Using grammatically correct sentences: a) provide a possible explanation for why \(\mathrm{IE}_{1}\) for \(\mathrm{He}\) is greater than \(\mathrm{IE}_{1}\) for \(\mathrm{H}\). b) provide a possible explanation for why \(\mathrm{IE}_{1}\) for \(\mathrm{Li}\) is less than \(\mathrm{IE}_{1}\) for He.

The periodic table is a tabular arrangement of chemical elements, organized by their atomic number, electron configurations, and recurring chemical properties. It is divided into periods (horizontal rows) and groups or families (vertical columns). As you move from left to right across a period, the elements become less metallic, and their ionization energies generally increase. This is because the positively charged nucleus exerts a stronger pull on the electrons due to the increasing number of protons, making it more difficult to remove an electron. In contrast, moving down a group, the ionization energy tends to decrease. That's because additional electron shells are added, which increases the atomic size and distances the outer electrons from the nucleus.

Understanding the layout of the periodic table helps us predict various properties of elements, including ionization energy, based on their relative positions. For instance, helium (He), being in the first period and eighteenth group, has a higher ionization energy compared to hydrogen (H) in the same period but in the first group. This is a direct consequence of the organized patterns within the periodic table.

Atomic size or atomic radius is a measure of the distance from the nucleus to the boundary of the surrounding cloud of electrons. Generally, as you move from top to bottom in a group on the periodic table, the atomic size increases. This is due to the addition of electron shells, which places the outer electrons farther from the nucleus. Conversely, moving from left to right across a period causes the atomic size to decrease; the electrons are pulled closer to the nucleus as the increased number of protons enhances the nuclear charge.

In the context of ionization energy, atomic size plays a crucial role. A larger atomic size typically means that the outer electrons are further from the influence of the nucleus and can be removed more easily, resulting in lower ionization energy. Therefore, lithium (Li), which has a larger atomic size than helium (He), also has a lower ionization energy.

Nuclear attraction refers to the force exerted by the positively charged nucleus on the negatively charged electrons around it. The strength of this attraction affects how tightly the electrons are held and, consequently, how much energy is required to remove an electron (ionization energy). A greater number of protons in the nucleus results in a stronger nuclear attraction, leading to a higher ionization energy.

Comparatively, helium (He) demonstrates a greater nuclear attraction than hydrogen (H) due to the larger number of protons in its nucleus. This intense pull on the electrons is why helium has a higher ionization energy. It requires more energy to overcome this strong attraction and remove an electron from a helium atom than from a hydrogen atom. Nuclear attraction is a fundamental concept that helps explain the trend of ionization energy across the periodic table.

Electron shielding is a phenomenon where inner electrons reduce the effective nuclear charge felt by the outermost electrons. This occurs because the inner electrons repel the outer electrons, which consequently shield them from the full attractive force of the nucleus. As a result, the further an electron is from the nucleus, the less nuclear attraction it experiences due to increased electron shielding. This leads to a lower ionization energy because the outer electrons can be removed more easily.

Lithium (Li) exhibits electron shielding with its second-level electron. This electron is shielded by the two innermost electrons, reducing the pull from the nucleus compared to an element like helium (He), which has electrons only in the first shell and experiences no electron shielding. Consequently, lithium's ionization energy is lower than that of helium. Recognizing the role of electron shielding is essential to understanding the variations in ionization energy among different elements.