Problem 31 - Chemistry An Atoms First Approach

Question

Without using Fig. \(3-4,\) predict the order of increasing electronegativity in each of the following groups of elements. a. \(C, N, O\) b. \(\mathbf{s}, \mathbf{S e}, \mathbf{C l}\) \(\mathbf{c}_{*} \mathrm{Si}, \mathrm{Ge}, \mathrm{Sn}\) d. \(\mathrm{TI}, \mathrm{S}, \mathrm{Ge}\)

Step-by-Step Solution

Answer

The order of increasing electronegativity for the given groups of elements are: a. C < N < O b. S < Se < Cl c. Sn < Ge < Si d. Tl < Ge < S

Step 1: a. C, N, O

In this group, all three elements C, N, and O belong to the same period. We will use the periodic trend stating that electronegativity increases from left to right in a period. The order in the periodic table from left to right is C, N, and O. Therefore, the order of increasing electronegativity is: C < N < O

Step 2: b. S, Se, Cl

In this group, all three elements S, Se, and Cl belong to the same period. We will use the periodic trend stating that electronegativity increases from left to right in a period. The order in the periodic table from left to right is S, Se, and Cl. Therefore, the order of increasing electronegativity is: S < Se < Cl

Step 3: c. Si, Ge, Sn

In this group, the elements Si, Ge, and Sn belong to the same group in the periodic table. We will use the periodic trend stating that electronegativity decreases as we move down a group. The order in the periodic table from top to bottom is Si, Ge, and Sn. Therefore, the order of increasing electronegativity is: Sn < Ge < Si

Step 4: d. Tl, S, Ge

In this group, the elements are not in the same period or group. Tl belongs to the last period in Group 13, S belongs to Period 3 in Group 16, and Ge belongs to Period 4 in Group 14. Since electronegativity generally decreases moving down a group and increases across a period, we can predict: - Tl has lower electronegativity than Ge, since Ge is on the right and in a higher period. - S has higher electronegativity than Ge, since S is on the right of the periodic table. - S has higher electronegativity than Tl, since S is in a much higher period and more to the right. Considering these trends, we can predict the order of increasing electronegativity as: Tl < Ge < S

Key Concepts

Periodic Trends

In the context of the periodic table, electronegativity is not a random property—it follows specific patterns known as periodic trends. One key trend is that electronegativity increases as we move across a period from left to right. This increase is due to the addition of protons in the nucleus which leads to a stronger attraction of the valence electrons, making atoms more likely to attract electrons in a bond.

Another trend is seen when moving down a group in the periodic table: electronegativity decreases. As the atomic number increases, so does the distance between the valence electrons and the nucleus, thanks to the addition of more electron shells. This distance weakens the nucleus' hold on the valence electrons, reducing the atom's ability to attract electrons towards itself. These two trends are instrumental in predicting the relative electronegativity of elements without the use of a chart or figure, as demonstrated in the exercise.

Periodic Table Groups

The periodic table is organized into columns called 'groups,' each representing elements with similar chemical and physical properties. This organization helps in understanding the properties of elements. For instance, elements in the same group have the same number of valence electrons, which significantly influences their chemical behavior, such as bonding and reactivity.

In the above exercise, when comparing elements such as Si, Ge, and Sn, we are looking at elements from Group 14. These elements have four valence electrons and tend to show a decrease in electronegativity as we move down the group. Understanding the commonalities within these groups and how properties change systematically helps to predict behaviors of the elements without needing to memorize each elements' specific properties.

Atomic Properties

Atomic properties such as atomic radius, ionization energy, and electronegativity are not standalone characteristics—they are interrelated. As the atomic radius increases, the electronegativity usually decreases because the nucleus is further from the valence electrons and exerts less pull on them. Conversely, a smaller atomic radius typically indicates higher electronegativity. Ionization energy also correlates with electronegativity, as elements with higher electronegativity generally have higher ionization energies. This is because the tight hold on their valence electrons requires more energy to remove an electron.

The atomic structure of an element, thus, underpins its electronegativity. By understanding these atomic properties and their interconnections, one can make educated predictions about an element's electronegativity and other chemical behaviors, enhancing the understanding of the step-by-step solutions provided in the textbook exercise.