Covalent Bond Lewis Structures - UNSOLVED PRACTICE SET

Q7. Define covalent bond. What are the conditions favourable for covalent bond formation? 

Q8. Draw the Lewis structures for:

(a) H₂O

(b) NH₃

Q9. What is a coordinate bond? How does it differ from a normal covalent bond? Give one example.

Q10. Why does carbon form four covalent bonds but not ionic bonds? 

Q11. Your elder sister is drawing rangoli with dots and connecting them with lines to make beautiful patterns. She explains that Lewis structures are similar — dots represent electrons and lines represent bonds. If she draws a molecule with 8 dots around one atom and 2 dots shared between two atoms, what type of bond is represented, and how many lone pairs does the first atom have?

Q12. Draw the Lewis structure of HNO₃ and identify all the bonds (single, double, coordinate) present in it. 

Q13. Explain the formation of covalent bonds with suitable examples. Describe the different types of covalent bonds: single, double, and triple. Draw Lewis structures for CH₄, C₂H₄, and C₂H₂, showing all bonds and lone pairs. How does the bond length and bond energy vary with bond multiplicity?

Q14. What are Lewis structures? Explain the steps involved in drawing a Lewis structure for any molecule or ion. Draw Lewis structures for:

(a) SO₄²⁻

(b) CO₃²⁻

(c) H₂SO₄

Identify any coordinate bonds present in these structures.

Q15. During a chemistry club activity, students are given a set of molecular formulas and asked to draw Lewis structures.

(a) Student A draws O₂ with a single bond and two lone pairs on each oxygen. Student B draws it with a double bond and two lone pairs on each oxygen. Which is correct? Explain using the octet rule.

(b) A student draws CO with a structure showing C=O with two lone pairs on oxygen and one lone pair on carbon. Another student suggests a triple bond with one lone pair on each atom. Which structure is more accurate, and what additional concept explains this?

(c) Draw the Lewis structure of ozone (O₃). Why does the central oxygen atom have a positive formal charge in one resonance structure?

(d) Why can't the Lewis structure of BF₃ satisfy the octet rule for boron? What does this tell us about the limitations of the octet rule?

Q16. For the following molecules, calculate the total number of valence electrons and draw their Lewis structures:

(a) CH₂Cl₂

(b) HCN

(c) C₂H₆O (two possible structures)

For each structure, identify:

Total valence electrons

Number of bonding pairs

Number of lone pairs on each atom

Q17. The cyanide ion (CN⁻) has a total of 10 valence electrons.

(a) Draw the Lewis structure of CN⁻.

(b) Identify the type of bond (single, double, or triple) between carbon and nitrogen.

(c) Calculate the formal charge on each atom in your structure.

(d) The isoelectronic molecule N₂ has a similar bond. Compare the bond order and bond length of CN⁻ with N₂.

Q18. Consider the following species: CO, CO₂, and CO₃²⁻.

(a) Draw the Lewis structure of each species.

(b) For CO, calculate the formal charges on carbon and oxygen for both possible Lewis structures (C≡O and C=O). Which structure is more stable?

(c) In CO₃²⁻, draw all possible resonance structures and calculate the average bond order of each C-O bond.

(d) Compare the C-O bond lengths in CO, CO₂, and CO₃²⁻. Arrange them in increasing order and explain your reasoning.