Sunday, 7 February 2016

The Periodic Table and Electronic Structure

History

The Periodic Table was invented in 1969 by the Russian scientist Dmitri Mendeleev (at that time, Professor of Chemistry at St. Petersburg University).

He arranged the 63 elements (all that were known at that time) by increasing atomic mass, leaving gaps where he predicted that elements, that were undiscovered at that time, would fit into the table.

Mendeleev arranged is so that elements that have similar properties are in the same groups.


This is the Periodic Table supplied by www.cie.org.uk
So this is probably the most reliable one you can find!


Key Terms


Groups = the vertical columns in the Periodic Table

Periods = the horizontal rows in the Periodic Table

Modern Periodic Table

There are currently 118 known elements!

Elements are arranged by proton number in the modern Periodic Table and elements with similar priorities are still in the same groups.


Groups:

There are eight groups of elements:

  • Group I
  • Group II
  • Group III
  • Group VI
  • Group V
  • Group VI
  • Group VII
  • Group 0 (sometimes known as Group VIII)
Group I is known as the alkali metals, Group II as the alkaline earth metals, Group VII as the halogens and Group 0 as the inert gases or noble gases.

The big block of elements between Groups II and III are the transition metals.


Periods:

The periods are numbered 1-7, going down the periodic table.

Across each period the properties of the elements change:

Courtesy of Bryan Earl and Doug Wilford's
Cambridge IGCSE Chemistry Third Edition


Metals and non-metals:

The bold line that starts beneath boron divides the Periodic Table into two parts.
The elements on the left of the line are metals and the elements on the right are non-metals.
Metalloids is the name given to the elements that are on this dividing line, they are metals that have properties of both metals and non-metals.



Electronic Structure and The Periodic Table

The group that an element is in determines how many electrons it has in it's outer shell.

N.B. This does not apply to the elements in Group 0, they have either 2 or 8 electrons

So the elements in Group I have 1 electron in their outer energy level and the elements in Group II have 2 electrons in their outer energy level, and so on and so on.

As you move down a group, the metallic characteristics of the elements increases, this happens because the outer energy shell becomes further away from the nucleus, as do the electrons in it. So there's less attraction between the outer energy shell's electrons and the nucleus, due to distance, so the electrons in the outer shell are easier to loose.
This doesn't happen in Group VII, in this group the reactivity DECREASES as you go down the group.

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