The Periodic Law

1
The Periodic Law
2
History of the Periodic Table

• Antoine Lavoisier
• (France, 1789)
• Earned reputation as
• father of chemistry
• Established a common
• naming system of compounds
• and elements.
• First to organize elements
• Grouped them into four
• categoriesGases, nonmetals, metals and earths
  (elements that could not be chemically separated
  at the time.)

3
History of the Periodic Table

• Dmitri Mendeleev
• (Russia, 1869)
• 1.Placed elements in groups in which
• they shared similar properties which resulted in
  order of
• increasing atomic mass with a few exceptions.
• 2. Ex If placed solely by atomic mass, iodine
  was
• not in group with chemically similar elements.
• 3. Left gaps for not-yet-discovered elements and
  predicted their properties gallium, germaniun
  scandium.

4
History of the Periodic Table

• Henry Mosely
• Moseley (1911) modified the table by organizing
  elements in order of increasing atomic numbers.
• Periodic Law The phsical and chemical properties
  of the elements are the periodic functions of
  their atomic numbers.
• Glenn Seaborg
• (UC Berkeley, 1944)
• Formed Actinide Series
• just like that of the
• lanthanides (58-71)

5
Basics of the Periodic Table

• periodic a repeating pattern
• table an organized collection of information

• period horizontal row on the P.T.
• Designates e- energy levels

group or family vertical column on the P.T.
Periodic Table an arrangement of elements in
order of atomic number elements with similar
properties appear at regular intervals (are in
the same group)
6
Electron Structures of Atoms
nucleus
1st energy level
1st Period Hydrogen (1) Helium (2)
2nd energy level
2nd Period Lithium (3) Neon (10)
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S block elements Group 1 2

• Chemically reactive metals, group 1 more reactive
  than group 2. Group Config ns 1-2
• Alkali metalssilvery appearance, soft enough to
  cut with a knife, not found in nature as free
  elements. H shares e-config but not properties.
• Alkaline-earth metalsharder, denser, stronger,
  and have a higher melting point than group 1.
  Too reactive to be found uncombined in nature. He
  shares e- config but not properties.

8
p block elements Group 13-18

• Includes all the three types of elements metals,
  non metals and metalloids.
• Group Config ns2 np 1-6
• Includes Halogens most reactive of the
  nonmetals. React vigorously with most metals to
  form salts.
• P block metals are generally harder denser than
  s block but softer less dense than d block
  metals.Found in nature solely as compounds except
  for bismuth.

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d block elements Group 3-12

• Transition Elements metals with typical
  properties good conductors, high luster.
• Less reactive than s block, many existing in
  nature as free elements.
• Electrons added to the d sublevel of the
  preceding energy level (n-1).
• Group configuration (n-1)d1-10ns 0-2
• Some deviations from orderly d sublevel filling
  occur in group 4-11(s electrons jumping to d
  sublevel)

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f-block elements

• F-block elements are wedged between groups 3 and
  4 in the sixth and seventh period, consisting of
  lanthanides and actinides
• Most elements are radioactive
• Trans Uranium elements are all synthetic
• Group Config ns 0-2 (n-1) d 0-1 (n-2)f 1-14

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• atomic radius

• Covalent Radius for Covalently Bonded Atoms half
  the distance between the nuclei of two covalently
  bonded atoms
• F-F bond length is 144 pm, so F covalent radius
  is 72 pm.
• H-F bond length is 109 pm, so H covalent radius
  is 37 nm.
• Atomic Radius for Elements like the Noble Gases
• Ar atomic radius is 131 pm
• Metallic Radius for Metals
• Al metallic radius is 143 pm.

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Trends in Atomic Size
6.3

• The atomic radius is one half of the distance
  between the nuclei of two atoms of the same
  element when the atoms are joined.

13
Trends in Atomic Size
6.3
14
6.3
15
Cationpositive ion, Anionnegative ion
Ionic Radii

• Forming a cation by losing electron(s) leads to a
  decrease in atomic radius, a smaller electron
  cloud.
• Forming an anion by electron(s) leads to an
  increase in atomic radius, less pull from the
  nucleus there is more repulsion between the
  greater number of electrons.

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Cations
6.3
17
Anions
6.3
18
Across a period atoms become smaller. Down a
group atoms become larger.
19
Ionization Energy

• Amount of energy required to remove an e from a
  neutral atom in its gaseous state.
• First Ionization Energy
• A(g) ? A(g) e-
• Second Ionization Energy
• A(g) ? A2(g) e-
• Third Ionization Energy
• A2(g) ? A3(g) e-

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Picture of IE trend
I.E. increases across a period and decreases down
a group.
21
Trends in Ionization Energy
6.3
22
Trends in Ionization Energy
6.3
23
Electron Affinity

• Amount of energy released when an e is added to a
  gaseous atom in its neutral state.
• First Electron Affinity
• A(g) e- ? A-(g)
• Second Electron Affinity
• A-(g) e- ? A2-(g)
• Third Electron Affinity
• A2-(g) e- ? A3-(g)

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Electronegativity

• A measure of the ability of an atom in a chemical
  compound to attract electrons.
• Fluorine, the most electronegative element, is
  arbitrarily assigned a value of 4.0. Values for
  other elements are calculated in relation to
  this.
• Tend to increase across a period
• Tend to decrease down a group or remain about the
  same.
• If an element does not form a compound, some
  noble gases, will not have a value.

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Trends in Electron Affinity and Electronegativity

• Both electron affinity and electronegativity
  increase from L to R across a period.
• Both electron affinity and electronegativity
  decrease down a group.

26
Two Factors Used to Explain Trends

• The principal energy level
• All other factors being equal, increased n for
  the orbitals in which electrons are found means
  increased size of orbitals, which leads to
  decreased attraction for electrons from the
  nucleus.

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Effective Nuclear Charge

• Effective charge is the approximate net nuclear
  charge felt by the highest energy electrons.
• All other factors being equal, increased
  effective charge means increased attraction for
  electrons, which leads to decreased size of
  orbitals.
• Effective charge depends upon two factors
• Total nuclear charge of protons (greater the
  total nuclear charge, higher the attraction felt
  by electrons)
• of shielding electrons (e present in between
  the nucleus and the valence shell electrons, the
  higher the number of shielding electrons, the
  lesser is the effective nuclear charge)

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SHIELDING

• The net nuclear charge felt by an outer electron
  is substantially lower than the actual nuclear
  charge. the outer electrons are shielded from the
  full charge of the nucleus by the inner
  electrons, which is called shielding effect.

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Explanation of Trends (1)
30
Explanation of Trends
31
Explanation of Trends
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Bibliography

• http//www.homeoint.org/morrell/british/originidea
  .htm
• http//www.chemsoc.org/viselements/pages/history.h
  tml
• http//www.library.upenn.edu/etext/smith/d/doberei
  ner.html
• http//www.library.upenn.edu/etext/smith/n/newland
  s.html
• http//www.ulb.ac.be/sciences/cudec/ressources/Men
  deleev.gif
• http//intro.chem.okstate.edu/1314F00/Lecture/Chap
  ter7/ATRADIID.DIR_PICT0003.gif
• http//scidiv.bcc.ctc.edu/wv/4/0004-000-IE.GIF