Prentice Hall © 2003Chapter 1 Chapter 1 Introduction: Matter & Measurement CHEMISTRY The Central Science 9th Edition David P. White

Prentice Hall © 2003 Chapter 1

Chapter 1Chapter 1Introduction: Matter & Introduction: Matter &

MeasurementMeasurement

CHEMISTRYCHEMISTRY The Central Science The Central Science

9th Edition9th Edition

David P. WhiteDavid P. White


Molecular Perspective of Chemistry

• Matter is anything that has mass and occupies space.

• Matter ultimately consists of atoms.• Atoms are nature’s building blocks.• Compounds may consist of the same type of

atoms or different types of atoms.

The Study of ChemistryThe Study of Chemistry


Why Study Chemistry?

• Chemistry is the study of the composition of matter and the changes that matter undergoes.• Five general types of chemistry: organic, inorganic, biochemistry, analytical, and physical.• CHEMISTRY STUDIES EVERYTHING!!

The Study of ChemistryThe Study of Chemistry


Three States of Matter:

• Gas, liquid, or solid.• Gases (vapors) have an indefinite shape, indefinite volume,

and can be compressed.• Liquids have an indefinite shape, but a definite volume.

Liquids are not compressible.• Solids are rigid, having a definite shape and volume. They

are not compressible.

Classification of MatterClassification of Matter


Substances, Elements and Compounds:• Substance: matter having distinct properties and the same

composition from sample to sample. Substances can be classified as elements or compounds.

• Element: simplest form of matter that has a unique set of properties. Cannot be decomposed. Each element contains a unique kind of atom. Ex.: H or C

• Compound: a substance containing two or more elements chemically combined in a fixed proportion and structure. Ex.: NaCl



Law of Constant Law of Constant Composition (c. 1800)Composition (c. 1800)

• Also called the Law of Definite Proportions.• A given compound always contains exactly

the same proportion of elements by mass regardless of the source of the compound.

• Example: A molecule of pure water (H2O) always is made up of two hydrogen atoms and one oxygen atom.


Pure Substances and Mixtures• Mixtures: combinations of different substances.

Composition of mixtures can vary from sample to sample.• If matter is not uniform throughout, then it is a

heterogeneous mixture. Ex.: vegetable soup• If matter is uniform throughout, it is homogeneous.

Ex.: air. • Homogeneous mixtures are called solutions. Solutions

can be gaseous, liquid, or solid!



Mixtures, Substances, Mixtures, Substances, CompoundsCompounds

• 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.• Each element is given a unique chemical symbol (one or

two letters).• Elements are building blocks of matter.• The main elements in the human body: CHNOPS!



Physical properties can be measured without changing the identity and composition of a substance. Ex.: color,

density, m.p. • Intensive properties do not depend on the amount of

substance. Ex.: m.p., b.p., density• Extensive properties depend on the amount of substance

present. Ex.: mass, volumeChemical properties describe how a substance reacts to

form other substances. Ex: flammability

Properties of MatterProperties of Matter

Types of propertiesTypes of properties


Physical and Chemical Changes

• When a substance undergoes a physical change, its physical appearance changes, not its composition!

Ex. Changes of state (ice melting)

• When a substance changes its composition, it undergoes a chemical change. Chemical changes = chemical reactions

Ex.: solid iron + gaseous oxygen form solid iron oxide

Changes of MatterChanges of Matter


Physical and Chemical Changes



Four Clues to Chemical Changes• Transfer of energy• Change in color• Production of gas• Formation of a precipitate: solid that forms or settles

out from a liquid mixture.• But…even with a clue, you cannot be sure of a

chemical change! You need to test the composition of the sample before and after to be sure!

Chemical ChangesChemical Changes


Separation of Mixtures

• Mixtures can be separated if their physical properties are different.

• Separation is based on differences in physical properties.

• Separation by filtration, distillation, chromatography, etc.




• 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• Chromatography separates 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.



SI Units (Système International d’Unités)

• There are two types of units:

– fundamental (or base) units;

– derived units.

Units of MeasurementUnits of Measurement




Prefixes in the Metric System



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



Temperature

There are three temperature scales:• Kelvin Scale (Absolute Scale)

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



Temperature

• Celsius Scale– Water freezes at 0 oC and boils at 100 oC.– To convert: K = oC + 273.15.

• Fahrenheit Scale _Water freezes at 32 oF and boils at 212 oF.– To convert:


Temperature



Density

• Used to identify substances.

• Units: g/cm3 or g/mL

• Density is an intensive property.

• Density is temperature dependent. Why?

• Water is different!



Exact and Inexact Exact and Inexact NumbersNumbers

• Two kinds of numbers are used in science.

• Exact numbers: values are known exactly; values are infinitely precise. Ex: definitions (12 inches = 1 foot), the number 1 in conversion factors, counting numbers

• Inexact numbers: values have uncertainty. Ex: all measurements (12.34 cm)


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

reported for the measurement.

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 Accuracy



Significant Figures

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

• Significant figures: all the figures known with certainty plus one extra figure.



Significant Figures Rules• 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 point are

significant.• Zeros at the end of a number with no decimal point are

ambiguous (e.g. 10,300 g). Use scientific notation or a decimal to indicate number of significant figures.



Significant Figures

• Multiplication & Division: report answer to the smallest number of significant figures in the calculation.

• Addition & Subtraction: report answer to the smallest number of decimal places in the calculation.


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Prentice Hall © 2003Chapter 1 Chapter 1 Introduction: Matter & Measurement CHEMISTRY The Central Science 9th Edition David P. White