MAIN GROUP ELEMENTS (GROUPS 1, 2, & 13 - 18)

GROUP

NAME

ENDING e-
CONFIG-
URATION

# OF
VALENCE
e-

e- DOT
DIAGRAM

WANTS TO (lose or
gain) e- TO
BE LIKE
NOBLE GAS?

CHARGE
OF
ION

OXID-
ATION
#

GETS (smaller or larger) WHEN ATOM TO ION?

1
or
I A

 

 

 

 

 

 

 

 

2
or
II A

 

 

 

 

 

 

 

 

13
or
III A

 

 

 

 

 

 

 

 

14
or
IV A

 

 

 

 

 

 

 

 

 

15
or
V A

 

 

 

 

 

 

 

 

 

16
or
VI A

 

 

 

 

 

 

 

 

 

17
or
VII A

 

 

 

 

 

 

 

 

 

18
or
VIII A

 

 

 

 

 

 

 

 

 

 

ION:
UNIT 5 - THE PERIODIC LAW

History of the Periodic Table

I.  Mendeleev and Chemical Periodicity
    A.  Wanted to organize elements according to their ____________________
    B.  When elements were arranged in order of increasing atomic mass*, similarities in chemical
         properties appeared at regular intervals (____________________)
    C.  *Several elements did not quite fit this pattern - Mendeleev put elements with similar
         ____________________ in the same column or group
    D.  1871 - Mendeleev predicted the existence and properties of several (then undiscovered)
         elements.  These elements were:
    E.  Within 15 years, those elements with those properties had been discovered

II.  Moseley and the Periodic Law
    A.  When elements were arranged in order of increasing ___________________________________,
         there was a distinct regular pattern.
    B.  ____________________:  The physical and chemical properties of the elements are periodic
        functions of their atomic numbers.
    C.  In other words, when elements are arranged in order of increasing atomic number, elements
         with similar properties appear at regular intervals.
    D.  Bottom line = elements in the same group have similar properties

III.  Modern Periodic Table:  arrangement of the elements in order of their atomic numbers so that elements with
       similar properties fall in the same group

Electron Configuration and the Periodic Table

I.  Stability of Noble Gases
    A.  Noble gases undergo very few chemical reactions - why?
    B.  Highest occupied energy level contains ________________________________________
    C.  Electrons in the highest occupied energy level are what determines an element's _____________

II.  Periods and Blocks of the Periodic Table
    A.  Horizontal row = ____________________;  7 on modern Periodic Table
    B.  Length of period determined by the sublevels being filled in that period
    C.  Period 1:  only _____ sublevel being filled;  can hold a maximum of _____ electrons; 
         period contains _____ elements
    D.  Period 4:  ____, ____, and ____ sublevels being filled;  s can hold ___ electrons, d can hold ___
         electrons, & p can hold ___ electrons; total of _____ electrons;
         Period 4 contains _____ elements
    E.  Period can be determined from the element's electron configuration
        1.  Bromine:  [Ar] 4s2 3d10 4p5
        2.  Highest number in front of letter is the element's highest occupied _____________________ -
             tells which period the element is in
        3.  For bromine, _____ is highest number, so it is in Period _____

III.  The "s" block elements:  Groups 1 and 2
    A.  Group 1 - Alkali Metals
        1.  generalized outermost energy level (valence) electron configuration:
        2.  silvery appearance
        3.  soft enough to cut with a knife
        4.  not found in nature as free elements - they're always part of a compound
        5.  with increasing atomic number, melting point ______________________
    B.  Group 2 - Alkaline Earth Metals
        1.  generalized valence electron configuration:
        2.  harder, stronger, more dense than Group 1
        3.  also have higher melting points than Group 1
        4.  less reactive than Group 1, but still not found in nature as free elements
    C.  Exceptions:  Hydrogen and Helium
        1.  Hydrogen (H)
            a.  electron configuration:
            b.  properties do not resemble those of any other element on the periodic table
        2.  Helium (He)
            a.  electron configuration:
            b.  in Group 18 because

IV.  The "d" block elements: Groups 3 - 12
    A.  called
    B.  have typical metallic properties:  ductile, malleable, shiny, solid, conduct electricity
    C.  less reactive than "s" block elements
    D.  found in nature as free elements
    E.  usual ending of electron configuration:

V.  The "p" block elements:  Groups 13 - 18
    A.  "s" and "p" block elements together referred to as _____________________________ elements
    B.  ending electron configurations of __________________ through ___________________
    C.  properties vary greatly b/c there are metals, metalloids, and nonmetals
    D.  Group 17 - Halogens
        1.  most reactive nonmetals
        2.  seven electrons in outermost energy level
    E.  "p" block metals are harder and more dense than "s" block , but not as hard or dense as the "d"
         block metals


LOCATING MAIN GROUP ELEMENTS ON THE PERIODIC TABLE

Given the electron configuration or noble gas configuration for an element, it is possible to determine its location on the Periodic Table without actually looking at a Periodic Table. 

* To tell which period this element is in...
    ~  find the highest occupied energy level for this element
You can do this by...
    ~  finding the largest coefficient number
The largest coefficient number is the number of the period where the element is located.

* To tell which "block" this element is in...  (like "s" block, "p" block, "d" block, etc)
    ~  find the highest occupied sublevel for this element
You can do this by...
    ~  finding the last lowercase letter written
The last lowercase letter written in the configuration is the "block" where the element is located.

* To tell which group this element is in...
    ~  find the highest occupied energy level for this element
You can do this by...
    ~  finding the largest coefficient number
Then...
    ~  add up the exponents of the largest coefficient number
This gives you the number of valence electrons in the element.
You will then know that 1 valence electron indicates that the element is in Group 1, 2 valence electrons indicates that the element is in Group 2, 3 valence electrons indicates that the element is in Group 13, 4 valence electrons indicates that the element is in Group 14, 5 valence electrons indicates that the element is in Group 15, 6 valence electrons indicates that the element is in Group 16, 7 valence electrons indicates that the element is in Group 17, and 8 valence electrons indicates that the element is in Group 18.

Look at the following example.

EXAMPLE:                [Ar] 4s2 3d10 4p5

It is possible to tell the period, group, and "block" where this element is located.
*  Period --  largest coefficient number is 4, so element is in Period 4
*  Block--  last lowercase letter written is "p", so element is in "p" block
*  Group--  largest coefficient number is 4...  2 electrons in 4s, 5 electrons in 4p -->  total of 7 valence electrons,
                 so this element is in Group 17.


 

Periodic Trends
Electronegativity/Electron Affinity (EN/EA)
:  measure of how much an atom wants to gain an electron

EN/EA Left to Right across a Period:                INCREASES  (not including Noble Gases)

Why?
* Elements on the left side of the P.T. (metals) want to lose electrons.  Elements on the right side of
   the P.T. (nonmetals) want to gain electrons.  Trend does not include Noble Gases because these   
   elements do not want to lose or gain electrons.

EN/EA Top to Bottom in a Group:                   DECREASES

Why?
*
This interference (and resulting decreased “hold”) is referred to as the SHIELDING EFFECT.

Ionization Energy (IE):  amount of energy required to remove an atom’s most loosely held electron

IE Left to Right across a Period:            INCREASES

Why?
* Elements on the left side of the P.T. (metals) want to lose electrons.  Therefore, it will not require
   much energy to remove an electron.  Elements on the right side of the P.T. (nonmetals) want to gain
   electrons.  Consequently, a lot of energy will be needed to remove (take away) an electron.

 

 

 

IE Top to Bottom in a Group:                DECREASES

Why?
*

Atomic Radius (AR):  distance from the nucleus to the H.O.E.L.

AR Top to Bottom in a Group:     INCREASES

Why? 
* There are more occupied energy levels as you move towards the
   bottom of the P.T.

 

 

 

 

 

AR Left to Right across a Period: DECREASES

Why?
*

Metallic Character:  how easily an atom will lose valence electrons (easier to lose = more metallic = more reactive METAL)
Which metal loses its valence electron(s) most easily? Fr
Why?
* Francium has one valence electron.  It is more reactive than elements at the top of Group 1 because there are many inner shell electrons that decrease the attraction the nucleus has for the valence electrons.

 

Nonmetallic Character:  how easily an atom will gain electrons (easier to gain = more nonmetallic = more reactive NONMETAL)
Which nonmetal gains electron(s) most easily?           F
Why?
* Fluorine has seven valence electrons.  It is more reactive than elements at the bottom of Group 17 because there are only a few inner shell electrons.  Consequently, the nucleus has a strong attraction for other electrons.

 

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