Chapter 1 Introduction: Matter

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Chapter 1Introduction Matter Measurement
CHEMISTRY The Central Science 10th Edition
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Why Study Chemistry

• 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.

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Chemistry Catastrophe Prevention?
The space shuttle Columbia disintegrated in 2003
upon reentry into the Earths atmosphere due to a
damaged thermal protection system.
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The Study of Chemistry

• 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.

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• Molecular Perspective of Chemistry

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Classification of Matter

• 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.

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Pure Substances and Mixtures
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.

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• Pure Substances and Mixtures

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Classification of Matter

• 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.

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Classification of Matter

• 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 earths crust consists of 5 main elements.
• The human body consists mostly of 3 main
  elements.

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Classification of Matter

• Elements

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Classification of Matter

• 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).

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Classification of Matter

• 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.

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Classification of Matter

• 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.

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Classification of Matter

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

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Properties of Matter

• 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.

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Properties 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.

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Properties of Matter
Physical and Chemical Changes
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Properties of Matter

• 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.

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Properties of Matter

• 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.

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Separation of Mixtures
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Units of Measurement

• 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.

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Units 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.

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Units of Measurement
Base SI Units
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Units of Measurement
SI Units
Selected Prefixes used in SI System
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Class 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?

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Units of Measurement

• 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.
• Temperature
• There 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.

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Units of Measurement

• 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

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Class Practice Example

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

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Units of Measurement
Temperature
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Units of Measurement

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

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Units of Measurement
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.

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Units of Measurement
Volume
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Units of Measurement

• 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).

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Class Practice Examples

• 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.

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Uncertainty in Measurement

• 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.

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Precision and Accuracy
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Uncertainty in Measurement

• 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).

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Uncertainty in Measurement

• 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).

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Dimensional Analysis

• 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

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Dimensional Analysis

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

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Class Practice Problem

• A persons 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.

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Dimensional Analysis

• 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?