Question

The correct value of ionisation energies (in \(\mathrm{kJ} \mathrm{mol}^{-1}\) ) of Si, P, CI and S respectively are: (a) \(786,1012,999,1256\) (b) \(1012,786,999,1256\) (c) \(786,1012,1256,999\) (d) \(786,999,1012,1256\)

Short answer

The correct answer is (d) 786, 999, 1012, 1256.

Step-by-step solution

Understanding Ionization Energy

Ionization energy is the energy needed to remove an electron from an atom. Elements with higher ionization energy require more energy to remove an electron.

Review Periodic Table Trends

Ionization energies generally increase across a period from left to right due to an increase in nuclear charge with the same electron shielding. Therefore, we expect the ionization energy to increase as we go from Si to P to S to Cl.

Identify Relative Ionization Energies

Among the given elements Si, P, Cl, and S, based on periodic table position, the order of increasing ionization energy is Si < P < S < Cl.

Match Ionization Energies to Values

We match the order Si < P < S < Cl to the numerical choices provided. Checking choice (d), we have Si = 786, P = 999, S = 1012, Cl = 1256, which matches the periodic trend we established.

Key concepts

Periodic Table Trends

The periodic table is a powerful tool for predicting and understanding the properties of elements. One such property is ionization energy, which tends to follow specific trends across the periodic table.
As one moves from left to right across a period, the ionization energy generally increases. This is because elements on the right have more protons in their nuclei compared to those on the left. More protons mean a stronger attraction between the nucleus and the electrons, which requires more energy to remove an electron. Moving down a group, ionization energy decreases due to increased electron shielding and greater distance from the nucleus.

• Across periods: Ionization energy increases.
• Down groups: Ionization energy decreases.

These trends help explain why elements like chlorine in the top right have high ionization energies compared to silicon, phosphorous, and sulfur.

Electron Shielding

Electron shielding describes how inner-shell electrons reduce the effective nuclear charge felt by outer-shell, or valence electrons. Even though the number of protons increases across a period, electrons that are added do little to shield this charge, thus the effective nuclear charge increases.
This shielding effect impacts the ionization energy. As we move across a period, the number of protons increases, resulting in a stronger nuclear pull on the valence electrons. However, within the same period, the electron shielding does not increase significantly, allowing the effect of the additional protons to be significant. Thus, the ionization energy increases.

• Inner electrons block nuclear charge.
• Effective nuclear charge increases across a period.
• Ionization energy rises due to less shielding.

Understanding this concept helps to grasp why elements further right in the same period require more energy for electron removal.

Nuclear Charge

Nuclear charge refers to the total charge of the protons in the nucleus of an atom. As the number of protons increases, so does the nuclear charge.
The nuclear charge has a direct impact on ionization energy. With a higher nuclear charge, the attraction between the nucleus and the electrons becomes stronger, leading to higher ionization energies. Therefore, as the nuclear charge increases, the ability of the atom to hold onto its electrons tightly also increases, making it harder to remove an electron.
This is why chlorine, with a higher atomic number than silicon, phosphorous, and sulfur, has the highest ionization energy among these elements. Effective nuclear charge, which is the perceived charge by valence electrons, also plays a considerable part in determining how tightly these outer electrons are held.

Comparative Ionization Energies

Comparing the ionization energies of elements involves understanding their position on the periodic table and their respective nuclear charges. For example, within the given group of elements—silicon (Si), phosphorus (P), sulfur (S), and chlorine (Cl)—the ionization energies follow the expected trend based on their positions.
Silicon, being the furthest left in its period, has the lowest ionization energy because its nuclear charge is the smallest. As you move right to phosphorus, sulfur, and then chlorine, the ionization energy increases.

• Silicon (Si): Lowest ionization energy.
• Phosphorus (P): Higher than Si but lower than S and Cl.
• Sulfur (S): Higher than Si and P.
• Chlorine (Cl): Highest ionization energy.

This comparison aligns with the understanding of periodic trends and highlights the role of effective nuclear charge in determining ionization energies.