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Transcript of Prentice Hall © 2003Chapter 1 Chapter 1 Introduction: Matter & Measurement CHEMISTRY The Central...
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
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
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
Classification of MatterClassification of Matter
Prentice Hall © 2003 Chapter 1
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.
Prentice Hall © 2003 Chapter 1
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!
Classification of MatterClassification of Matter
Prentice Hall © 2003 Chapter 1
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.
Prentice Hall © 2003 Chapter 1
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!
Classification of MatterClassification of Matter
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
Physical and Chemical Changes
Properties of MatterProperties of Matter
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
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.
Properties of MatterProperties of Matter
Prentice Hall © 2003 Chapter 1
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.
Properties of MatterProperties of Matter
Separation of Mixtures
Prentice Hall © 2003 Chapter 1
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.
Properties of MatterProperties of Matter
Prentice Hall © 2003 Chapter 1
SI Units (Système International d’Unités)
• There are two types of units:
– fundamental (or base) units;
– derived units.
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
Prefixes in the Metric System
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
Derived Units• Derived units are obtained from base SI units.• Example:
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
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.
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
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:
Units of MeasurementUnits of Measurement
Temperature
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
Density
• Used to identify substances.
• Units: g/cm3 or g/mL
• Density is an intensive property.
• Density is temperature dependent. Why?
• Water is different!
Units of MeasurementUnits of Measurement
Prentice Hall © 2003 Chapter 1
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)
Prentice Hall © 2003 Chapter 1
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
Prentice Hall © 2003 Chapter 1
Precision and Accuracy
Uncertainty in MeasurementUncertainty in Measurement
Prentice Hall © 2003 Chapter 1
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.
Uncertainty in MeasurementUncertainty in Measurement
Prentice Hall © 2003 Chapter 1
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.
Uncertainty in MeasurementUncertainty in Measurement
Prentice Hall © 2003 Chapter 1
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.
Uncertainty in MeasurementUncertainty in Measurement