chapter 11 matter and measurement chapter 1. 2 the study of chemistry what is chemistry? chemistry...

Chapter 1 1

Matter and MeasurementMatter and Measurement

Chapter 1Chapter 1

Chapter 1 2

The Study of ChemistryThe Study of Chemistry

What is Chemistry?What is Chemistry?

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

Matter – anything that occupies space and has mass.

Chapter 1 3

Classification of MatterClassification of Matter

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.

Chapter 1 4


Chapter 1 5


Chapter 1 6


Pure Substances and MixturesPure Substances and Mixtures

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.• Solution (Air, gasoline)

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

Chapter 1 7


Separation of MixturesSeparation of MixturesMixtures can be separated by physical means.

– Filtration.

– Chromatography.

– Distillation.

Chapter 1 8


Separation of MixturesSeparation of Mixtures

Chapter 1 9


Pure Substances and MixturesPure Substances and MixturesIt is also possible for a homogeneous substance to be

composed of a single substance – pure substance.• Element – A substance that can not be separated into

simpler substances by chemical means.• Atom – the smallest unit of an element that retains a

substances chemical activity.

Chapter 1 10


ElementsElements• 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

Chapter 1 11


Pure Substances and MixturesPure Substances and MixturesIt is also possible for a homogeneous substance to

be composed of a single substance – pure substance.

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

Chapter 1 12


CompoundsCompounds• 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.

Chapter 1 13

Properties of MatterProperties of Matter

Physical and Chemical ChangesPhysical and Chemical ChangesPhysical Property – A property that can be measured

without changing the identity of the substance.

Example: color, odor, density

Chapter 1 14


Physical and Chemical ChangesPhysical and Chemical Changes

Intensive properties – independent of sample size.

Extensive properties - depends on the quantity of the sample (sample size).

Chapter 1 15

Units of MeasurementUnits of Measurement

DensityDensityDensity – mass per unit volume of an object.

volume

massDensity

Chapter 1 16



Physical change – the change in the physical properties of a substance.

– Physical appearance changes, but the substances identity does not.

Water (ice) Water (liquid)

Chapter 1 17



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.

Chapter 1 18



Chapter 1 19


m/ssecondsmeters

time of unitsdistance of units

velocity of Units

SI UnitsSI 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:

Chapter 1 20

SI UnitsSI Units


Chapter 1 21


SI UnitsSI Units

Chapter 1 22


TemperatureTemperature

Chapter 1 23


TemperatureTemperatureKelvin Scale

Used 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.

Chapter 1 24


TemperatureTemperature

Converting between Celsius and Fahrenheit

32-F95

C

32C59

F

Chapter 1 25


VolumeVolume• 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.

Chapter 1 26

MassMassMass 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.


Chapter 1 27

• 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

Chapter 1 28

Precision and AccuracyPrecision and Accuracy• 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.


Chapter 1 29

Precision and AccuracyPrecision and Accuracy


Chapter 1 30


Significant FiguresSignificant Figures• The number of digits reported in a measurement

reflect the accuracy of the measurement and the precision of the measuring device.

• The last digit to the right in a number is taken to be inexact.

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

Chapter 1 31


Significant FiguresSignificant Figures• 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 book, it a decimal point is used the zeros are significant.

– 10,300 has 3 significant figures.

– 10,300. has 5 significant figures.

• Physical constants are “infinitely” significant.

Chapter 1 32


Significant FiguresSignificant Figures• Multiplication / DivisionMultiplication / Division

– The result must have the same number of significant figures as The result must have the same number of significant figures as the least accurately determined datathe 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

Chapter 1 33


Significant FiguresSignificant Figures• Addition / Subtraction.Addition / Subtraction.

– The result must have the same number of digits to the right of The result must have the same number of digits to the right of the decimal point as the least accurately determined data.the decimal point as the least accurately 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 = 24.0.15.152 + 1.76 + 7.1 = 24.0.

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

Chapter 1 34


Rounding rules• 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.”

Chapter 1 35

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?

Chapter 1 36


• Method of calculation using a conversion factor.

inches

footor

foot

inches

footinches

12

11

1

121

112

Chapter 1 37


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

(12in = 1 ft)

in

ftin

12

18

Chapter 1 38


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

(12in = 1 ft)

ftin

ftin 67.0

12

18

Chapter 1 39


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

First we will need to determine the conversion factors

Centimeter (cm) Meter (m)

Meter (m) Nanometer (nm)

Or

1 cm = 0.01 m

1 x 10-9 m = 1 nm

Chapter 1 40



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

Chapter 1 41



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

Chapter 1 42



nmm

nm

cm

mcm 23.0

101

1

1

01.0103.2

98

Chapter 1 43


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

First we will need to determine the conversion factors

Mile (mi) Meter (m)

Meter (m) kilometer (km)

Or

1 mile = 1.6093km

1000m = 1 km

Chapter 1 44




1 mile = 1.6093km 1000m = 1 km

km

m

mi

kmmi 2820,31

Chapter 1 45




1 mile = 1.6093km 1000m = 1 km

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

22

2820,31km

m

mi

kmmi

Chapter 1 46



22

2

1

1000

1

6093.1820,31

km

m

mi

kmmi

Chapter 1 47



2

26

2

22

1

101

1

5898.2820,31

km

m

mi

kmmi

Chapter 1 48



2102

26

2

22 102407.8

1

101

1

5898.2820,31 m

km

m

mi

kmmi

Chapter 1 49


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

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

Chapter 1 50



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

Chapter 1 51



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

Chapter 1 52



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

Chapter 1 53

End of Chapter ProblemsEnd of Chapter Problems

4, 10, 14, 20, 26, 34, 42, 60

chapter 11 matter and measurement chapter 1. 2 the study of chemistry what is chemistry? chemistry...

Documents

classification of matter

properties of matter

density slide

measurement chapter

behavior of matter

physical properties

element cobalt slide

homogeneous substance