Ch' 4 Atomic Structure

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Ch. 4 Atomic Structure
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Atoms

• The smallest particle of an element that retains
  its identity in a chemical reaction

Iron Atoms on Copper
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Democritus (460 B.C.-370 B.C)

• First to suggest existence of atoms
• Said atoms are indivisible and indestructible

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Daltons Atomic Theory 1766-1844

• Elements are made up of tiny particles called
  atoms
• All atoms of a given element are identical
• All atoms of a given element are different from
  those of a different element
• Atoms combine to form compounds. These compounds
  contain specific ratios.
• Atoms are indivisible by chemical processes- a
  solid indivisible mass

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Errors in Daltons Theory

• Atoms are now known to be divisible
• Protons
• Neutrons
• Electrons

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Electrons

• Negatively charged subatomic particles
• (-1)
• Discovered by J.J. Thompson

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J.J Thomson 1856-1940

• Discovered the electron in 1897 using a cathode
  ray tube
• Cathode ray repelled by the negative charge of
  electric field

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Protons

• Positively charged subatomic particles (1)
• Discovered in 1886 by Eugene Goldstein ? Canal
  Rays travel in opposite directions as cathode
  rays

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Neutrons

• Subatomic particles with no charge but mass
  nearly equal to that of a proton

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Ernest Rutherford 1871-1937

• Proposed the nuclear atom-electrons surround a
  nucleus
• 1911 directed a particles at gold foil- some were
  reflected or deflected, indicating a concentrated
  positive charge
• Discovered the proton-later experiments showed
  the presence of neutrons

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Nucleus

• Protons and Neutrons are located in the nucleus

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Atomic Number

• The number of protons in the nucleus of an atom
• Elements are different because they contain
  different numbers of protons

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Mass Number

• The total number of protons and neutrons in an
  atom

Mass Number
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Calculating Neutrons
6 Neutrons
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Isotopes

• Atoms that have the same number of protons but
  different numbers of neutrons
• Because they have different numbers of neutrons
  they also have different mass numbers

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Atomic Mass

• Is the weighted average mass of the atoms in a
  naturally occurring sample of the element
• Reflects both the mass and the relative abundance
  of the isotopes as they occur in nature

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Atomic Mass (AMU)

• To calculate the atomic mass of an element,
  multiply the mass of each isotope by its natural
  abundance, expressed as a decimal, and then add
  the products

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Example
98.89
1.11
Atomic Mass 12.00 amu X 0.9889
13.003 amu X .0111 12.011 amu
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Periodic Table
Group or Family
Period