Problem 92 - A Complete Resource Book in Chemistry for JEE Main

Question

The electronegativity of the following elements increases in the order: (a) Si, P, C, N (b) \(\mathrm{N}, \mathrm{Si}, \mathrm{C}, \mathrm{P}\) (c) \(\mathrm{P}, \mathrm{Si}, \mathrm{N}, \mathrm{C}\) (d) \(\mathrm{C}, \mathrm{N}, \mathrm{Si}, \mathrm{P}\)

Step-by-Step Solution

Answer

The correct order is (a) Si, P, C, N.

Step 1: Understanding Electronegativity

Electronegativity is the tendency of an atom to attract electrons towards itself in a chemical bond. In general, electronegativity increases across a period from left to right and decreases down a group in the periodic table.

Step 2: Identify Periods and Groups

Locate the elements on the periodic table: Silicon (Si) and Phosphorus (P) are in period 3, Carbon (C) and Nitrogen (N) are in period 2. Group-wise, Si is in group 14, P is in group 15, C is in group 14, and N is in group 15.

Step 3: Compare Electronegativity Across Periods

Within the same period (Si and P in period 3), electronegativity increases from left to right. Thus, P is more electronegative than Si. Similarly, in period 2, N is more electronegative than C.

Step 4: Compare Electronegativity Down Groups

Comparing elements across the same group, electronegativity decreases down a group. Thus, between N (group 15) and P (group 15), N is more electronegative, and C is more electronegative than Si (group 14).

Step 5: Arrange Based on Electronegativity

Considering the trends, the correct order of increasing electronegativity is Si, P, C, N. Si (least electronegative), followed by P, then C, and lastly N (most electronegative).

Key Concepts

Periodic Table

The periodic table is an organized display of all known chemical elements, arranged by increasing atomic number. It is designed in a way that reveals the recurring 'periodic' trends in the properties of the elements.
As we look from left to right across the periodic table, elements show a progression from metals on the left to nonmetals on the right. This arrangement also allows us to observe shifts in properties like electronegativity.
This table is divided into rows, called "periods," and columns, known as "groups". Understanding these arrangements helps in predicting and explaining the chemical behavior of elements, including their electronegativity.

Chemical Bonding

Chemical bonding refers to the forces that hold atoms together in compounds. There are primarily three types of bonds: ionic, covalent, and metallic. Electronegativity plays a crucial role in determining the type and strength of these bonds.

• Ionic bonds: Formed when electrons are transferred from one atom to another, typically between metals and nonmetals with a large difference in electronegativity.
• Covalent bonds: Involve the sharing of electrons between atoms, usually nonmetals with similar electronegativities.
• Metallic bonds: Consist of positive metal ions surrounded by a sea of electrons, characteristic of metal atoms.

The concept of electronegativity is essential in understanding how and why atoms attract electrons differently and form bonds as they do. Stronger electronegative atoms "pull" electrons towards themselves in a covalent bond, resulting in partial charges.

Element Groups

Element groups are columns of the periodic table and are numbered 1 to 18. Elements within the same group tend to exhibit similar chemical and physical properties because they have the same number of valence electrons.
Group 14, for example, includes elements like carbon (C) and silicon (Si). They have four valence electrons, and this commonality affects their bonding patterns and overall chemical behavior.
On the other hand, Group 15 includes nitrogen (N) and phosphorus (P), which have five valence electrons. This difference in valence electrons is key to understanding the differences in their electronegativity. Higher valence electrons generally mean higher electronegativity because these atoms are more inclined to complete their valence shell by attracting electrons.

Atomic Properties

Atomic properties, such as size, ionization energy, and electronegativity, are fundamental in predicting and explaining the chemical behavior of elements.

• Atomic Size: Generally decreases across a period due to the increase in nuclear charge pulling the electron cloud closer to the nucleus, and increases down a group because of the addition of electron shells.
• Ionization Energy: The energy required to remove an electron from an atom, increases across a period and decreases down a group.
• Electronegativity: The tendency of an atom to attract electrons in a chemical bond, increases across a period from left to right and decreases down a group.

Understanding these atomic properties helps predict how elements will behave in reactions, including which will attract electrons more strongly in a compound. For example, nitrogen is more electronegative than phosphorus because it is higher up in group 15, aligning with these trends.