Lewis Structure Worksheet | Lecture notes Geometry | Docsity - Free Printable
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Step-by-step solution for: Lewis Structure Worksheet | Lecture notes Geometry | Docsity
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Show Answer Key & Explanations
Step-by-step solution for: Lewis Structure Worksheet | Lecture notes Geometry | Docsity
Here are the Lewis structures for the requested problems. I have calculated the total valence electrons for each, arranged the atoms to satisfy the octet rule (8 electrons around most atoms, 2 for Hydrogen), and minimized formal charges where possible.
1. CH₄ (Methane)
* Total Valence Electrons: $4 + 4(1) = 8$
* Structure: Carbon is in the center with 4 single bonds to Hydrogen atoms. No lone pairs on Carbon.
2. BF₃ (Boron Trifluoride)
* Total Valence Electrons: $3 + 3(7) = 24$
* Structure: Boron is in the center bonded to 3 Fluorine atoms. Each Fluorine has 3 lone pairs. Boron is an exception to the octet rule and is stable with only 6 electrons here.
3. HCN (Hydrogen Cyanide)
* Total Valence Electrons: $1 + 4 + 5 = 10$
* Structure: H — C ≡ N :
* Single bond between H and C.
* Triple bond between C and N.
* One lone pair on Nitrogen.
4. NH₃ (Ammonia)
* Total Valence Electrons: $5 + 3(1) = 8$
* Structure: Nitrogen in the center bonded to 3 Hydrogens. Nitrogen has 1 lone pair.
5. H₂O (Water)
* Total Valence Electrons: $2(1) + 6 = 8$
* Structure: Oxygen in the center bonded to 2 Hydrogens. Oxygen has 2 lone pairs.
6. ClO₃ (Chlorate Ion)
* Total Valence Electrons: $7 + 3(6) + 1 = 26$
* Preferred Structure (Minimizing Formal Charge): Chlorine in the center. Two double bonds to Oxygen, one single bond to Oxygen. The single-bonded Oxygen has a -1 charge. Chlorine has 1 lone pair.
* *(Note: A structure with all single bonds is also valid but has higher formal charges).*
7. H₃O⁺ (Hydronium Ion)
* Total Valence Electrons: $3(1) + 6 - 1 = 8$
* Structure: Oxygen in the center bonded to 3 Hydrogens. Oxygen has 1 lone pair and a +1 formal charge.
8. NH₂⁻ (Amide Ion)
* Total Valence Electrons: $5 + 2(1) + 1 = 8$
* Structure: Nitrogen in the center bonded to 2 Hydrogens. Nitrogen has 2 lone pairs and a -1 formal charge.
9. NH₄⁺ (Ammonium Ion)
* Total Valence Electrons: $5 + 4(1) - 1 = 8$
* Structure: Nitrogen in the center bonded to 4 Hydrogens. No lone pairs. Nitrogen has a +1 formal charge.
10. C₂H₄ (Ethene)
* Total Valence Electrons: $2(4) + 4(1) = 12$
* Structure: Two Carbons connected by a double bond ($C=C$). Each Carbon is bonded to 2 Hydrogens.
11. C₂H₂ (Ethyne/Acetylene)
* Total Valence Electrons: $2(4) + 2(1) = 10$
* Structure: Two Carbons connected by a triple bond ($C \equiv C$). Each Carbon is bonded to 1 Hydrogen.
12. N₃⁻ (Azide Ion)
* Total Valence Electrons: $3(5) + 1 = 16$
* Resonance Structures:
1. $[ \ddot{N} = N = \ddot{N} ]^-$ (Central N has no lone pairs; outer N's have 2 lone pairs each). (Most Preferred)
2. $[ :N \equiv N - \ddot{N}: ]^{2-}$ ... actually $[ :N \equiv N - \ddot{N}: ]^-$ (Left N has 1 lone pair, Right N has 3 lone pairs).
3. Mirror image of #2.
13. NOF (Nitrosyl Fluoride)
* Total Valence Electrons: $5 + 6 + 7 = 18$
* Structure: Nitrogen is central. Double bond to Oxygen ($N=O$), single bond to Fluorine ($N-F$). Nitrogen has 1 lone pair. Oxygen has 2 lone pairs. Fluorine has 3 lone pairs.
14. ClF₂⁺
* Total Valence Electrons: $7 + 2(7) - 1 = 20$
* Structure: Chlorine in the center bonded to 2 Fluorines. Chlorine has 2 lone pairs. Each Fluorine has 3 lone pairs. Chlorine has a +1 formal charge.
15. FNO₂ (Nitryl Fluoride)
* Total Valence Electrons: $7 + 5 + 2(6) = 24$
* Structure: Nitrogen is central. Single bond to Fluorine. Double bond to one Oxygen, single bond to the other Oxygen (which has a -1 charge). Nitrogen has a +1 charge.
* *Resonance:* The double bond can switch between the two oxygens.
16. XeF₅⁺
* Total Valence Electrons: $8 + 5(7) - 1 = 42$
* Structure: Xenon in the center bonded to 5 Fluorines. Xenon has 1 lone pair. Each Fluorine has 3 lone pairs. Xenon has a +1 formal charge.
17. BrF₅
* Total Valence Electrons: $7 + 5(7) = 42$
* Structure: Bromine in the center bonded to 5 Fluorines. Bromine has 1 lone pair. Each Fluorine has 3 lone pairs.
18. IF₄⁻
* Total Valence Electrons: $7 + 4(7) + 1 = 36$
* Structure: Iodine in the center bonded to 4 Fluorines. Iodine has 2 lone pairs. Each Fluorine has 3 lone pairs. Iodine has a -1 formal charge.
19. SF₆
* Total Valence Electrons: $6 + 6(7) = 48$
* Structure: Sulfur in the center bonded to 6 Fluorines. No lone pairs on Sulfur. Each Fluorine has 3 lone pairs.
20. N₂O₄
* Total Valence Electrons: $2(5) + 4(6) = 34$
* Structure: Two Nitrogens bonded together ($N-N$). Each Nitrogen is double-bonded to one Oxygen and single-bonded to another Oxygen (with a negative charge). The Nitrogens carry a positive charge.
* *Resonance:* The double/single bonds on the oxygens can swap places.
21. C₂O₄²⁻ (Oxalate Ion)
* Total Valence Electrons: $2(4) + 4(6) + 2 = 34$
* Structure: Two Carbons bonded together ($C-C$). Each Carbon is double-bonded to one Oxygen and single-bonded to another Oxygen (with a -1 charge).
* *Resonance:* The double bonds can flip positions on both sides.
22. H₂PO₄⁻ (Dihydrogen Phosphate)
* Total Valence Electrons: $2(1) + 5 + 4(6) + 1 = 32$
* Structure: Phosphorus in the center. Double bond to one Oxygen. Single bonds to two OH groups. Single bond to one Oxygen with a -1 charge.
23. XeO₄
* Total Valence Electrons: $8 + 4(6) = 32$
* Preferred Structure (Minimizing Formal Charge): Xenon in the center with four double bonds to the four Oxygens ($Xe=O$). Xenon has 0 lone pairs. Each Oxygen has 2 lone pairs.
24. PO₃⁻
* Total Valence Electrons: $5 + 3(6) + 1 = 24$
* Structure: Phosphorus in the center. Double bond to one Oxygen. Single bonds to two Oxygens (each with a -1 charge).
* *Note:* This ion is less common than phosphate ($PO_4^{3-}$), but structurally similar to nitrate ($NO_3^-$) but with P expanding its octet.
25. NO (Nitric Oxide)
* Total Valence Electrons: $5 + 6 = 11$
* Structure: $: \dot{N} = \ddot{O} :$
* Double bond between N and O.
* Nitrogen has 1 lone pair and 1 unpaired electron (radical).
* Oxygen has 2 lone pairs.
* *(This is an odd-electron molecule, so it cannot satisfy the octet rule for both atoms).*
1. CH₄ (Methane)
* Total Valence Electrons: $4 + 4(1) = 8$
* Structure: Carbon is in the center with 4 single bonds to Hydrogen atoms. No lone pairs on Carbon.
2. BF₃ (Boron Trifluoride)
* Total Valence Electrons: $3 + 3(7) = 24$
* Structure: Boron is in the center bonded to 3 Fluorine atoms. Each Fluorine has 3 lone pairs. Boron is an exception to the octet rule and is stable with only 6 electrons here.
3. HCN (Hydrogen Cyanide)
* Total Valence Electrons: $1 + 4 + 5 = 10$
* Structure: H — C ≡ N :
* Single bond between H and C.
* Triple bond between C and N.
* One lone pair on Nitrogen.
4. NH₃ (Ammonia)
* Total Valence Electrons: $5 + 3(1) = 8$
* Structure: Nitrogen in the center bonded to 3 Hydrogens. Nitrogen has 1 lone pair.
5. H₂O (Water)
* Total Valence Electrons: $2(1) + 6 = 8$
* Structure: Oxygen in the center bonded to 2 Hydrogens. Oxygen has 2 lone pairs.
6. ClO₃ (Chlorate Ion)
* Total Valence Electrons: $7 + 3(6) + 1 = 26$
* Preferred Structure (Minimizing Formal Charge): Chlorine in the center. Two double bonds to Oxygen, one single bond to Oxygen. The single-bonded Oxygen has a -1 charge. Chlorine has 1 lone pair.
* *(Note: A structure with all single bonds is also valid but has higher formal charges).*
7. H₃O⁺ (Hydronium Ion)
* Total Valence Electrons: $3(1) + 6 - 1 = 8$
* Structure: Oxygen in the center bonded to 3 Hydrogens. Oxygen has 1 lone pair and a +1 formal charge.
8. NH₂⁻ (Amide Ion)
* Total Valence Electrons: $5 + 2(1) + 1 = 8$
* Structure: Nitrogen in the center bonded to 2 Hydrogens. Nitrogen has 2 lone pairs and a -1 formal charge.
9. NH₄⁺ (Ammonium Ion)
* Total Valence Electrons: $5 + 4(1) - 1 = 8$
* Structure: Nitrogen in the center bonded to 4 Hydrogens. No lone pairs. Nitrogen has a +1 formal charge.
10. C₂H₄ (Ethene)
* Total Valence Electrons: $2(4) + 4(1) = 12$
* Structure: Two Carbons connected by a double bond ($C=C$). Each Carbon is bonded to 2 Hydrogens.
11. C₂H₂ (Ethyne/Acetylene)
* Total Valence Electrons: $2(4) + 2(1) = 10$
* Structure: Two Carbons connected by a triple bond ($C \equiv C$). Each Carbon is bonded to 1 Hydrogen.
12. N₃⁻ (Azide Ion)
* Total Valence Electrons: $3(5) + 1 = 16$
* Resonance Structures:
1. $[ \ddot{N} = N = \ddot{N} ]^-$ (Central N has no lone pairs; outer N's have 2 lone pairs each). (Most Preferred)
2. $[ :N \equiv N - \ddot{N}: ]^{2-}$ ... actually $[ :N \equiv N - \ddot{N}: ]^-$ (Left N has 1 lone pair, Right N has 3 lone pairs).
3. Mirror image of #2.
13. NOF (Nitrosyl Fluoride)
* Total Valence Electrons: $5 + 6 + 7 = 18$
* Structure: Nitrogen is central. Double bond to Oxygen ($N=O$), single bond to Fluorine ($N-F$). Nitrogen has 1 lone pair. Oxygen has 2 lone pairs. Fluorine has 3 lone pairs.
14. ClF₂⁺
* Total Valence Electrons: $7 + 2(7) - 1 = 20$
* Structure: Chlorine in the center bonded to 2 Fluorines. Chlorine has 2 lone pairs. Each Fluorine has 3 lone pairs. Chlorine has a +1 formal charge.
15. FNO₂ (Nitryl Fluoride)
* Total Valence Electrons: $7 + 5 + 2(6) = 24$
* Structure: Nitrogen is central. Single bond to Fluorine. Double bond to one Oxygen, single bond to the other Oxygen (which has a -1 charge). Nitrogen has a +1 charge.
* *Resonance:* The double bond can switch between the two oxygens.
16. XeF₅⁺
* Total Valence Electrons: $8 + 5(7) - 1 = 42$
* Structure: Xenon in the center bonded to 5 Fluorines. Xenon has 1 lone pair. Each Fluorine has 3 lone pairs. Xenon has a +1 formal charge.
17. BrF₅
* Total Valence Electrons: $7 + 5(7) = 42$
* Structure: Bromine in the center bonded to 5 Fluorines. Bromine has 1 lone pair. Each Fluorine has 3 lone pairs.
18. IF₄⁻
* Total Valence Electrons: $7 + 4(7) + 1 = 36$
* Structure: Iodine in the center bonded to 4 Fluorines. Iodine has 2 lone pairs. Each Fluorine has 3 lone pairs. Iodine has a -1 formal charge.
19. SF₆
* Total Valence Electrons: $6 + 6(7) = 48$
* Structure: Sulfur in the center bonded to 6 Fluorines. No lone pairs on Sulfur. Each Fluorine has 3 lone pairs.
20. N₂O₄
* Total Valence Electrons: $2(5) + 4(6) = 34$
* Structure: Two Nitrogens bonded together ($N-N$). Each Nitrogen is double-bonded to one Oxygen and single-bonded to another Oxygen (with a negative charge). The Nitrogens carry a positive charge.
* *Resonance:* The double/single bonds on the oxygens can swap places.
21. C₂O₄²⁻ (Oxalate Ion)
* Total Valence Electrons: $2(4) + 4(6) + 2 = 34$
* Structure: Two Carbons bonded together ($C-C$). Each Carbon is double-bonded to one Oxygen and single-bonded to another Oxygen (with a -1 charge).
* *Resonance:* The double bonds can flip positions on both sides.
22. H₂PO₄⁻ (Dihydrogen Phosphate)
* Total Valence Electrons: $2(1) + 5 + 4(6) + 1 = 32$
* Structure: Phosphorus in the center. Double bond to one Oxygen. Single bonds to two OH groups. Single bond to one Oxygen with a -1 charge.
23. XeO₄
* Total Valence Electrons: $8 + 4(6) = 32$
* Preferred Structure (Minimizing Formal Charge): Xenon in the center with four double bonds to the four Oxygens ($Xe=O$). Xenon has 0 lone pairs. Each Oxygen has 2 lone pairs.
24. PO₃⁻
* Total Valence Electrons: $5 + 3(6) + 1 = 24$
* Structure: Phosphorus in the center. Double bond to one Oxygen. Single bonds to two Oxygens (each with a -1 charge).
* *Note:* This ion is less common than phosphate ($PO_4^{3-}$), but structurally similar to nitrate ($NO_3^-$) but with P expanding its octet.
25. NO (Nitric Oxide)
* Total Valence Electrons: $5 + 6 = 11$
* Structure: $: \dot{N} = \ddot{O} :$
* Double bond between N and O.
* Nitrogen has 1 lone pair and 1 unpaired electron (radical).
* Oxygen has 2 lone pairs.
* *(This is an odd-electron molecule, so it cannot satisfy the octet rule for both atoms).*
Parent Tip: Review the logic above to help your child master the concept of lewis dot worksheet.
