The Atom

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Chapter 2

• The Atom

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What is an atom?

• It is the smallest particle of an element that
  retains the properties of that element.
• It is composed of 3 subatomic particles
• 1. Protons
• 2. Neutrons
• 3. Electrons

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Regions of the AtomThe nucleus

• Nucleus very small dense region near the
  center of the atom. It contains protons AND
  neutrons
• Protons positively charged (1) subatomic
  particle with a mass of 1 amu (atomic mass unit)
  or
• 1.673X 10-27 kg Symbol p

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Regions of the AtomThe nucleus contd

• Neutrons neutrally charged (0) subatomic
  particle with a mass of 1 amu. Symbol n0
• Note The mass of both the proton and neutron are
  almost exactly the same. These 2 particles give
  the atom nearly all its mass!

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Discovery of the Nucleus

• Ernest Rutherford
• Used a lead-shielded box and polonium which emits
  positively charged particles.
• Shot a beam of these particles at a VERY thin
  sheet of gold foil.
• Most of the particles passed straight through the
  foil and others were deflected or bounced back.
• The results were totally unexpected.
• For more info see your text p. 64

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Regions of the AtomThe Electron Cloud

• Electron Cloud region surrounding the nucleus
  occupied by negatively charged subatomic
  particles called electrons. This region is very
  LARGE, but is mostly EMPTY SPACE!
• Electrons negatively charged subatomic particle
  (-1) with a mass of approximately 0 amu or 9.109
  X 10-31kg. Its mass is so small that it
  practically makes NO contribution to the total
  mass of the atom.
• Symbol e-

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Discovery of the Electron

• J.J. Thompson and the cathode-ray tube
• Used a vacuum tube, electrodes, and an electric
  source.
• Concluded that cathode rays are made up of
  invisible, negatively charged particles called
  electrons.
• See page 61 in your text for more info.

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Summary of Subatomic Particles
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The Atom as a Whole

• Atoms are electrically neutral, they do not have
  a charge, and therefore, the total number of
  electrons MUST equal the number of protons.
• p e-
• An atoms total mass is primarily composed of
  only the protons neutron (the nucleus). The
  electrons mass is so small that it does not
  contribute.

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Counting Parts of the Atom

• Atomic Number (Z) the number of protons in the
  nucleus of each atom of that element.
• Atomic (Z) of protons
• The atomic number identifies an element. If the
  number of protons changes then the identity of
  the element changes with it!

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Counting Parts of the Atom

• Mass number (A) the total of protons and
  neutrons in the nucleus of an atom
• Mass (A) protons neutrons

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Isotopes

• Isotopes atoms of the same element with
  different numbers of neutrons. This causes them
  to have different mass numbers.

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Table The Isotopes of Hydrogen
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ExampleFill in the following table
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Average Atomic Mass

• Average atomic mass the weighted average of the
  atomic masses of the naturally occurring isotopes
  of an element.
• The abundance of each isotope MUST be considered
  when finding the average.

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Calculating the average mass

• Average mass
• (m11)(relative abundance) (m12)(relative
  abundance) .
• m mass of each isotope
• Relative abundance is the percentage in decimal
  form.
• (Move the decimal two places to the left)

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Example 1

• Naturally occurring copper is composed of 77.2
  copper-63, which has a mass of 63 amu and 22.8
  copper-65, which has a mass of 65 amu. What is
  the average atomic mass of copper?
• Average (63amu) (0.772)
• (65 amu) (0.228)
• Average 63.456 amu

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Example 2

• Oxygen has 3 isotopes. The masses and relative
  abundances are listed below. Calculate the
  average atomic mass of oxygen.
• Mass Abundance
• 15.994915 amu 99.762
• 16.999131 amu 0.038
• 17.999160 amu 0.200

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Atomic Mass Unit

• Equal to exactly 1/12 of the carbon-12 atom.
  (1.660 540 x 10-24g)

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Moles

• Mole the amount of a substance that contains as
  many particles as there are atoms in 12 g of
  carbon-12. This is a counting unit.

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Avogadros Number

• The number of particles in a mole has been
  determined to be 6.022 x 1023. (Named after
  Avogadro, a scientist whose ideas suggested a
  relationship between mass and the number of
  atoms.)

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Molar Mass

• The mass of 1 mole of a pure substance
• Equal to the average atomic mass in grams. (Use
  mass from the periodic table)

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New conversion factor(s)

• 6.022 x 1023 atoms 1 mole Periodic Table Mass
  (g)

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Practice Problems