Electron Configuration and Periodicity

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Chapter 8

• Electron Configuration and Periodicity

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Development of Periodic Table

• Dmitri Mendeleev and Lothar Meyer independently
  came to the same conclusion about how elements
  should be grouped.

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Meyer vs Mendeleev

• Meyer group the elements according to their
  chemical and physical properties based on their
  volumes whereas Mendeleev based it on their
  masses.

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Development of Periodic Table

• Because Mendeleev was able to predict the
  existence of certain elements. His table of
  elements was adopted.
• Mendeleev predicted the discovery of germanium
  (which he called eka-silicon) as an element with
  an atomic weight between that of zinc and
  arsenic, but with chemical properties similar to
  those of silicon.

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Periodic Law

• Original The chemical and physical properties
  of the elements are periodic functions of their
  atomic masses
• Current The chemical and physical properties of
  the elements are periodic functions of their
  atomic numbers (THINK! PROTONS ELECTRONS)

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Atomic Properties and Periodic Trends

• There are a number of important physical and
  chemical properties of elements that the periodic
  table can put in perspective
• metallic character (includes luster,
  conductivity, malleability, and ductility)
• atomic radius half the distance between the
  nuclei of two like atoms
• ionic radius the radius of the nuclei after
  they have formed positive or negative ions

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Atomic Properties and Periodic Trends

• Ionization energies the energy required to
  remove electrons
• electron affinity the energy change that occurs
  when an atom gains an electron
• electronegativity the ability of an atom to
  attract a pair of electrons when bonded to
  another atom

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Atomic Size

• Atomic size generally increases as you move down
  a group of the periodic table. Electrons are
  added to higher principal energy levels, and
  nuclear charge increases. The outermost orbitals
  are larger, and the shielding by electrons at
  lower energy levels increases.

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• As you proceed across a period, the size of
  successive atoms decreases. How is this
  explained?
• With each successive element, the positive
  nuclear charge increases (as does the electron
  charge). The electrons are being added to the
  same principal energy level, so the force of
  attraction between the nucleus and the electron
  increases, leading to a decrease in size.

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• Effective nuclear charge the net nuclear charge
  felt by an electron after shielding from other
  electrons in the atom is taken into account.
  Zeff Zact ? Zshield.
• In any period, the number of electrons between
  the nucleus and the principal energy level is the
  same for all elements. The shielding effect of
  these electrons is a constant within a period.

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Atomic Radius
Fig. 8.17 Atomic Radii for Main Group Elements

• Atomic radii actually decrease across a row in
  the periodic table. Due to an increase in the
  effective nuclear charge.
• Within each group (vertical column), the atomic
  radius tends to increase with the period number.

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Example 1

• Choose the larger atom in each pair
• Na or Si
• P or Sb
• Al or Cl
• Al or In

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Ionic Radius

• If positively charged the radius decreases
• If negatively charged the radius increases
  (relative to the atom).
• When substances have the same number of electrons
  (isoelectronic), the radius will depend upon
  which has the largest number of protons.

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Example 2

• Choose the larger particle in each pair
• Na or Na
• Co3 or Co2
• Al3 or Al

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Example 3

• Predict which of the following substances has
  the largest radius P3?, S2?, Cl?, Ar, K, Ca2.

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IONIZATION ENERGY

• Ionization energy, Ei minimum energy required to
  remove an electron from the ground state of atom
  (molecule) in the gas phase. M(g) h? ? M e?.
• Ei related to electron configuration. Higher
  energies stable ground states.
• Sign of the ionization energy is always positive,
  i.e. it requires energy for ionization to occur.

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• The ionization energy is inversely proportional
  to the radius and directly related to Zeff.
• Exceptions to trend
• B, Al, Ga, etc. their ionization energies are
  slightly less than the ionization energy of the
  element preceding them in their period.
• Before ionization ns2np1.
• After ionization is ns2. Higher energy ? smaller
  radius.
• Group 6A elements.
• Before ionization ns2np4.
• After ionization ns2np3 where each p electron in
  different orbital (Hunds rule).
• Electron-electron repulsion by two electrons in
  same orbital increases the energy (lowers EI).

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Ionization Energy Periodic table
Ionization Energy vs atomic
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Example 4

• Choose the atom with te larger ionization energy
  in each pair
• B or C
• O or S
• Cl or As

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HIGHER IONIZATION ENERGIES

• The energies for the subsequent loss of more
  electrons are increasingly higher. For the
  second ionization reaction written as
• M(g) h? ? M2 e? Ei2.
• Large increases in the ionization energies vary
  in a zig-zag way across the periodic table.
• States with higher ionization energies have
  1s22s22p6 (stable).

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Example 5

• Use the periodic table to select the element in
  the following list for which there is the largest
  difference between the second and third
  ionization energies.

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Electronegativity

• Electronegativity is the tendency of atoms of an
  element to attract electrons when they are
  chemically combined with another element. They
  are expressed in a qualitative measurement called
  the Pauling electronegativity scale. Ionization
  energy and electron affinity are used to
  calculate electronegativity.

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Electronegativity Trends

• Going across a period, electronegativity
  increases. Metals have low electronegativities,
  non-metals high electronegativities.
• Electronegativity decreases going down a group.
  The ability of an atom to attract electrons
  decreases with increasing nuclear charge.
• What is electronegativity good for? Predicting
  the type of bonding that exists between atoms in
  compound. The term was first coined by Linus
  Pauling in his landmark 1932 paper, "The Nature
  of the Chemical Bond."

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Example 6

• Predict the order of increasing electronegativity
  in each of the following groups of elements.
• C, N, O
• S, Se, Cl
• Si, Ge, Sn

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ELECTRON AFFINITY

• Electron Affinity, Eea, is the energy change that
  occurs when an isolated atom in the gas phase
  gains an electron.
• E.g. Cl e? ? Cl? Eea ?348.6 kJ/mol
• Energy is often released during the process.

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• Magnitude of released energy indicates the
  tendency of the atom to gain an electron.
• From the data in the table the halogens clearly
  have a strong tendency to become negatively
  charged
• Inert gases and group I II elements have a very
  small Eea.

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Example 7

• Order the atoms in each of the following sets
  from the least exothermic electron affinity to
  the most
• N, O, F
• F, Cl, Br, I

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Magnetic Properties

• Although an electron behaves like a tiny magnet,
  two electrons that are opposite in spin cancel
  each other. Only atoms with unpaired electrons
  exhibit magnetic susceptibility.

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• A paramagnetic substance is one that is weakly
  attracted by a magnetic field, usually the result
  of unpaired electrons.
• A diamagnetic substance is not attracted by a
  magnetic field generally because it has only
  paired electrons.

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Example 8

• Indicate whether the following are diamagnetic or
  paramagnetic
• Fe
• Ca
• Sb