Ionization Energy Within a Group

As the atomic radius increases from top to bottom in a group, the valence electrons become more further away from the nucleus and the nuclear attraction forces on these electrons decrease. Therefore, as the atomic radius increases, the amount of energy required to remove an electron decreases. As a result, we can say that within a group ionization energy of elements decrease from top to bottom. 

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Explain the trend in ionization energy within a group on the periodic table. 

Describe the trend in first ionization energies within a group. 

Therefore, ionization energy tends to decrease from top to bottom within a group, since the valence electrons (the first ones to be lost) are farther away from the nucleus. 

Why are the properties of elements within a group similar hut not identical Although elements within a group have the same number of valence electrons, they have different numbers of nonvalence electrons. Remember what happens as the atomic number increases within a group. As new levels of electrons are added, the atomic radius increases and the shielding effect increases. As a result, the ionization energy decreases. A lower ionization energy makes it easier for an element to lose electrons. 

The first ionization energies for the Group IIIA elements (B, Al, Ga, In, Tl) are exceptions to the general horizontal trends. They are lower than those of the IIA elements in the same periods because the IIIA elements have only a single electron in their outermost p orbitals. Less energy is required to remove the first p electron than the second s electron from the outermost shell, because the p orbital is at a higher energy (less stable) than an r orbital within the same shell (w value). 

A general trend within a group, the first ionization energy decreases down the group because in the same direction the atomic radii and principal quantum number n increase. There are only a few exceptions to this trend, and they are found in Groups 13 and 14. 

We have seen that the metallic character of the elements decreases from left to right across a period and increases from top to bottom within a group. On the basis of these trends and the knowledge that metals usually have low ionization energies while nonmetals usually have high electron affinities, we can frequently predict the outcome of a reaction involving some of these elements. 

In this section, we focus on three atomic properties that reflect the central importance of electron configuration and effective nuclear charge atomic size, ionization energy, and electron affinity. Most notably, these properties are periodic, which means they generally exhibit consistent changes, or trends, within a group or period. 

Note that the values tend to increase within each period, except for small drops in ionization energy at the Group IIIA and VIA elements. Large drops occur when a new period begins. 

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