As described above, ionization energies are dependent upon the atomic radius. Since going from right to left on the periodic table, the atomic radius increases, and the ionization energy increases from left to right in the periods and up the groups.
Exceptions to this trend is observed for alkaline earth metals group 2 and nitrogen group elements group Typically, group 2 elements have ionization energy greater than group 13 elements and group 15 elements have greater ionization energy than group 16 elements. Groups 2 and 15 have completely and half-filled electronic configuration respectively, thus, it requires more energy to remove an electron from completely filled orbitals than incompletely filled orbitals.
Alkali metals IA group have small ionization energies, especially when compared to halogens or VII A group see diagram 1. In addition to the radius distance between nucleus and the electrons in outermost orbital , the number of electrons between the nucleus and the electron s you're looking at in the outermost shell have an effect on the ionization energy as well.
This effect, where the full positive charge of the nucleus is not felt by outer electrons due to the negative charges of inner electrons partially canceling out the positive charge, is called shielding.
The more electrons shielding the outer electron shell from the nucleus, the less energy required to expel an electron from said atom. The higher the shielding effect the lower the ionization energy see diagram 2. It is because of the shielding effect that the ionization energy decreases from top to bottom within a group.
From this trend, Cesium is said to have the lowest ionization energy and Fluorine is said to have the highest ionization energy with the exception of Helium and Neon. Each succeeding ionization energy is larger than the preceding energy. Electron orbitals are separated into various shells which have strong impacts on the ionization energies of the various electrons.
For instance, let us look at aluminum. This is usually determined by how easily electrons can be removed we call it ionization energy! Elements with high electronegativity will be very reactive, as will elements with low ionization energy. Alkali metals, for example, are very reactive but there is difference between sodium and ceasium. Ceasium has lower ionization energy so it is VERY reactive.
Sodium has greater ionization energy than ceasium so it is not so reactive as caesium but it's still very reactive. Add a comment. Active Oldest Votes. Improve this answer. Zuhair Zuhair 11 1 1 bronze badge. Sign up or log in Sign up using Google. Sign up using Facebook. Sign up using Email and Password. Post as a guest Name. Email Required, but never shown. Featured on Meta. Now live: A fully responsive profile. Version labels for answers.
Linked 1. Ionization energies measure the tendency of a neutral atom to resist the loss of electrons. It takes a considerable amount of energy, for example, to remove an electron from a neutral fluorine atom to form a positively charged ion.
The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion. A fluorine atom in the gas phase, for example, gives off energy when it gains an electron to form a fluoride ion. Electron affinities are more difficult to measure than ionization energies and are usually known to fewer significant figures.
The electron affinities of the main group elements are shown in the figure below. Several patterns can be found in these data. Electron affinities generally become smaller as we go down a column of the periodic table for two reasons. First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom.
Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger. Electron affinity data are complicated by the fact that the repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. As a result, these elements have a smaller electron affinity than the elements below them in these columns as shown in the figure below.
From that point on, however, the electron affinities decrease as we continue down these columns. At first glance, there appears to be no pattern in electron affinity across a row of the periodic table, as shown in the figure below. When these data are listed along with the electron configurations of these elements, however, they make sense. These data can be explained by noting that electron affinities are much smaller than ionization energies. As a result, elements such as helium, beryllium, nitrogen, and neon, which have unusually stable electron configurations, have such small affinities for extra electrons that no energy is given off when a neutral atom of these elements picks up an electron.
These configurations are so stable that it actually takes energy to force one of these elements to pick up an extra electron to form a negative ion. There is no doubt that sodium reacts vigorously with chlorine to form NaCl.
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