Why do the metallic elements of a given period (horizontal row) typically have much lower ionization energies than do the nonmetallic elements of the same period?
Typically, the outside electrons in the non-metals "see" a higher positive charge and that makes it harder to dislodge the electron. For example, look at Na and Cl. Na has 11+ charges in the nucleus; Cl has 17+.
11* 2e shell 8e shell 1e shell
17* 2e shell 8e shell 7e shell.
So the Na (11+ charges) pulls on that outside single electron by about 1 charge (11 + shielded by 10 e = about 1)
The Cl sees about (17+ charges shielded by 10 e = about 7+ charges); therefore, the outside electrons in Cl are harder to pull away, thus a higher ionization energy.
The metallic elements of a given period typically have much lower ionization energies than the nonmetallic elements of the same period due to the following reasons:
1. Effective nuclear charge: The metallic elements have fewer electrons in their outermost energy level, known as valence electrons, compared to the nonmetallic elements. This results in a weaker effective nuclear charge experienced by the valence electrons, making it easier to remove them and lowering the ionization energy.
2. Atomic size: The metallic elements tend to have larger atomic sizes than the nonmetallic elements in the same period. This means that the valence electrons in metallic elements are relatively farther away from the positively charged nucleus, resulting in a weaker attractive force between the valence electrons and the nucleus. Consequently, it requires less energy to remove an electron and therefore, lower ionization energy.
3. Electron shielding: The metallic elements have greater electron shielding effects compared to the nonmetallic elements in the same period. Electron shielding is the phenomenon where the inner electron shells shield the valence electrons from the full effect of the positive nuclear charge. This reduces the attraction between the valence electrons and the nucleus, making it easier to remove the valence electrons and thus, lowering the ionization energy.
Overall, these factors collectively contribute to the lower ionization energies of metallic elements compared to nonmetallic elements in the same period.
The trend of metallic elements in a given period having lower ionization energies than nonmetallic elements can be explained by their electronic configuration and atomic structure.
Ionization energy is the energy required to remove an electron from an atom, resulting in the formation of a positively charged ion. So, elements with lower ionization energies require less energy to lose an electron.
Metallic elements tend to have lower ionization energies because they have a larger atomic radius and fewer valence electrons in their outermost shell compared to nonmetallic elements. This is due to their position on the periodic table.
In a period, the atomic radius generally decreases from left to right. This means that the metallic elements, found on the left side of the periodic table, have larger atomic radii compared to nonmetallic elements. The larger atomic radius allows metallic atoms to hold their valence electrons less tightly, making it easier for them to lose an electron and form positive ions.
Additionally, metallic elements often have a weak attraction between their positively charged nucleus and the valence electrons due to their relatively low effective nuclear charge. This weak attraction makes it easier for the valence electrons to be removed, resulting in lower ionization energies.
On the other hand, nonmetallic elements, found on the right side of the periodic table, have smaller atomic radii and higher effective nuclear charges. The smaller atomic radius means that the valence electrons are held more tightly, requiring a higher amount of energy to remove them and resulting in higher ionization energies.
In summary, the trend of metallic elements having lower ionization energies than nonmetallic elements in a given period can be explained by their larger atomic radius and weaker attraction between the nucleus and valence electrons. These factors make it easier for metallic elements to lose their valence electrons and form positive ions.