Chemistry Reference

The Periodic Table
The Periodic Table − Physical and chemical properties as well as background information on all the elements.  Each element includes links to other sites that feature that element.
Atomic Radius
Within a group, atomic radius increases from top to bottom.  This is a consequence of increased energy levels as one moves down a group.  Increased energy levels equates to larger orbitals and therefore more room for electrons to travel.  Across a period, atomic radius decreases from left to right.  While it is true that the number of electrons increases from left to right, so does the number of protons.  Since there is no increase in energy level, orbital sizes should be expected to remain constant.  However, the attraction of the protons (and recall they are about 1820 times more massive than electrons) for the electrons shrinks the orbitals and makes the atom smaller.
Electronegativity
Electronegativity measures how strongly an atom will attract electrons to itself when bonded to another element.  The opposite of this is electropositivity.  There are some generalizations that can be made that predict whether an element will be electronegative or electropositive and how strong it will be in that regard.  The table given below covers these points.
  Electronegative Electropositive
Type of element Nonmetals Metals and metalloids
Atomic radius Smaller radii (top of groups/right of periods) Larger radii (bottom of groups/left of periods)
The closer you get to... Fluorine (most electronegative) Francium (most electropositive)
 

Consider the two periodic tables below.  Note the strong correlation between atomic radius and electronegativity.  Can you see that the elements with the smallest radii have the highest electronegativity values?  This is due to the absence, or at least weak, shielding effect shown by smaller elements. 

Ionization Energy
Ionization energy is a measure of the amount of energy needed to remove a specified number of electrons from an atom. Most often, the first ionization energy is discussed, and hence describes the energy needed to remove a single electron from an atom. However, there can be as many ionization energies for an element as there are electrons, although removing large numbers of electrons is not often feasible. Generally, the smaller an atom is (see the atomic radius table above), the more difficult it is to remove an electron. As a consequence, ionization energies for small elements tend to be very large. For large elements, shielding plays a significant role and makes it easier for electrons to be removed. Large elements tend to have small ionization energies. Valence electrons also have a role in determining ionization energy. Using barium as an example, one can see it is one of the larger elements of the periodic table. It is predicted to have low ionization energies, which it does when the first two electrons are removed. However, after removing two electrons, barium now has 54 electrons like xenon. Xenon is a noble gas, and although it is a large noble gas, it has a large ionization energy. This is because removing electrons from xenon destroys the octet, the s2p6 configuration that makes the noble gases stable.
footer

Home
Contact the website administrator.