not the same as weight which is dependant on gravity.gravity.

Units of MeasurementUnits of MeasurementMass

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Units of MeasurementUnits of MeasurementTemperature

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Kelvin ScaleUsed in science.Same temperature increment as Celsius scale.Lowest temperature possible (absolute zero) is zero Kelvin. Absolute zero: 0 K = -273.15oC.

Celsius ScaleAlso used in science.Water freezes at 0oC and boils at 100oC.To convert: K = oC + 273.15.

Fahrenheit ScaleNot generally used in science.Water freezes at 32oF and boils at 212oF.

Temperature

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Converting between Celsius and Fahrenheit

32-F95

C 32C59

F

Temperature

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• The units for volume are given by (units of length)3.– i.e., SI unit for volume is 1 m3.

• A more common volume unit is the liter (L)– 1 L = 1 dm3 = 1000 cm3 = 1000 mL.

• We usually use 1 mL = 1 cm3.

Volume

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Density – mass per unit volume of an object.

volume

massDensity

Density

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• All scientific measures are subject to error.• These errors are reflected in the number of figures

reported for the measurement.• These errors are also reflected in the observation

that two successive measures of the same quantity are different.

Uncertainty in MeasurementUncertainty in Measurement

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• Measurements that are close to the “correct” value are accurate.

• Measurements which are close to each other are precise.

• Measurements can be– accurate and precise

– precise but inaccurate

– neither accurate nor precise


Precision and Accuracy

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Precision and Accuracy


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• The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device.

• All the figures known with certainty plus one extra figure are called significant figures.

• In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).

Significant Figures

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• Non-zero numbers are always significant.• Zeros between non-zero numbers are always

significant.• Zeros before the first non-zero digit are not

significant. Zeros at the end of the number after a decimal place are significant.

• Zeros at the end of a number before a decimal place are ambiguous. For this course we will consider these to be significant.– Example – so for this class, the number 10,300 has 5

significant figures.

Significant Figures

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• Multiplication / DivisionMultiplication / Division– The result must have the same number of significant The result must have the same number of significant

figures as the least accurately determined datafigures as the least accurately determined data

Example: Example:

12.512 (5 sig. fig.) 12.512 (5 sig. fig.)

5.1 (2 sig. fig.)5.1 (2 sig. fig.)

12.512 x 5.1 = 64 12.512 x 5.1 = 64

Answer has only 2 significant figuresAnswer has only 2 significant figures

Significant Figures

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• Addition / Subtraction.Addition / Subtraction.– The result must have the same number of digits to the The result must have the same number of digits to the

right of the decimal point as the least accurately right of the decimal point as the least accurately determined data.determined data.

Example:Example:15.152 (5 sig. fig., 3 digits to the right),15.152 (5 sig. fig., 3 digits to the right),1.76 (3 sig. fig., 2 digits to the right),1.76 (3 sig. fig., 2 digits to the right),7.1 (2 sig. fig., 1 digit to the right).7.1 (2 sig. fig., 1 digit to the right).

15.152 + 1.76 + 7.1 = 24.015.152 + 1.76 + 7.1 = 24.0

24.0 (3 sig. fig., but only 1 digit to the right of the decimal 24.0 (3 sig. fig., but only 1 digit to the right of the decimal point)point)

Significant Figures

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• If the leftmost digit to be removed is less than 5, the preceding number is left unchanged.“Round down.”

• If the leftmost digit to be removed is 5 or greater, the preceding number is increased by 1.“Round up.”

Rounding rules

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Dimensional AnalysisDimensional Analysis• In dimensional analysis always ask three questions:• What data are we given?• What quantity do we need?• What conversion factors are available to take us from

what we are given to what we need?

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Dimensional AnalysisDimensional Analysis

• Method of calculation using a conversion factor.

inches

footor

foot

inches

footinches

12

11

1

121

112

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Example: we want to convert the distance 8 in. to feet.

(12in = 1 ft)

in

ftin

12

18

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Example: we want to convert the distance 8 in. to feet.

(12in = 1 ft)

ftin

ftin 67.0

12

18

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ProblemConvert the quantity from 2.3 x 10-8 cm to nanometers (nm)

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First we will need to determine the conversion factors

Centimeter (cm) Meter (m)

Meter (m) Nanometer (nm)

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Centimeter (cm) Meter (m)

Meter (m) Nanometer (nm)

Or

1 cm = 0.01 m

1 x 10-9 m = 1 nm

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1 cm = 0.01 m

1 x 10-9 m = 1 nm

Now, we need to setup the equation where the cm cancels and nm is left.

m

nm

cm

mcm8103.2

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1 cm = 0.01 m

1 x 10-9 m = 1 nm

Now, fill-in the value that corresponds with the unit and solve the equation.

m

nm

cm

mcm

98

101

1

1

01.0103.2

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nmm

nm

cm

mcm 23.0

101

1

1

01.0103.2

98

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ProblemConvert the quantity from 31,820 mi2 to square meters (m2)

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Mile (mi) kilometer (km)

kilometer (km) meter (m)

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Mile (mi) kilometer (km)

kilometer (km) meter (km)

Or

1 mile = 1.6093km

1000m = 1 km

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Now, we need to setup the equation where the mi cancels and m is left.

1 mile = 1.6093km 1000m = 1 km

km

m

mi

kmmi 2820,31

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Now, we need to setup the equation where the mi cancels and m is left.

1 mile = 1.6093km 1000m = 1 km

Notice, that the units do not cancel, each conversion factor must be “squared”.

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2820,31km

m

mi

kmmi

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22

2

1

1000

1

6093.1820,31

km

m

mi

kmmi

60



2

26

2

22

1

101

1

5898.2820,31

km

m

mi

kmmi

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2102

26

2

22 102407.8

1

101

1

5898.2820,31 m

km

m

mi

kmmi

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ProblemConvert the quantity from 14 m/s to miles per hour (mi/hr).

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Determine the conversion factors

Meter (m) Kilometer (km) Kilometer(km) Mile(mi)

Seconds (s) Minutes (min) Minutes(min) Hours (hr)

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Determine the conversion factors

Meter (m) Kilometer (km) Kilometer(km) Mile(mi)

Seconds (s) Minutes (min) Minutes(min) Hours (hr)

Or

1 mile = 1.6093 km 1000m = 1 km

60 sec = 1 min 60 min = 1 hr

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1 mile = 1.6093 km 1000m = 1 km

60 sec = 1 min 60 min = 1 hr

hr

min

min

s

km

mi

m

km/14 sm

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1 mile = 1.6093 km 1000m = 1 km

60 sec = 1 min 60 min = 1 hr

hr1

min60

min1

s60

km6093.1

mi1

m1000

km1/14 sm

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1 mile = 1.6093 km 1000m = 1 km

60 sec = 1 min 60 min = 1 hr

hrmi

sm

/31

hr1

min60

min1

s60

km6093.1

mi1

m1000

km1/14

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End of Chapter ProblemsEnd of Chapter Problems

1.2, 1.16a, 1.18, 1.20, 1.26, 1.36,

1.38, 1.44, 1.52, 1.63, 1.67

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1 Introduction: Matter and Measurement Chapter 1.

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