Answer
The trend in electron affinities from aluminum to chlorine can be rationalized by considering the increasing effective nuclear charge and decreasing atomic radius as we move across the period. These factors result in a stronger attraction for an additional electron, leading to higher electron affinity values.
Work Step by Step
The trend in electron affinities of the elements from aluminum to chlorine can be rationalized based on their atomic structure and the concept of electron affinity.
Electron affinity is defined as the energy change that occurs when an atom gains an electron to form a negative ion. It is a measure of the atom's ability to attract and hold an additional electron. A higher electron affinity indicates a stronger attraction for an electron.
In this case, we observe the following trend in electron affinities:
Aluminum (-44 kJ/mol) < Silicon (-120 kJ/mol) < Phosphorus (-74 kJ/mol) < Sulfur (-200.4 kJ/mol) < Chlorine (-384.7 kJ/mol)
To rationalize this trend, we need to consider the atomic structure and the factors that influence electron affinity.
1. Effective Nuclear Charge: The effective nuclear charge experienced by an electron in an atom increases as we move across a period from left to right. This is because the number of protons in the nucleus increases, while the shielding effect remains relatively constant. The increased effective nuclear charge leads to a stronger attraction for an additional electron, resulting in higher electron affinity. This explains the general trend from aluminum to chlorine.
2. Atomic Radius: The atomic radius decreases as we move across a period from left to right. This is due to the increasing number of protons in the nucleus, which pulls the electrons closer to the nucleus. As the atomic radius decreases, the outermost electron is closer to the nucleus, resulting in a stronger attraction and higher electron affinity. This factor also contributes to the trend observed.
Based on these factors, we can explain the trend in electron affinities as follows:
Aluminum (Al) has a relatively low electron affinity because it has a larger atomic radius and a lower effective nuclear charge compared to the other elements in the series.
As we move from silicon (Si) to sulfur (S), the effective nuclear charge increases, and the atomic radius decreases. This leads to a gradual increase in electron affinity.
Chlorine (Cl) has the highest electron affinity among the elements listed because it has the highest effective nuclear charge and the smallest atomic radius. The strong attraction between the nucleus and the incoming electron results in the highest electron affinity value.
Phosphorus is out of line. The electron affinity equation for P is:
P(g)+e$\rightarrow$ P$^-$(g)
From the configuration $[Ne]3s^23p^3$ the additional electron leads to the configuration $[Ne]3s^2 3p^4$, which means the additional electron will go into an orbital that already has one electron. The result is that it will be greater repulsion between the paired electrons in P$^-$ , causing the electron affinity of P to be less favorable than predicted based solely on attractions to the nucleus.
