The Periodic Table and Periodic law

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					                   The Periodic Table and Periodic law

In the 1700’s Lavoisier compiled a list of all known elements of the time.

See table 1

    The 1800’s brought large amounts of information and scientists
     needed a way to organize knowledge about elements
    John Newlands proposed an arrangement were elements are ordered
     by increasing atomic mass.
    Newlands noticed when elements are arranged by increasing atomic
     mass, their properties repeated every eight element.
    Meyer and Mendeleev both demonstrated a connection between
     atomic mass and elemental properties.
    Moseley rearranged the table by increasing atomic number and
     resulted in a clear periodic pattern.
    Periodic repetition of chemical and physical properties of the elements
     when they are arranged by increasing atomic number is called the
     periodic law.

Contributions to the classification of elements. See Table 6.2

The modern periodic table contains boxes which contain the elements name,
symbol, atomic number, and atomic mass. See fig

    Columns of elements are called groups.
    Rows of elements are called periods
    Elements in groups 1, 2, and 13 -18 possess a wide variety of
     chemical and physical properties and are called the representative
     elements.
    Elements in groups 3 – 12 are known as the transition metals.
    Elements are classified as metals, non-metals and metalloids.
    Metals are elements that are generally shiny when smooth and clean,
     solid at room temperature, and good conductors of heat and
     electricity.
    Alkali metals are all the elements in group 1 except hydrogen, and
     are very reactive.
    Alkaline earth metals are in group 2, and are also highly reactive.
 The transition elements are divided into transition metals and inner
  transition metals.
 The two sets of inner transition metals are called the lanthanide
  series and actinide series and are located at the bottom of the periodic
  table.
 Non-metals are elements that are generally gases, or brittle, dull
  looking solids, and poor conductors of heat and electricity.
 Group 17 is composed of highly reactive elements called halogens.
 Group 18 gases are extremely unreactive and commonly called noble
  gases.
 Metalloids have physical and chemical properties of both metals and
  non-metals such as silicon and germanium.

Main idea – Elements are arranged into different blocks in the periodic
table according to their electron configurations.

Organizing the elements by electron configuration

    Recall electrons in the highest principle energy level are called
     valence electrons.
    All group 1 elements have one valence electron. See table 6.3
    The energy level of an element’s valence electrons indicates the
     period on the periodic table in which it is found.
    The number of valence electrons for elements in group 13 – 18 is
     ten less than their group number

   The s-, p-, d-, and f- block Elements

 The shape of the periodic table becomes clearer if it is divided into
  blocks representing the atom’s energy sublevel being filled with
  valence electrons.
 S-block elements consist of group 1 and 2, and the element helium.
 Group 1 elements have a partially filled s orbital with one electron.
 Group 2 elements have a completely filled s orbital with two electrons
 After the s orbital is filled, valence electrons occupy the p-orbital
 Groups 13 – 18 contain elements with completely or partially filled p
  orbitals. See table 6.4
 The d-block contains the transition metals and is the largest block
    There are exceptions, but d-block elements usually have filled
     outermost s orbital, and filled or partially filled d-orbital.
    The five d orbitals can hold 10 electrons, so the d-block spans ten
     groups on the periodic table.
    The f-block contains the inner transition metals.
    f-block elements have filled or partially filled outermost s orbitals and
     filled or partially filled 4f and 5f orbitals.
    The 7f orbitals hold 14 electrons, and the inner transition metals span
     14 groups.

   Main Idea Trends among elements in the periodic table include their size
   and ability to lose or attract electrons.

Atomic Radius

    Atomic size is a periodic trend influenced by electron configuration.
    For metals, atomic radius is half the distance between adjacent nuclei
     in a crystal of the element.
    For elements that occur as molecules, the atomic radius is half the
     distance between nuclei of identical atoms.
    There is a general decrease in atomic radius from left to right, caused
     by increasing positive charge in the nucleus.
    Valence electrons are not shielded from the increasing nuclear charge
     because no additional electrons come between the nucleus and
     valence electrons.
    Atomic radius generally increases as you move down a group.
    The outermost orbital size increases down a group, making the atom
     larger

Ionic Radius

    An ion is an atom or bonded group of atoms with a positive or
     negative charge.
    When atoms lose electrons and form positively charged ions, they
     always become smaller for two reasons:
    The lose of a valence electron can leave an empty outer orbital
     resulting in a small radius
    Electrostatic repulsion decrease allowing the electrons to be pulled
     closed to the radius.
    When atoms gain electrons, they can become larger, because the
     addition of an electron increases electrostatic repulsion.
    The ionic radii of positive ions generally decrease from left to right,
     beginning with group 15 or 16.
    Both positive and negative ions increase in size moving down a group.

Ionization Energy

    Ionization energy is defined as the energy required to remove an
     electron from a gaseous atom.
    The energy required to remove the first electron is called the first
     ionization energy.
    Removing the second electron requires more energy, and is called the
     second ionization energy.
    Each successive ionization requires more energy, but it is not a steady
     increase.
    The ionization at which the large increase in energy occurs is related
     to the number of valence electrons.
    First ionization energy increases from left to right across a period
    First ionization energy decreases down a group because the atomic
     size increases and less energy is required to remove an electron farther
     from the nucleus.

    The octet rule states that atoms tend to gain, lose or share electrons in
     order to acquire a full set of eight valance electrons.
    The octet rule is useful for predicting what type of ions an element is
     likely to form.
    The electronegativity of an element indicates its relative ability to
     attract electrons in a chemical bond.
    Electronegativity decreases down a group and increase from left to
     right across a period.
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