Periodic Properties of the Elements

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Periodic Properties of the Elements
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The Periodic Table

• The modern periodic table was developed in 1872
  by Dmitri Mendeleev (1834-1907). A similar table
  was also developed independently by Julius Meyer
  (1830-1895).
• The table groups elements with similar
  properties (both physical and chemical) in
  vertical columns. As a result, certain
  properties recur periodically.

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The Periodic Table

• Mendeleev left empty spaces in his table for
  elements that hadnt yet been discovered. Based
  on the principle of recurring properties, he was
  able to predict the density, atomic mass, melting
  or boiling points and formulas of compounds for
  several missing elements.

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The Periodic Table
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The Periodic Table
metal/non-metal line
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The Periodic Table

• The periodic table is based on observations of
  chemical and physical behavior of the elements.
  It was developed before the discovery of
  subatomic particles or knowledge of the structure
  of atoms.
• The basis of the periodic table can be
  explained by quantum theory and the electronic
  structure of atoms.

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Periodic Trends

• Many of the properties of atoms show clear
  trends in going across a period (from left to
  right) or down a group.
• In going across a period, each atom gains a
  proton in the nucleus as well as a valence
  electron.

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Periodic Trends

• The increase of positive charge in the nucleus
  isnt completely cancelled out by the addition of
  the electron.
• Electrons added to the valence shell dont
  shield each other very much. As a result, in
  going across a period, the effective nuclear
  charge (Zeff) increases.

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Effective Nuclear Charge

• The effective nuclear charge (Zeff) equals the
  atomic number (Z) minus the shielding factor (s).
• Zeff Z-s

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Effective Nuclear Charge

• The effective nuclear charge (Zeff) equals the
  atomic number (Z) minus the shielding factor (s).
• Zeff Z-s
• Within the valence shell, the shielding factor
  is approximately 0.35, so going across a period
  results in an increase in Zeff of roughly .65
  for each element.

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Effective Nuclear Charge

• Zeff Z-s

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Effective Nuclear Charge

• Electrons in the valence shell are partially
  shielded from the nucleus by core electrons.

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Effective Nuclear Charge

• Electrons in p or d orbitals dont get too
  close to the nucleus, so they are less shielding
  than electrons in s orbitals. As a result,
  effective nuclear charge increases across a
  period.

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Periodic Trends
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Periodic Trends

• In going down a group or family, a full quantum
  level of electrons, along with an equal number of
  protons, is added.
• As n increases, the valence electrons are, on
  average, farther from the nucleus, and experience
  less nuclear pull due to the shielding by the
  core electrons. As a result, Zeff decreases
  slightly going down a group.

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Trends- Atomic Radii

• Atomic radii are obtained in a variety of ways
• 1. For metallic elements, the radius is half the
  internuclear distance in the crystal, which is
  obtained from X-ray data.
• 2. For diatomic molecules, the radius is half
  the bond length.
• 3. For other elements, estimates of the radii
  are made.

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Trends- Atomic Radii

• Atomic radii follow trends directly related to
  the effective nuclear charge. As Zeff increases
  across a period, the electrons are pulled closer
  to the nucleus, and atomic radii decrease.
• As Zeff decreases down a group, the valence
  electrons experience less nuclear attraction, and
  the radius increases.

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Trends- Atomic Radii

• Atomic size roughly halves across a period, and
  doubles going down a group.

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Trends Ionization Energy

• Ionization energy is the energy required to
  remove an electron from a mole of gaseous atoms
  or ions.
• X(g) energy ? X(g) e-
• Elements can lose more than one electron, so
  there are 1st, 2nd, 3rd, etc., ionization
  energies.

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Ionization Energy

• It always requires energy to remove an electron
  from a neutral atom.
• As more electrons are removed and the ion
  becomes positively charged, it requires
  increasingly greater energy to remove electrons.

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Trends Ionization Energy

• Ionization energy is a measure of how tightly
  the electrons in the highest occupied orbitals
  are held by the nucleus. As a result, it is
  directly related to the effective nuclear charge.
• Ionization energy increases going across a
  period, and decreases going down a group.

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Trends Ionization Energy
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Trends Ionization Energy
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Ionization Energy
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Trends Electron Affinity

• Electron Affinity involves the addition of an
  electron to a mole of gaseous atoms. There are
  different conventions to defining electron
  affinity. Your text defines the EA as the energy
  released during the following process
• X(g) e- ? X-(g)

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Electron Affinity

• Your text defines the EA as the energy released
  during the following process
• X(g) e- ? X-(g)
• A positive value for EA indicates that the
  process releases significant energy. Thus, the
  halogens tend to have high electron affinities.

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Trends Electron Affinity

• There is less of a predictable trend in
  electron affinities. In going across a period
  (ignoring the noble gases), the electron affinity
  should become more negative. Although this is
  observed, there are many inconsistencies.

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Trends Electron Affinity
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Trends - Electron Affinity
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Trends- Electron Affinity

• In going down a group, the electron affinity
  should become become smaller. Although this
  trend is observed, there is only a slight change
  in electron affinities within a group. There may
  also be inconsistencies in the general trend.

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Electron Affinity

• It should be noted that the addition of a
  second electron to an anion is always highly
  unfavorable. The electron affinity of oxygen is
  141 kJ/mol to form O. Addition of the second
  electron to form the oxide ion (O2) requires 744
  kJ/mol.

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Metallic Character

• Metals are shiny, malleable and ductile. They
  are generally good conductors of heat and
  electricity, and low ionization energies.
• In reaction with non-metals, metals tend to
  lose electrons and form cations.

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Metallic Character
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Metallic Character

• Across a period, metallic behavior decreases.
  Non-metals are often crumbly solids, liquids or
  gases at room temperature.

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Metallic Character

• Metallic behavior increases going down a group.

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Electron Configurations of Ions

• The atoms of the main group elements (groups
  IA-VIIA) will form ions by losing or gaining
  electrons. The resulting ion will have the same
  electron configuration as a noble gas (group
  VIIIA). These configurations are usually very
  stable.

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Electron Configurations of Ions

• Atoms or ions with the same electron
  configuration (or number of electrons) are called
  isoelectronic.
• For example, Na, Mg2, Ne, F-, and O2- are
  isoelectronic. The size will decrease with
  increasing positive charge.
• O2- gt F- gtNegt Nagt Mg2

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Electron Configurations of Ions

• When atoms lose electrons, the electrons are
  always removed from the highest quantum level
  first.
• For the first row of transition metals, this
  means that the electrons in 4s subshell are lost
  before the 3d subshell.
• Fe Ar4s23d6 Fe2 Ar 3d6 or Ar4s03d6

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Common Ionic Charges

• The charges of ions of elements in groups 1A-7A
  (the main groups) are usually predictable.
• Group 1A metals form 1 ions, group 2A metals
  form 2 ions, etc.
• The non-metals of group 5A have a -3 charge,
  those of group 6A have a -2 charge, and the
  halogens form ions with a -1 charge.

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Typical Ionic Charges
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Trends Ionic Size

• Cations are always smaller than the neutral
  atom. The loss of one or more electrons
  significantly increases Zeff, resulting in the
  valence electrons being pulled closer to the
  nucleus.

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Ionic Size - Cations
Within a group, assuming the same ionic charge,
the size of the ion increases going down the
group, due to more core electrons shielding the
nucleus as n increases.
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Trends Ionic Size

• Across period, the cations get more positive,
  and as a result, considerably smaller.

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Trends Ionic Size

• Anions are always larger in size than the
  neutral atom. The addition of one or more
  electrons results in greater electron-electron
  repulsion, which causes the valence electrons to
  spread out a bit.

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Size of Anions
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• Anions are always larger than the neutral atom.

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Size of Anions

• Within a group, assuming the same ionic charge,
  the size of the ion increases going down the
  group, due to more core electrons shielding the
  nucleus as n increases.

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Trends Ionic Size
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Group IA the Alkali Metals
In discussing the chemistry, preparation and
properties of the group IA elements, it is
important to remember that hydrogen is not a
group IA metal. Its properties and reactivity
would place it within group 7A (diatomic
non-metals), rather than group IA.
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Group 1A Metals

• The group 1A metals are soft shiny metals with
  fairly low densities (Li, Na and K are less dense
  than water) and low melting points. Sodium melts
  at 98oC, and cesium melts at 29oC.
• The softness, low density and low melting
  points are the result of weaker metallic bonding
  due to only one valence electron in this group.

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Group 1A Metals - Production

• Due to the high reactivity with oxygen and
  water, all of the metals are found in nature in
  ionic form (M1).
• The pure metal must be produced in an oxygen
  and water-free environment. Typically, an
  electrical current is passed through the melted
  chloride salt. The metal and the chlorine gas
  are collected separately.

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Reactivity Trends

• The chemical behavior of the group IA metals
  illustrates periodic trends. As the valence
  electron occupies a higher quantum level, it
  experiences less nuclear attraction, and is more
  easily removed.

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Group 1A Metals Water

• The reaction with water forms hydrogen gas and
  the aqueous metal hydroxide. The reaction is so
  vigorous, that the hydrogen may ignite.
• 2 M(s) 2 H2O(l) ? H2(g) 2 MOH(aq)

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Metallic Character

• The group IA metals react with water to produce
  hydrogen and the metal hydroxide.

Metallic behavior increases going down a group.