Question

Without using Fig. 8.3, predict the order of increasing electronegativity in each of the following groups of elements. a. \(\mathrm{C}, \mathrm{N}, \mathrm{O} \quad\) c. \(\mathrm{Si}, \mathrm{Ge}, \mathrm{Sn}\) b. \(\mathrm{S}, \mathrm{Se}, \mathrm{Cl} \quad\) d. \(\mathrm{TI}, \mathrm{S}, \mathrm{Ge}\)

Short answer

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

Step-by-step solution

Determine Groups and Periods

Check the periodic table to find the group and period of each element. a. C, N, O -> second period, Group 14, 15, and 16 b. S, Se, Cl -> third period, Group 16, 16, and 17 c. Si, Ge, Sn -> Group 14, with Si in period 3, Ge in period 4, and Sn in period 5 d. Tl, S, Ge -> Tl is in Group 13, period 6; S is in Group 16, period 3; and Ge is in Group 14, period 4

Apply Electronegativity Trends

Use the periodic table trends to determine the order of increasing electronegativity for each group of elements. a. C, N, O -> Electronegativity increases going from left to right across a period (Group 14 to 16), so the order is C < N < O. b. S, Se, Cl -> Electronegativity decreases going down a group (S and Se are in the same group), and it increases from left to right (Group 16 to 17). So, the order is Se < S < Cl. c. Si, Ge, Sn -> Electronegativity decreases going down a group, so the order is Sn < Ge < Si. d. Tl, S, Ge -> Electronegativity decreases going down a group (Tl and Ge in different periods) and increases from left to right (Group 13 to 14). So, the order is Tl < Ge < S.

Final Answer

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

Key concepts

Periodic Table Trends

Electronegativity is a key concept in understanding chemical bonding and reactions. On the periodic table, electronegativity shows a predictable trend. As you move across a period from left to right, electronegativity tends to increase. This happens because the elements have more protons in the nucleus, which pulls the electrons closer and increases the atom's attraction for electrons.

On the other hand, as you move down a group in the periodic table, electronegativity decreases. The additional electron shells shield the nucleus, making it less effective in attracting additional electrons. Thus, hydrogen at the top of its group has higher electronegativity than heavier elements in the same group.

These trends help in predicting how elements will interact during chemical reactions. For example:

• Fluorine, a halogen, is the most electronegative element.
• Caesium, an alkali metal, has low electronegativity.

Remembering these trends can make solving problems about element reactivity much easier.

Chemical Periodicity

Chemical periodicity refers to the recurring trends that you observe in the periodic table as you move through periods and down groups. A significant aspect of chemical periodicity is the gradual progression in properties of elements as you go across a period or down a group.

For instance, the periodic table is arranged in such a way that each column, or group, contains elements with similar chemical and physical properties. These properties include electronegativity, ionization energy, and atomic radius. As previously discussed, electronegativity generally increases across a period because of added protons in the nucleus, making the atom more likely to attract electrons.

Another important periodic trend is ionization energy, which also increases across a period and decreases down a group. All these principles of chemical periodicity play a crucial role in predicting how different elements interact with each other in chemical reactions.

Group and Period Identification

Group and period identification on the periodic table is fundamental to understanding properties and behaviors of elements like electronegativity. Each row in the periodic table is called a period, and each column is called a group.

Elements in the same period have the same number of electron shells. As a result, they have different chemical properties but exhibit gradual changes in reactivity and electronegativity. For example, elements like carbon (C), nitrogen (N), and oxygen (O) are in the second period and show increasing electronegativity.

Groups, however, are groups of elements with similar chemical properties due to having the same number of electrons in their outer shell. Hence, sulfur (S) and selenium (Se) are in Group 16 and share similar reactivity patterns. Understanding these groupings helps predict how one element might behave when combined with others. This identification is crucial for organizing chemical reactions and predicting outcomes.