Introduction: Matter and Measurement

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Introduction Matter and Measurement

• Chapter 1 BLB 11th

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1.1 The Study of Chemistry

• Chemistry is everywhere!
• Matter is everywhere!
• Thus, chemistry matters!
• Chemistry involves the study of matter its
  properties and behavior.
• Macroscopic observations are rooted in
  microscopic structure.

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Why study chemistry?

• Its required.
• It sounds interesting.
• Its unavoidable.
• It truly is the central science.
• Name an element
• Name a compound
• Name a mixture

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The Periodic Table of the Elements
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Molecules

• O2, H2O, CO2, C2H5OH, C2H6O2, C9H8O4
• Models
• Shown on p. 2

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Expectations

• Classify matter
• Properties of matter
• g ? mL (using density)
• Solve for any variable in a formula.
• Metric unit conversions
• Other conversions temperature, metric-English,
  etc.
• Identify and work with significant figures.

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1.2 Classification of Matter

• Matter anything which has mass and takes up
  space.
• States of matter (p. 5)
• Solid rigid, regular
• Liquid fluid, irregular
• Gas open, random
• Phases of matter

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States of Matter
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States of Matter
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Alternative diagram to that on p. 9
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Atoms are the building blocks of matter.
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The Periodic Table of the Elements
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Elements
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Group Activity

• Get into groups of four or five. No solo flyers!
• Introduce yourself.
• Work together.
• Discuss, argue, and intellectually engage.
• Choose a person to record and report the groups
  result.

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Problem 1

• Describe the contents of the containers.
• Devise a plan to determine which liquid is in
  each of the two containers.

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Description
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Strategy for identification
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1.3 Properties of Matter

• physical measured or observed without changing
  the identity of a substance, e.g. physical state,
  color, odor, density, boiling point
• chemical describes a substances reactivity,
  e.g. flammability, corrosiveness
• extensive depends on the amount of matter
  present, e.g. mass, volume
• intensive does not depend on the amount of
  matter present, e.g. density, color, temperature

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Physical Chemical Changes

• Physical change in appearance, not in
  composition, e.g. phase changes, separation of
  mixtures filtration, distillation,
  chromatography (p. 12)
• Chemical new substance is formed as the
  chemical identities change, e.g. any chemical
  reaction (pp. 10-11)
• Dissolve vs. react
• Explode vs. ignite

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1.4 Units of Measurement (SI Units)
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Volume a derived unit
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Metric Prefixes
Angstrom Å 10-10 m
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Temperature Scales
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Temperature Conversions

• F ? C
• C ? F
• C ? K

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Density

• Density mass per unit volume
• D m/V (g/cm3 or g/mL)
• Measured at a specific temperature
• Useful as a conversion factor (g ? mL)
• Most substances become more dense at lower
  temperatures.
• Specific gravity density of a substance divided
  by the density of a reference substance (usually
  water) no units

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Difference in density values is the reason some
things float and others sink.
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Density of Water
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1.5 Uncertainty in Measurement

• Exact numbers have a defined value, e.g.
  12-dozen, 2.54 cm/in 1000 g 1 kg count of
  objects
• All measurements have some degree of uncertainty
  inexact
• Types of error systematic random
• The last digit of a measured quantity is
  uncertain.
• The more significant figures, the greater the
  certainty.
• precision agreement among data
• accuracy agreement of data with true value

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Different measuring devices have different uses
and different degrees of accuracy and precision.
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Significant Figures
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Significant Figures in Calculations

• A calculated result can be no more certain than
  the data measured.
• Mathematical operations (p. 23)
• Averaging least number of decimal places
• and - least number of decimal places
• x and least number of sig. figs.
• Round off at the end at the end of a multi-step
  problem.

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1.6 Dimensional Analysis

• Problem-solving strategies (pp. 26 28)
• Estimate and then calculate your answer. Do the
  two agree?
• Get your units correct and your answer should be
  correct.
• Report to correct number of sig. figs.
• Practice, practice, practice!