CBSE Class 11 Chemistry Chapter 4: Chemical Bonding and Molecular Structure NCERT Solutions

Solution: A chemical bond is fundamentally an attractive force that binds together the constituent particles, such as atoms or ions, within a chemical species, leading to a more stable entity. The formation of these bonds is driven by the inherent tendency of systems to achieve a state of lower energy and greater stability. It has been observed that elements with fully filled outermost electron shells, like noble gases, are exceptionally stable and unreactive. Conversely, atoms with incomplete outermost shells are less stable and thus more reactive. To attain stability, these atoms tend to combine with each other. This combination can occur through two primary mechanisms: sharing of one or more electrons between atoms, which results in a covalent bond, or the complete transfer of one or more electrons from one atom to another, forming an ionic bond. In both cases, the atoms involved achieve a stable electron configuration, often resembling that of the nearest noble gas (octet rule or duplet rule for hydrogen).

Solution: Lewis dot symbols represent the valence electrons of an atom. The number of dots around the element's symbol corresponds to the number of electrons in its outermost shell.

Magnesium (Mg): Magnesium is in Group 2, so it has 2 valence electrons. Its Lewis dot symbol is: Mg with two dots.

Sodium (Na): Sodium is in Group 1, so it has 1 valence electron. Its Lewis dot symbol is: Na with one dot.

Boron (B): Boron is in Group 13, so it has 3 valence electrons. Its Lewis dot symbol is: B with three dots.

Oxygen (O): Oxygen is in Group 16, so it has 6 valence electrons. Its Lewis dot symbol is: O with six dots.

Nitrogen (N): Nitrogen is in Group 15, so it has 5 valence electrons. Its Lewis dot symbol is: N with five dots.

Bromine (Br): Bromine is in Group 17, so it has 7 valence electrons. Its Lewis dot symbol is: Br with seven dots.

Solution: Lewis symbols for atoms show valence electrons, while for ions, they show valence electrons adjusted for the charge.

1. Sulphur (S) and S²⁻ ion:

   Sulphur (S) is in Group 16 and has 6 valence electrons. Its Lewis dot symbol is S with six dots around it.

   The S²⁻ ion has gained two electrons, so it has 6 + 2 = 8 valence electrons. The Lewis symbol for S²⁻ is represented by S enclosed in square brackets with a charge of 2- outside, and eight dots around the S.

   $$[\; \backslash underset\{\backslash cdot\; \backslash cdot\}\{\backslash stackrel\{\backslash cdot\; \backslash cdot\}\{S\}\}\; \backslash underset\{\backslash cdot\; \backslash cdot\}\{\backslash cdot\; \backslash cdot\}\; ]^\{2-\}$$

2. Aluminium (Al) and Al³⁺ ion:

   Aluminium (Al) is in Group 13 and has 3 valence electrons. Its Lewis dot symbol is Al with three dots around it.

   The Al³⁺ ion has lost its three valence electrons. Therefore, it has 3 - 3 = 0 valence electrons to show. The Lewis symbol for Al³⁺ is represented by Al enclosed in square brackets with a charge of 3+ outside.

   $$[\; Al\; ]^\{3+\}$$

3. Hydrogen (H) and H⁻ ion:

   Hydrogen (H) has 1 valence electron. Its Lewis dot symbol is H with one dot next to it.

   The H⁻ ion has gained one electron, so it has 1 + 1 = 2 valence electrons, achieving the stable duplet configuration like Helium. The Lewis symbol for H⁻ is represented by H enclosed in square brackets with a charge of 1- outside, and two dots around the H.

   $$[\; \backslash dot\{H\}\; ]^\{-\}$$

Solution: Lewis structures represent the bonding and lone pair electrons in a molecule or ion.

1. Hydrogen Sulphide (H₂S):

   Sulphur (S) has 6 valence electrons, and each Hydrogen (H) has 1 valence electron. Total valence electrons = 6 + 2(1) = 8. S is the central atom. Two single bonds are formed between S and each H. The remaining 4 electrons on S form two lone pairs. $$H\; -\; \backslash underset\{\backslash cdot\; \backslash cdot\}\{\backslash stackrel\{\backslash cdot\; \backslash cdot\}\{S\}\}\; -\; H$$

2. Silicon Tetrachloride (SiCl₄):

   Silicon (Si) has 4 valence electrons, and each Chlorine (Cl) has 7 valence electrons. Total valence electrons = 4 + 4(7) = 32. Si is the central atom. Four single bonds are formed between Si and each Cl. Each Cl atom has 3 lone pairs to complete its octet. (Each Cl has 3 lone pairs not shown for simplicity)

3. Beryllium Fluoride (BeF₂):

   Beryllium (Be) has 2 valence electrons, and each Fluorine (F) has 7 valence electrons. Total valence electrons = 2 + 2(7) = 16. Be is the central atom. Two single bonds are formed between Be and each F. Each F atom has 3 lone pairs to complete its octet. Be has only 4 electrons around it, which is an exception to the octet rule. $$F\; -\; Be\; -\; F$$ (Each F has 3 lone pairs not shown for simplicity)

4. Carbonate Ion (CO₃²⁻):

   Carbon (C) has 4 valence electrons, Oxygen (O) has 6 valence electrons, and the charge is 2-. Total valence electrons = 4 + 3(6) + 2 = 24. C is the central atom. One C=O double bond, one C-O single bond, and one C-O single bond are formed. The single-bonded O atoms have 3 lone pairs each, and the double-bonded O has 2 lone pairs. The C atom has an octet. The overall structure has a 2- charge. (Lone pairs on oxygen atoms are not shown for simplicity, but are present to complete their octets.)

5. Formic Acid (HCOOH):

   Carbon (C) has 4 valence electrons, Hydrogen (H) has 1 valence electron, and Oxygen (O) has 6 valence electrons. Total valence electrons = 1 + 1 + 4 + 2(6) = 18. C is bonded to one O with a double bond, to another O with a single bond, and this second O is bonded to an H. The first O has 2 lone pairs. The second O has 2 lone pairs. The H attached to C has no lone pairs. The H attached to O has no lone pairs. $$\backslash begin\{bmatrix\}\; H\; \backslash \backslash \; |\; \backslash \backslash \; C=O\; \backslash \backslash \; |\; \backslash \backslash \; O-H\; \backslash end\{bmatrix\}$$ (Lone pairs on oxygen atoms are not shown for simplicity.)