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Chemistry-140 Lecture 3

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Chemistry-140 Lecture 3 Chapter 1: Matter & Measurement Chapter Highlights physical properties & states of matter basic kinetic molecular theory – PowerPoint PPT presentation

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Title: Chemistry-140 Lecture 3


1
Chemistry-140 Lecture 3
  • Chapter 1 Matter Measurement
  • Chapter Highlights
  • physical properties states of matter
  • basic kinetic molecular theory
  • identify elements by name and symbol
  • atom, element, molecule compound
  • physical versus chemical changes
  • homogeneous heterogeneous mixtures
  • metric system
  • precision accuracy and significant figures

2
Chemistry-140 Lecture 3
  • Three States of Matter
  • In a chemical context, matter, can exist in one
    of three distinct forms, phases or states of
    matter.
  • Solids are rigid, have fixed volume and are
    noncompressible.
  • Liquids change shape but have a reasonably fixed
    volume and are essentially noncompressible.
  • Gases (vapours) readily change shape and volume
    are highly compressible.

3
Chemistry-140 Lecture 3
  • Physical Chemical Properties
  • Properties of substances can be divided into
    physical and chemical.
  • Physical properties are those that describe a
    substance as it is. For example colour,
    density, hardness, melting point etc.
  • Chemical properties are those that describe how a
    substance can be transformed. For example
    flammability, solubility, corrosiveness, etc.

4
Chemistry-140 Lecture 3
  • Basics of the Kinetic Molecular Theory
  • Used to interpret the behaviour of solids,
    liquids gases
  • Definition all matter consists of extremely
    small particles which are in constant motion.
  • the higher the temperature the faster the
    particles move
  • the kinetic energy of motion competes with the
    forces of attraction between particles
  • this allows for melting of a solid or
    vapourization of a liquid as the temperature
    increases

5
Chemistry-140 Lecture 3
  • Substances
  • Substances have fixed compositions and distinct
    properties
  • Pure substances are substances that are free from
    impurities.
  • Impurities are trace quantities of substances in
    a sample other than the principal component.
  • Most naturally occurring samples of matter
    consist of two or more substances.

6
Chemistry-140 Lecture 3
  • Mixtures
  • Mixtures are combinations of more than one
    substance
  • Homogeneous mixtures all species are in the same
    phase and uniformly distributed throughout the
    sample.
  • Heterogeneous mixtures all species are not in
    the same phase or they are not uniformly
    distributed

7
Chemistry-140 Lecture 3
  • Elements
  • Elements are substances that cannot be
    decomposed into simpler substances by chemical
    means.
  • Only a few of the elements are abundant on Earth.
  • Two elements comprise the majority of the Earth's
    crust
  • The human body concentrates three elements

8
Chemistry-140 Lecture 3
  • Compounds
  • Compounds consist of more than one element united
    chemically in definite mass proportions.
  • The law of constant composition or law of
    definite proportions states that, for a given
    chemical compound, the mass proportions of the
    elements are always the same.
  • butane has chemical formula C4H10

9
Chemistry-140 Lecture 3
  • Measurement
  • The metric system of units is attractive since
    all related units differ by powers of ten
  • For example, there are 100 centimeters in a
    meter.
  • SI (Système International d'Unités) is the
    recognized set of fundamental metric units for
    scientific measurement.
  • mass Kilogram (kg)
  • length Metre (m)
  • time Second (s)
  • electrical current Ampere (A)
  • temperature Kelvin (K)
  • light intensity Candela (cd)
  • amount of substance Mole (mol)

10
Chemistry-140 Lecture 3
  • Metric System
  • Metric uses prefixes to indicate subunits.
  • For example
  • 1 centimeter (cm) 1/100 meter 1 x 10-2 m
  • 1 millimeter (mm) 1/1,000 meter 1 x 10-3 m
  • 1 nanometer (nm) 1/1,000,000,000 meter 1 x
    10-9 m
  • Others quantities are derived from these.
  • Volume of a box is given by the product of the
    lengths.
  • One Liter is the volume of a box that is 10
    centimeters on each side 1 L 1000 cm3 or 1 mL
    1 cm3.
  • Density is the ratio of mass to volume of a
    sample (g/mL). For a pure substance, density is
    constant.

11
Chemistry-140 Lecture 3
  • Precision Accuracy
  • Precision is a measure of random error
  • (the less the random error, the higher the
    precision)
  • Precision refers to the reproducibility of a
    measurement.
  • Accuracy is a measure of systematic error
  • (the less the systematic error, the higher the
    accuracy)
  • Accuracy refers to the agreement of a measurement
    with the "correct" value.

12
Chemistry-140 Lecture 3
  • Precision Accuracy in Numbers
  • Exact numbers are those that are infinitely
    precise.
  • Most exact numbers are integers, such as.
  • 12 eggs 1 dozen, and are not measurements
  • (1 mole 6.022 x 1023).
  • Inexact numbers result from measurements
  • and are subject to error.
  • Precision and accuracy represent the two types of
    error.

13
Chemistry-140 Lecture 3
  • Significant Figures
  • Significant figures reflect precision.
  • All of the known digits are reported plus a last,
    estimated, digit.
  • For example
  • a measurement with a metric ruler, graduated in
    mm, might be reported as "27.6 mm" showing what
    the scale actually has markings for (27 mm) and
    an estimate between-the-lines" (.6 mm).

14
Chemistry-140 Lecture 3
  • Application of Significant Figures

Question What is the difference between 4.0 g
and 4.00 g? Answer 4.0 has two significant
figures 4.0 0.1 4.00 has three significant
figures 4.00 0.01 Question A balance
has precision of 0.001 g. How many significant
figures should be reported for a sample that
weighs about 25 g? Answer Five!!!
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