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Chapter 1 Introduction: Matter

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Title: Chapter 1 Introduction: Matter


1
Chapter 1 Introduction Matter Measurement
CHEMISTRY The Central Science 10th Edition
2
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.

3
Chemistry Catastrophe Prevention?
The space shuttle Columbia disintegrated in 2003
upon reentry into the Earths atmosphere due to a
damaged thermal protection system.
4
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.

5
  • Molecular Perspective of Chemistry

6
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.

7
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.

8
  • Pure Substances and Mixtures

9
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.

10
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.

11
Classification of Matter
  • Elements

12
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).

13
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.

14
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.

15
Classification of Matter
  • Mixtures
  • Heterogeneous mixtures are not uniform
    throughout.
  • Homogeneous mixtures are uniform throughout.
  • Homogeneous mixtures are called solutions.

16
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17
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.

18
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.

19
Properties of Matter
Physical and Chemical Changes
20
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.

21
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.

22
Separation of Mixtures
23
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.

24
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.

25
Units of Measurement
Base SI Units
26
Units of Measurement
SI Units
Selected Prefixes used in SI System
27
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?

28
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.

29
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

30
Class Practice Example
  • Make the following temperature conversions (a)
    68 oF to oC (b) -36.7 oC to oF

31
Units of Measurement
Temperature
32
Units of Measurement
  • Derived Units
  • Derived units are obtained from the 7 base SI
    units.
  • Example

33
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.

34
Units of Measurement
Volume
35
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).

36
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.

37
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.

38
Precision and Accuracy
39
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).

40
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).

41
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

42
Dimensional Analysis
  • Using Two or More Conversion Factors
  • Example to convert length in meters to length in
    inches

43
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.

44
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?
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