Question

Without using Fig. \(3-4,\) predict the order of increasing electronegativity in each of the following groups of elements. a. \(\mathrm{Na}, \mathrm{K}, \mathrm{Rb}\) \(\mathbf{b} . \mathbf{B}, \mathbf{O}, \mathbf{G a}\) c. \(F, C\), \(B r\) d. \(s, O, F\)

Short answer

The orders of increasing electronegativity for each group are: a. K < Rb < Na b. Ga < B < O c. Br < C < F d. S < O < F

Step-by-step solution

Group a: Na, K, Rb

All three elements (Na, K, and Rb) are in the same group of the periodic table (Group 1: alkali metals). They are ordered as Na is above K, and K is above Rb. As electronegativity decreases down the group, the order of increasing electronegativity will be: K < Rb < Na.

Group b: B, O, Ga

B and Ga are in Group 13 (boron group), while O is in Group 16 (chalcogens). Since O is further to the right compared to B and Ga, it will have the highest electronegativity. Between B and Ga, B is above Ga, so B will have higher electronegativity. The order of increasing electronegativity will be: Ga < B < O.

Group c: F, C, Br

F and Br are in Group 17 (halogens), while C is in Group 14 (carbon group). F is the top element in its group, and Br is below it. C is in between these two elements horizontally. The electronegativity decreases down the group, and increases from left to right across a period. Thus, the order of increasing electronegativity will be: Br < C < F.

Group d: S, O, F

All three elements (S, O, and F) are in the same period (Period 2). S is in Group 16 (chalcogens), O is in Group 15 (pnictogens), and F is in Group 17 (halogens). Electronegativity increases from left to right across a period. Thus, the order of increasing electronegativity will be: S < O < F.

Key concepts

Periodic Table Trends

Understanding the periodic table is fundamental to grasping various chemical properties, including electronegativity. Electronegativity measures an atom's tendency to attract and bond with electrons. This value generally increases from left to right across a period, as the number of charges on the nucleus increases, which attracts more electrons towards itself. Conversely, electronegativity decreases down a group as the number of electron shells increases, creating more distance between the nucleus and the valence electrons, thus, weakening their attraction. These trends are pivotal in predicting the behavior of elements in different parts of the table, like alkali metals, halogens, and others based on their respective position.

Alkali Metals

The alkali metals, which include elements such as sodium (Na), potassium (K), and rubidium (Rb), are located in Group 1 of the periodic table. These metals are characterized by their single valence electron, which they readily lose to form positive ions.
Alkali metals have low electronegativity values because of their larger atomic radius and the fact that their valence electron is further away from the nucleus. As one moves down this group, from Na to Rb for instance, the electronegativity decreases. This explains why in increasing order, K (<) Rb (<) Na, potassium has a higher electronegativity than rubidium but a lower electronegativity than sodium.

Halogens

Halogens are located in Group 17 of the periodic table, which includes elements like fluorine (F), chlorine (Cl), and bromine (Br). They are highly reactive nonmetals with seven valence electrons, making them one electron shy of a full octet. This high electron affinity is why halogens are among the most electronegative elements.
Fluorine is the most electronegative element in the group because it is at the top of the group and has fewer electron shells, strengthening the nuclear attraction to valence electrons. Bromine, being lower on the group, has a lower electronegativity due to additional inner electron shells that shield the valence electron from the nucleus.

Atomic Structure

Atomic structure is central to understanding why electronegativity trends exist on the periodic table. An atom consists of a nucleus composed of protons and neutrons, surrounded by orbiting electrons. Electronegativity is influenced by two main factors: the number of protons and the distance between the nucleus and the valence electrons.
Atoms with a higher number of protons exert a stronger attractive force on electrons, thus increasing the atom's electronegativity. Moreover, as the number of electron shells increases, this creates more distance between the nucleus and valence electrons, resulting in reduced electronegativity. Elements on the same period have the same number of shells but differ in nuclear charge, leading to varying degrees of electronegativity.