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.
Trends - 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.
Trends - 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.
|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...||Flourine (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.
Trends - Ionization Energy
The "nth" ionization
energy is the energy required to remove "n" electron(s) from an element.
Ionization energy increases greatly as:
• successive electrons are removed. This is because the remaining electrons can be more strongly attracted the protons in the nucleus.
• atomic radius decreases. Thus, removing an electron from an atom becomes more difficult (requires more energy) from left to right across a period and from top to bottom in a group. In other words, the smaller an atom is, the more the electrons can be equally attracted to the nucleus. This makes ionization energy increase.
• energy levels are removed. For example, the first ionization energy for sodium is relatively small, since it is a lone electron in the 3s sublevel. Furthermore, removing that electron makes the highest energy level for sodium its second, which has a full 2s and 2p sublevels. Removing a second electron from sodium is significantly more difficult, as the remaining electrons are part of a noble gas configuration (octet).