The Periodic Table

1
The Periodic Table

• Elements are arranged in a way that shows a
  repeating, or periodic, pattern.
• Dmitri Mendeleev created the first periodic table
  of the elements in 1869.
• He ordered the 70 known elements by their atomic
  masses and their chemical properties.
• He found that some elements could not be put into
  groups with similar properties and at the same
  time stay in order.

2
Modern Periodic Table

• Later, Henry Moseley carried on the work.
• Moseley put the elements in order of increasing
  atomic NUMBER.
• He found that the position of the element
  corresponded to its properties.
• The modern periodic table shows the position of
  the element is related to
• Atomic number AND
• Arrangement of electrons in its energy levels

3
Electron Shells

• Move down P. table Principal quantum number (n)
  increases.
• Distribution of electrons in an atom is
  represented with a radial electron density graph.
• Radial electron density is probability of finding
  an electron at a particular distance from the
  nucleus.
• Electron shells are diffuse and overlap a great
  deal.

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Examples of Electron shells

• He 1s2
• Radial plot shows 1 maximum
• Ne 1s2 2s2 2p6
• Radial plot shows 2 maxima ( 1 each for the 1st
  and 2nd energy levels )
• Ar 1s2 2s2 2p6 3s2 3p6
• Radial plot shows 3 maxima ( 1 each for the
  1st,2nd and 3rd energy levels )

5
Atomic Sizes Single atoms

• Colliding argon atoms ricochet apart because
  electron clouds cannot penetrate one another to a
  significant extent.
• The apparent radii are determined by the closest
  distances separating the nuclei during such
  collisions.
• This radius is called the nonbonding radius.

6
Atomic Sizes Bonded atoms

• The distance between two nuclei is called the
  bond distance.
• If the two atoms making up the molecule are the
  same, then ½ the bond distance is called the
  bonding atomic radius of the atom.
• This radius is shorter than the nonbonding
  radius.

7
Atomic Sizes using Periodic Table

• As we move down a group, atoms become larger.
• Larger n more shells larger radius
• As we move across a period, atoms become smaller.
• More protons more effective nuclear charge,
  Zeff
• More positive charge increases the attraction of
  nucleus to the electrons in the outermost shell,
  so the electrons are pulled in more tightly,
  resulting in smaller radius

8
Ionization energy

• Ionization energy of an ion or atom is the
  minimum energy required to remove an electron
  from the ground state of the isolated gaseous
  atom or ion.
• The first ionization energy, I1 is the energy
  required to remove one electron from an atom.
• Na(g) ? Na(g) e-
• The 2nd ionization energy, I2, is the energy
  required to remove an electron from an ion.
• Na(g) ? Na2(g) e-
• Larger ionization energy, harder to remove
  electron.

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Periodic Trends in Ionization Energy

• Highest Fluorine
• Ionization energy decreases down a group.
• Easier to remove electrons that are farther from
  the nucleus.
• Ionization energy increases across a period.
• Zeff increases, so its harder to remove an
  electron.
• Exceptions Removing the 1st and 4th p electrons

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Electron Affinity

• Electron affinity is the energy change when a
  gaseous atom gains an electron to form a gaseous
  ion.
• Electron affinity Cl(g) e- ? Cl-(g)
• Ionization energy Cl(g) ? Cl(g) e-
• Affinity for reaction above is exothermic ?E
  -349 kJ/mol
• If adding the electron makes the species more
  stable, it will be exothermic.

Gain
Lose
11
Coulombs law

• Which law can best be used to explain why
  addition of an electron to the O2 ion is an
  endothermic process?
• Coulombs law The energy required for the
  process is necessary to overcome the
  electrostatic repulsion between the electron and
  the already negatively charged O2 ion.

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Ion size

• The oxide ion is isoelectronic (has exactly the
  same number and configuration of electrons) with
  neon, and yet O2 is bigger than Ne. Why?
• This is Coulomb's law at work. In any
  isoelectronic series the species with the highest
  nuclear charge will have the smallest radius.

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Metals

• Metallic character increases down a group and
  from left to right across a period.
• Metal properties
• Lustrous (shiny)
• Malleable (can be shaped)
• Ductile (can be pulled into wire)
• Conduct electricity
• Metal oxides form basic ionic solids
• Metal oxide water ? metal hydroxide
• Metal oxides react with acids to form salt and
  water

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Metals

• Metal oxides form basic ionic solids
• Metal oxide water ? metal hydroxide
• MgO(s) H2O(l) ? Mg(OH)2(s)
• Metal oxides react with acids to form salt and
  water
• MgO(s) 2HCl(aq) ? MgCl2(aq) H2O(l)
• Most neutral metals are oxidized rather than
  reduced.
• Metals have low ionization energies.

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Metal reactivity

• Which of the alkali metals would you expect to
  react most violently with water? Li, Na, K, Rb
• Of these four, rubidium has the lowest ionization
  energy, making it the most reactive. Rubidium
  reacts explosively with water.

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Nonmetals

• Lower melting points than metals
• Diatomic molecules are nonmetals.
• Most nonmetal oxides are acidic
• Nonmetal oxide water ? acid
• P4O10(s) 6H2O(l) ? 4H3PO4(aq)
• Nonmetal oxides react with bases to form salt and
  water
• CO2(g) 2NaOH(aq) ? Na2CO3(aq) H2O(l)

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Nonmetallic oxides

• Which nonmetallic oxide would you expect to be
  the strongest acid? NO2, N2O, N2O4, N2O5
• N2O5 Nitrogen has an oxidation state of 5 in
  this compound. In general, the higher the
  oxidation state of the nonmetal, the more acidic
  the nonmetal oxide.

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General Trend Summary
Electronegativity, Ionization Energy, Electron
Affinity
F
Atomic Radius, Metallic Character
Electronegativity, Ionization Energy, Electron
Affinity
Atomic Radius, Metallic Character
Fr