The Structure of the Atom

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The Structure of the Atom

• CHAPTER 4

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Early Theories of Matter
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460-370 BC Democritus

• Greek Philosopher
• Named atom
• smallest unit of matter
• means indivisible

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1807 John Dalton Atomic Theory

• Revived and revised Democritus ideas and began
  developing the modern atomic theory

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Daltons Atomic Theory

• all elements are composed of tiny indivisible
  particles called atoms
• atoms of same element alike. Each element is
  different from atoms of other elements
• atoms of different elements combine in simple
  whole number ratios to form compounds
• chemical reactions occur when atoms are
  rearranged

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• Picture shows
• Conservation of Mass
• Element combing in simple whole number ratios

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Subatomic Particles the Nuclear Atom

• Chapter 4

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A. Discovering the Electron
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1879 William Crookes

• investigated electrical discharge in gases
• cathode ray tube
• Cathode rays are streams of negatively charged
  particles.
• The particles are found in all forms of matter

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1897 J.J. Thomson

• determined nature of cathode ray
• determined charge to mass ratio of electron
• Found that atoms were divisible - Dalton
  Democritus were wrong

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1901 J.J. Thomson

• positive beam experiments
• plum pudding model of atom or chocolate-chip
  cookie dough model of the atom

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1909 Robert Milliken

• determined charge of electron
• oil drop experiment
• with Thomsons charge to mass ratio able to
  determine the mass of e-
• Mass of electron 9.1x10-28 grams

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1911 Ernest Rutherford

• discovered nucleus
• gold foil experiment disproved plum pudding model
• small dense central part of atom nucleus
• () charge

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1920 Rutherford

• Refined concept of nucleus
• Concluded that nucleus contained positively
  charged particles called protons

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1932 James Chadwick

• identified neutron
• same mass as proton
• no charge

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How Atoms Differ
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NUCLEUS

• Protons with () charge
• Neutrons with no charge.
• Protons neutrons have about the same mass.
• () charge is responsible for most of mass of
  atom (dense central part).

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ELECTRONS

• move around nucleus
• responsible for most of volume of atom
• (-) charge
• negligible mass

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ATOMIC NUMBER MASS NUMBER
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ATOMIC NUMBER

• of protons in nucleus
• it identifies the element
• elements in Periodic Table are listed in
  increasing order of atomic
• if atom is neutral the of protons equals of
  electrons

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MASS NUMBER

• sum of protons neutrons in nucleus
• written as part of name must be given to you
• Neon-20 mass 20 p n 20
• atomic 10 p 10
• n 10

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• mass
• p n
• Symbol
• atomic
• p
• ? neutrons

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• Oxygen-17
• mass 17
• atomic 8
• p n 17
• p _8_
• 9 n

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• To calculate electrons for an ion you must look
  at the charge written in the upper right corner
• To determine the number of electrons
• If the charge is positive then subtract that
  number from the number of protons.
• If the charge is negative then add that number to
  the number of protons

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ISOTOPES

• Atoms of the same element are not all identical -
  may differ in of neutrons
• Isotopes - atoms of the same elements (same of
  protons), but different mass (different
  neutrons) and therefore different masses

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• 12 13 14
• C C C
• 6 6 6
• ? neutrons
• 6 7 8

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ATOMIC MASS
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ATOM

• smallest unit of an element that can exist alone
  and still have the properties of that element

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Atomic Mass

• Average Atomic Mass - weighted average of the
  masses of the naturally occuring isotopes
• relative mass based on carbon-12 as the standard
• Carbon-12 is defined as having a mass of exactly
  12 amu
• atomic mass unit (amu) - 1/12 of the mass of a
  carbon-12 atom

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Weighted Average Example

• 50 test 70
• 30 Lab 80
• 20 Daily 90
• 50(70) 30(80) 20(90)
• .5(70) .3(80) .2(90)
• 35 24 18
• 77

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Isotopes of Hydrogen

• H-1 protium 1.0078 99.985
• H-2 deuterium 2.0140 0.015
• H-3 tritium 3.0160 -------

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Calculating Average Atomic Mass

• Multiply the percent (as a decimal ) by the mass
  and then add each together.
• 99.985 (1.0078 amu) 0.015 (2.0140 amu)
• .99985 (1.0078 amu) .00015 (2.0140 amu)
• 1.0076488 amu 0.0003021 amu
• 1.00795 amu

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Example

• chlorine - 35 75.8
• chlorine - 37 24.2
• Will the average atomic mass be closer to 35 or
  37?
• (35 because higher )
• 75.8(35) 24.2(37)
• 35.5 amu

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Example

• Calculate the average atomic mass for boron. Be
  as accurate as possible. Use values from table
  on p. 281. Check your answer by looking at the
  periodic table.