1 brown_et al chapter_1

Dr. Curry’s Life Mississippi

“Place of birth” Military Tour

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University of Alabama “Analytical Chemistry, Ph.D.”

University of West Alabama “Bachelor of Science”

Chapter 1Introduction: Matter &

Measurement

CHEMISTRY The Central Science

10th Edition

• Chemistry is the study of the properties of materials and the changes that materials undergo.

• Chemistry is central to our understanding of other sciences.

• Chemistry is also encountered in everyday life.

Why Study Chemistry

Chemistry: Catastrophe Prevention?

The space shuttle Columbia disintegrated in 2003 upon reentry into the Earth’s atmosphere due to a damaged thermal protection system.

The Molecular Perspective of Chemistry• Matter is the physical material of the universe.• Matter is made up of relatively few elements.• On the microscopic level, matter consists of atoms and

molecules.• Atoms combine to form molecules.• As we see, molecules may consist of the same type of

atoms or different types of atoms.

The Study of Chemistry

Molecular Perspective of Chemistry

States of Matter• Matter can be a gas, a liquid, or a solid.• These are the three states of matter.• Gases take the shape and volume of their container.• Gases can be compressed to form liquids.• Liquids take the shape of their container, but they do

have their own volume.• Solids are rigid and have a definite shape and volume.

Classification of Matter

• Elements consist of a unique type of atom.• Molecules can consist of more than one type of element.

– Molecules that have only one type of atom (an element).– Molecules that have more than one type of atom (a compound).

• If more than one atom, element, or compound are found together, then the substance is a mixture.

Pure Substances and Mixtures


• Pure Substances and Mixtures

Pure Substances and Mixtures• If matter is not uniform throughout, then it is a

heterogeneous mixture.• If matter is uniform throughout, it is homogeneous.• If homogeneous matter can be separated by physical

means, then the matter is a mixture.• If homogeneous matter cannot be separated by physical

means, then the matter is a pure substance.• If a pure substance can be decomposed into something

else, then the substance is a compound.


Elements• If a pure substance cannot be decomposed into

something else, then the substance is an element.• There are 114 elements known.• Each element is given a unique chemical symbol (one or

two letters).• Elements are building blocks of matter.• The earth’s crust consists of 5 main elements.• The human body consists mostly of 3 main elements.


Elements


Elements• Chemical symbols with one letter have that letter

capitalized (e.g., H, B, C, N, etc.)• Chemical symbols with two letters have only the first

letter capitalized (e.g., He, Be).


Compounds• Most elements interact to form compounds.

• Example, H2O• 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.


Compounds• If water is decomposed, then there will always be twice

as much hydrogen gas formed as oxygen gas.• Pure substances that cannot be decomposed are elements.


Mixtures• Heterogeneous mixtures are not uniform throughout.• Homogeneous mixtures are uniform throughout.• Homogeneous mixtures are called solutions.


Physical vs. Chemical Properties• Physical properties can be measure without changing the

basic identity of the substance (e.g., color, density, odor, melting point)

• Chemical properties describe how substances react or change to form different substances (e.g., hydrogen burns in oxygen)

• Intensive physical properties do not depend on how much of the substance is present.– Examples: density, temperature, and melting point.

• Extensive physical properties depend on the amount of substance present.– Examples: mass, volume, pressure.

Properties of MatterProperties of Matter

Physical and Chemical Changes• When a substance undergoes a physical change, its

physical appearance changes. – Ice melts: a solid is converted into a liquid.

• Physical changes do not result in a change of composition.

• When a substance changes its composition, it undergoes a chemical change:– When pure hydrogen and pure oxygen react completely, they

form pure water. In the flask containing water, there is no oxygen or hydrogen left over.


Physical and Chemical Changes


Separation of Mixtures• Mixtures can be separated if their physical properties are

different.• Solids can be separated from liquids by means of

filtration.• The solid is collected in filter paper, and the solution,

called the filtrate, passes through the filter paper and is collected in a flask.


Separation of Mixtures• Homogeneous liquid mixtures can be separated by

distillation.• Distillation requires the different liquids to have different

boiling points.• In essence, each component of the mixture is boiled and

collected.• The lowest boiling fraction is collected first.


Separation of Mixtures

Separation of Mixtures• Chromatography can be used to separate mixtures that

have different abilities to adhere to solid surfaces.• The greater the affinity the component has for the surface

(paper) the slower it moves.• The greater affinity the component has for the liquid, the

faster it moves.• Chromatography can be used to separate the different

colors of inks in a pen.

Units of MeasurementUnits of Measurement

SI Units• There are two types of units:

– fundamental (or base) units;– derived units.

• There are 7 base units in the SI system.


Units of MeasurementUnits of MeasurementBase SI Units

SI Units

Selected Prefixes used in SI System


Class Practice ExamplesClass Practice Examples

• What is the name given to the unit that equals (a) 10-9 grams; (b) 10-6 second; (c) 10-3 meter

• What fraction of a meter is a nanometer?

SI Units• Note the SI unit for length is the meter (m) whereas the SI unit for

mass is the kilogram (kg).– 1 kg weighs 2.2046 lb.

TemperatureThere are three temperature scales:• Kelvin Scale

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


Temperature• Celsius Scale

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

• Fahrenheit Scale– Not generally used in science.– Water freezes at 32 oF and boils at 212 oF.– To convert:

32-F95C 32C

59F


• Make the following temperature conversions: (a) 68 oF to oC; (b) -36.7 oC to oF

32-F95C 32C

59F

Class Practice ExampleClass Practice Example

Temperature


Derived Units• Derived units are obtained from the 7 base SI units.• Example:


m/ssecondsmeters

timeof unitsdistance of units velocityof Units

Volume

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

m3.• We usually use 1 mL = 1

cm3.• Other volume units:

– 1 L = 1 dm3 = 1000 cm3 = 1000 mL.


Volume


Density• Used to characterize substances.• Defined as mass divided by volume:

• Units: g/cm3.• Originally based on mass (the density was defined as the

mass of 1.00 g of pure water).


volumemassDensity

• Answer the following problems:• (a) Calculate the density of mercury if 1.0 x 102 g

occupies a volume of 7.36 cm3.

• (b) Using the density for mercury, calculate the mass of 65.0 cm3 of mercury.

Class Practice ExamplesClass Practice Examples

• 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.Precision and Accuracy

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

• Measurements that are close to each other are precise.

Uncertainty in MeasurementUncertainty in Measurement

Precision and AccuracyPrecision and Accuracy

Significant Figures• 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• Non-zero numbers are always significant.• Zeros between non-zero numbers are always significant.• Zeros before the first non-zero digit are not significant.

(Example: 0.0003 has one significant figure.)• 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 (e.g. 10,300 g).


• Method of calculation utilizing a knowledge of units.• Given units can be multiplied or divided to give the

desired units. • Conversion factors are used to manipulate units:• Desired unit = given unit (conversion factor)• The conversion factors are simple ratios:

unitgiven unit desiredfactor Conversion

Dimensional AnalysisDimensional Analysis

Using Two or More Conversion Factors• Example to convert length in meters to length in inches:

cm 2.54

in 1m

cm 100m ofnumber in ofNumber

in cm conversioncm m conversionm ofnumber in ofNumber


• A person’s height is measured to be 67.50 in. What is this height in centimeters?

• Perform the following conversions: (a) 2 days to s; (b) 20 Kg to g.

Class Practice ProblemClass Practice Problem

Using Two or More Conversion Factors• 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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