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									                               The Periodic Law Notes (Chapter 5)

I. History of the Periodic Table – About 70 elements were known by 1850 (no noble gases) but
there didn’t appear to be a good way of arranging or relating them to study.

A. Mendeleev and Chemical Periodicity – Mendeleev placed known information of elements on
cards (atomic mass, density, etc…). He arranged them in order of increasing atomic masses,
certain similarities in their chemical properties appeared at regular intervals. Such a repeating
pattern is referred to as periodic.
   1. Mendeleev’s table was published in 1869.
   2. He left blanks in his periodic table for undiscovered elements and he predicted their
   properties. Later elements were discovered with properties he predicted!
   3. Problems with his table – a few elements did not fit – the atomic mass arrangement did not
   match with other similar properties.
   4. Recognition – Mendeleev never received the Nobel Prize – the importance of the Periodic
   Table was not realized in his lifetime!

B. Mosleley and the Periodic Law
  1. Moseley, with Rutherford, in 1911 discovered a new pattern. The positive charge of the
     nucleus increased by one unit from one element to the next when elements are arranged as
     they are now on the Periodic Table
  2. This discovery led to definition of atomic number and the reorganization of the Periodic Table
  based on atomic number not atomic mass.
  3. The Periodic Law – the physical and chemical properties of the elements are periodic
  functions of their atomic numbers.
  4. Moseley died at the age of 28 – victim of WWI

C. The Modern Periodic Table - an arrangement of the elements in order of their atomic numbers
so that elements with similar properties fall in the same column (or group).
  Groups: vertical columns (#1-18)
  Periodic: horizontal rows (# 1-7)
  2. Periodicity – the similarities of the elements in the same group is explained by the
  arrangement of the electrons around the nucleus.

II. Electron Configuration and the Periodic Table

A. The s-block Elements: Groups 1 and 2
  1. Group 1: Alkali metals
      -soft silvery metals
      -most reactive of all metals, never found free in nature
      -reacts with water to form alkaline or basic solutions – store under kerosene
      -whenever you mix Li, Na, K, Rb, Cs, or Fr with water it will explode and produce an alkaline
       -ns1 (ending of all electron configurations for this group)

  2. Group 2: Alkaline earth metals
      -less reactive than Alkali, but still react in water to produce an alkaline solution
      -never found free in nature
      -harder, denser, stronger than alkali
      - ns2 (ending of all electron configurations for this group), because they have 2 electrons in
      the s sublevel, this makes them a little less reactive then the Alkali metals in group 1
B. The d-Block Elements: Groups 3-12
      - are all metals with metallic properties (malleability, luster, good conductors, etc…); are
      referred to as the Transition Metals
      -Harder and denser than alkali or alkaline
      -Less reactive than alkali or alkaline
      -For the most part their outermost electrons are in a d sublevel
      -Exceptions to the predicted electron configurations are found in these groups (Ex: Ni, Pd, Pt)

Irregular Electron configurations – sometimes the electron configuration is NOT what we would
predict it to be. Sometimes electrons are moved because (l) it will result in greater stability for that
atom or (2) for some unknown reason??

It is very important to define “stable” here. STABLE means:
        l. all degenerate (equal energy) orbital’s are FULL
        2. all degenerate orbital’s are half-full
        3. all degenerate orbital’s are totally empty.

Examples – draw the orbital notation (lines and arrows) for the predicted electron configuration for
Cr #24:

However, the real E. C. is [Ar]4s13d5. The 4s1 electron has been moved to achieve greater
The elements with “irregular” electron configurations are Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, La, Pt, Au,
Ac, Gd, Th, Pa, U, Np, and Cm. We will use the actual electron configuration (found on the periodic
table at the back of your book) instead of the predicted configuration for these elements. You do
NOT have to memorize these!

C. The p – Block Elements: Groups 13 – 18
      -Contain metals and nonmetals
      -Metalloids, along zigzag line, have characteristics of both metals and nonmetals (many are
      good conductors but are brittle). The metalloids are boron, silicon, germanium, arsenic,
      antimony, and tellurium.

       1. Group 17 - Halogens – most reactive nonmetals
       -7 electrons in outermost (s and p) energy levels (that is why they are so reactive – only need
       one electron to have 8)
       -called the salt formers (they react vigorously with metals to form salts). A salt is a metal and
       a nonmetal bonded together.
       -most are gases

       2. Group 18 - Noble gases – very unreactive
       - 8 electrons in outermost s and p energy levels
       - all are gases

The s and p blocks are called the main group or representative elements!

D. The f-Block Elements: Inner Transition Metals
  -final electrons fill an f sublevel

  1. Lanthanides – shiny reactive metals; Ce-Lu (fill the 4f sublevel)
  2. Actinides – unstable and radioactive; Th-Lr (fill the 5f sublevel)
Hydrogen and Helium - Oddballs
-Hydrogen is NOT an Alkali metal, it is a very reactive gas. It is placed with the Alkali metals
because 1s1 is its electron configuration.
-Helium is a Noble gas, it is unreactive, but it does not have 8 electrons in outermost energy level,
because it only has 2 total electrons!

For the test memorize the names, symbols, and placement of the
following elements: H-Kr (1-36), Hg, Ag, Au, Xe, Rn
III. Electron Configuration and Periodic Properties

A. Valence Electrons – electrons in the outermost (largest n) s and p orbital’s; electrons available
to be lost, gained, or shared in the formation of chemical compounds

When all of the valence orbital’s are full (have 8 electrons), the atom tends to be unreactive (like the
Noble Gases)
      # of Valence Electrons       Group #               Ending Configuration
              1                    1                     ns1 very reactive
              2                    2                     ns2
              3                    13                    np1
              4                    14                    np2
              5                    15                    np3
              6                    16                    np4
              7                    17                    np5 very reactive
              8                    18                    np6 very unreactive

                                                        B. Atomic Radii
          Atomic Radii Trend                            1. the size of an atomic radius cannot be
                                                        measured exactly because it does not have a
                                                        sharply defined boundary. However the
                                                        atomic radius can be thought of as ½ the
                                                        distance between the nuclei of identical atoms
                                                        joined in a molecule.

2. Period trend – atomic radii decrease as you move across a period.
       As you move across a period, from left to right, the size of an atom decreases because the
       nucleus is getting larger and more positive, but it is still pulling on the same number of energy
       levels. Example: Which atom is larger?           (A) Zr or Sn                   (B) Li or Cs

3. Group trend - atomic radii decrease as you move up a group (or increase as you go down a
       Shielding effect - an invisible barrier made of core electrons serve to decrease the pull of
       the nucleus on the outer (valence) electrons.
       Shielding increases as you go down a group because there are more core electrons.
       Shielding is considered to be constant as you move across a period because the number of
       energy levels is staying the same.
       Which atom has more shielding? (A) K or Ca              (B) Na or K
       Which atom is smaller? (A) N or P                (B) Li or K
                                                     C. Ionization Energy (IE) – a measure of the
                                                     ease with which an electron can be removed
                                                     from a gaseous atom or ion (sometimes called
             IE Trend                                ionization potential).

                                                     1. First ionization energy – the energy required
                                                     to remove one outermost electron from a
                                                     gaseous atom. Second ionization energy – the
                                                     additional amount of energy needed to remove
                                                     an outermost electron from the gaseous +1 ion.

2. Period trend – ionization energy increases as you move across the period.
   As you move across a period the ionization energy increases because the atoms get smaller.
   Another way to think of it: the number of valence electrons increases (the amount of energy
   needed to remove one electron is less then what is needed to remove 7 or 8 electrons).

3. Group trend – ionization energy increases as you move up a group (or decreases as you move
down a group).
   In general, as you do down a group the ionization energy decreases because the size of the
   atom is increasing and the outermost electrons are further from the nucleus.

Which atom has the higher first ionization energy?     (A) Hf or Pt           (B) Cl or Ar

                                                     D. Electron Affinity (EA) – the energy change
                                                     that accompanies the addition of an electron to
                                                     a gaseous atom. The desire of an atom to have
             EA Trend                                another electron. Chlorine for example would
                                                     like to have another electron and therefore has a
                                                     very high electron affinity.

                                                     1. Period trend – electron affinity increases as
                                                     you move across a period because atoms
                                                     become smaller and the nuclear charge
                                                     increases. This means there is a greater pull
                                                     from the nucleus.

  2. Group trend – electron affinity increases as you move up a group (or decreases as you move
  down a group) because the size of the atom increases and shielding increases. This means
  there is less pull from the nucleus.
Example: Which element has the greater electron affinity? Pb or Sn
                                                          E. Ionic Radii
       Ionic Radii Trend                                  1. Period trend – The size of ions decreases
                                                          as you move across a period because you
                                                          have more protons pulling on the same number
                                                          of energy levels.

                                                          2. Group trend – the size of the ions decrease
                                                          as you move up a group (or increase as you
                                                          move down a group) because the number of
                                                          energy levels increases.

Example: Which would be larger? K+1 or Ge +4                Which is larger? P ion or Cl ion

3. Metals – metal ions (cations) always have a smaller radius than its corresponding atom
a. it loses its outer energy level electrons (valence electrons)
b. the proton to electron ratio is greater in the ion than in the atom
        sodium atom 11p+/11e- = 1.0                sodium ion 11p+/10e- = 1.1
The value of p+ to e- ratio varies inversely to the size of the ion. The sodium ion is smaller because
it has a larger proton to electron ratio.

Which would be larger?        Potassium atom or potassium ion

4. Nonmetals - nonmetal ions are always larger than their corresponding atoms because:
a. repulsion between electrons
b. the p+ to e- ratio is less in the ion
       chlorine atom 17p+/17e- = 1.0             chlorine ion 17p+/18e- = 0.94
The chlorine ion is larger because it has a smaller proton to electron ratio.

Which would be larger? Sulfur atom or sulfur ion

                                                      F. Electronegativity (EN) – the tendency of an
                                                      atom to attract electrons to itself when it is
     Electronegativity Trend                          chemically combined (bonded) to another element.

                                                      In general, metals have low EN and nonmetals have
                                                      high EN. The actual amount of EN an atom has is
                                                      indicated by a number of the Pauling Electronegativity
                                                      Scale that goes from 0 to 4. Dr. Linus Pauling set up
                                                      this scale and gave the element having the greatest EN
                                                      an arbitrary number of 4, and he assigned numbers to
                                                      the others relative to this element.

Fluorine is the most electronegative element at 4. (3.98) and Francium is the least electronegative at 0.7.

1. Period trend - EN increases as you go across a period (excluding the noble gases) because size

2. Group trend - EN increases as you go up a group (or decreases as you go down a group) because there
is less pull from the nucleus as the electrons get further away.

Which would have the greater EN? (A) Ca or Se               (B) Be or O

Electronegativity enables us to predict what type of bond will be formed when two elements combine.
Homework (Periodic Table Day 1):
Periodic Table Set 1 Problems: Page 155 - 1bc, 3, 4, 6, 7a, 8, 9, 11, 13, 14, 15,
16a, 27ab, 28, 30, 40

Homework (Periodic Table Day 2):

                                 Periodic Table Trends Worksheet
Trends-Atomic size, Shielding, Ionization Energy, Electron Affinity

1. Circle the one from each pair that would be the larger in size:

(A) F atom or O atom                (B) Ba atom or Ra atom # 88            (C) Hf atom #72 or Ti atom #22

(D) Cs ion or Ba ion                (E) Al ion or Al atom                  (F) Po ion #84 or At ion #85

(G) I ion or I atom                 (H) Dy atom #66 or Cf atom #98         (I) As #33 atom or Cl atom

(J) Ca atom or Ca ion               (K) W atom #74 or Gd atom #74          (L) Mg ion or Na ion

(M) Sb-3 #51 or Sb +4               (N) Si+4 or P-3

2. Circle the element that has more shielding:

(A) B#5 or In #49                   (B) Mg or S or neither                 (C) Tl #81 or Y#39

3. Circle the element with the greater first ionization energy (IE):

(A) Pb or Sn                        (B) B or C                             (C) Ba or At #85

(D) Lr #103 or Ra #88               (E) Cs #55 or V #23                    (F) Si # 14 or Ag

4. Circle the element with the greater electron affinity (EA):

(A) F or Cl             (B) Sr #38 or Rb #37          (C) Os #76 or Co #27         (D) Am #95 or Eu #63

5. Circle the element with the lower electronegativity (EN):

(A) C or N                (B) Na or K                   (C) Ta #73 or Cu               (D) Pd #46 or Mo #42

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