Chapter 1: Matter and Measurements

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Chapter 1 Matter and Measurements

• Outline
• 1.1 Types of Matter
• 1.2 Measurements
• 1.3 Properties of Substances

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Introduction

• We will
• I. Consider the different types of matter pure
  substances vs mixtures, elements vs compounds.
• II. Look at the kinds of measurements on which
  chemistry is founded, the uncertainties
  associated with these measurements and a method
  used to convert measured quantities from one unit
  to another.
• III. Focus on certain physical properties
  including color, density, and water solubility
  that can be used to identify substances.

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Types of Matter

• Matter is anything that has mass and takes up
  space.
• Matter can exist as one of three phases solid,
  liquid, or gas
• a. solid has a rigid shape and fixed volume
• b. liquid fixed volume but no rigid in shape
• c. gas has neither a fixed volume or rigid shape

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Matter (Contd)

• Matter is classified into two categories
• I. Pure substances-has a fixed composition and a
  unique set of properties.
• II. Mixtures-composed of two or more substances.
• Pure substances are either elements or compounds,
  where as mixture can be either homogeneous or
  heterogeneous.

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Matter (Contd)

• An element is a type of matter that cannot be
  broken down into two or more pure substances.
• There are 112 elements, 91 of which occur
  naturally.
• Elements are identified by their elemental
  symbols. These consist of one or two letters,
  usually derived from the name of the element.
• e.g. Aluminum-Al, Carbon-C, Copper-Cu, Mercury-Hg

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Elements (Contd)

• Table 1.1 Common Elements and their percentage
  abundances on earth
• A compound is a pure substance made up of more
  than 1 element.
• e.g. water (H2O) methane (CH4) acetylene (C2H2).

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Compounds (Contd)

• Compounds have fixed compositions. In
  otherwords, a given compound always contains the
  same elements in the same percentages by mass.
• Mixtures, however, can vary in composition.

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Compounds (Contd)

• Properties of compounds differ greatly from those
  of the elements they contain.
• e.g. Sodium Chloride (Table Salt)
• NaCl ? Na Cl-
• Compounds can be resolved into its elements by
  electrolysis (passing a current through a
  compound, usually a liquid).
• e.g. H2O ? H2(g) O2(g)

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Mixtures

• A mixture contains two or more susbstances
  combined in such a way that each substance
  retains its chemical identity.
• I. Homogeneous uniform mixtures in which the
  composition is the same throughout also called a
  solution.
• a. Solvent substance present in greatest
  amount, usually a liquid.

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Mixtures (Contd)

• b. solute substance present in least amount
  can be gas, liquid, or solid.
• II. Heterogeneous nonuniform mixture is one in
  which the composition varies throughout.
• Mixtures can be separated by techniques such as
  filtration (heterogeneous) or distillation
  (homogeneous).

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Measurements

• Scientific measurements are expressed in the
  metric system, which is a base-10 system.
• Table 1.2 Metric Prefixes
• 106 mega M 10-3 milli m
• 103 kilo k 10-6 micro µ
• 10-1 deci d 10-9
  nano n
• 10-2 centi c 10-12 pico p

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Instruments and Units

• The meter is the standard metric unit of length.
• 1 cm 10-2 m 1 mm 10-3m 1 km 103m
• Volume is most commonly expressed in one of three
  units
• 1 cm3 (10-2 m)3 10-6 m3
• 1 L 10-3 m3 103 cm3
• 1 mL 10-3 L 10-6 m3

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Units (Contd)

• The gram is the standard metric unit for mass.
  It can also be expressed in kilograms or
  milligrams.
• e.g. 1 g 10-3g 1 mg 10-3g
• Mass is a measure of the amount of matter in an
  object weight is a force, it is a measure of the
  gravitational force acting on an object.

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Units (Contd)

• Temperature is the factor that determines the
  direction of heat flow. Temperature is measure
  in degrees Celsius in the laboratory.
• For many purposes in chemistry, temperature is
  expressed in the units of kelvin (K)
• TK tºC 273.15

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Uncertainties in Measurements Significant Figures

• The uncertainty in measurement comes from the
  instrument used in the measuring.
• e.g. 8 ? 1 mL (large graduated cylinder)
• 8.0 ? 0.1 mL (small graduated cylinder)
• 8.00 ? 0.01 mL (buret)
• It is understood that there is an uncertainty of
  at least 1 unit in the last digit. It is often
  described in terms of significant figures
  (meaningful digits obtained in a measurement).

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Significant Figures

• E.g. 8.00 3 sig. figs.
• 8.0 2 sig. figs.
• 8 1 sig. fig.
• Suppose a piece of metal is reported to weigh
  500g. You do no know how many of these figures
  are significant.
• If the metal was weighed to the nearest gram,
• 500 ? 1g, than there are 3 significant figures.

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Significant Figures (Contd)

• If the metal had been weighed to the nearest 10g,
  500 ? 10g, then there are only 2 significant
  figures.
• If the technician had used scientific notation,
  the number of significant figures is easily
  identified.
• e.g. 5.00 x 102g 3 sig. figs.
• 5.0 x 102g 2 sig. figs.
• 5 x 102g 1 sig. fig.

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Significant Figures (Contd)

• Most measured quantities are used to calculate
  other values. The precision of these final
  values is limited by that of the measurements on
  which it is based.
• When measured quantities are multiplied or
  divided, the number of significant figures is
  that same as that in the quantity with the
  smallest number of significant figures.

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Example

• Average speed distance traveled/time elapsed
• Average speed 5.6 x 103 km/8.50 hours
• 658.8235294 km/hr
• 6.6 x 102 km/hr

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Significant Figures (Contd)

• When measured quantities are added or subtracted,
  the number of decimal places in the result is the
  same as that in the quantity with the greatest
  uncertainty and hence the smallest number of
  decimal places.
• E.g. Find the total mass of a solution made up of
  10.21g of instant coffee, 0.2g of sugar, and 256g
  of water.

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Solution to Example 2

• Instant Coffee 10.21g ? 0.01g 2 decimal
• Sugar 0.2 g ? 0.1g 1 decimal
• Water 256 g ? 1g 0 decimal
• Total Mass 266g

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Significant Figures (Contd)

• In calculations, there are certain number which
  are exact rather than estimates.
• e.g. tF 1.8tC 32
• If in a problem, the amount is spelled out (i.e.
  one kilogram vs 1.0 kilogram), the implication is
  that the number is exact and therefore there is
  no uncertainty.

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Conversion of Units

• Do some examples on the board.

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Properties of Substances

• The properties used to identify an uknown
  substance must be intensive (independent of
  amount).
• Extensive properties are those that are dependent
  on amount (mass, volume, etc).

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Examples of Intensive Properties

• Chemical Properties observations of the
  substance in a chemical reaction.
• Physical Properties observed without changing
  the chemical identity of the substance.
• a. melting point
• b. boiling point

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Density

• The density of a substance is the ratio of the
  mass to volume
• density mass/volume d m/v
• Do example 1.6.

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Solubility

• Solubility is stated as the amount of a substance
  in grams that dissolves in 100g of solvent at a
  given temperature.
• e.q. At 20C, 32g of potassium nitrate dissolves
  in 100g of water.

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Color Absorption Spectrum