Lewis Structure Practice Questions with Answers

Concept Explanation

A Lewis structure is a simplified graphic representation of the distribution of valence electrons around atoms within a molecule or polyatomic ion. These diagrams, also known as Lewis dot structures or electron dot structures, help chemists visualize how atoms bond to achieve stability. According to the octet rule, most atoms are most stable when they have eight electrons in their outer shell, mimicking the electron configuration of a noble gas. Hydrogen is a notable exception, following the duet rule with only two electrons.

To draw an accurate Lewis structure, you must first calculate the total number of valence electrons available by summing the valence electrons for each atom in the chemical formula. For ions, you add electrons for negative charges and subtract them for positive charges. The central atom is typically the least electronegative element (excluding hydrogen). Once the skeletal structure is drawn with single bonds, the remaining electrons are distributed as lone pairs to satisfy the octet rule for terminal atoms, then the central atom. If the central atom lacks an octet, multiple bonds (double or triple) are formed by shifting lone pairs from surrounding atoms. Understanding these structures is a prerequisite for more advanced topics, such as calculating formal charges or using redox reaction practice questions to track electron shifts.

Solved Examples

1. Draw the Lewis structure for Water (H₂O).

   1. Count valence electrons: H (1x2) + O (6) = 8 valence electrons.

   2. Oxygen is the central atom; Hydrogen atoms are terminal.

   3. Place two electrons (a single bond) between each H and O. (4 electrons used).

   4. Distribute remaining 4 electrons as lone pairs on Oxygen.

   5. Check octets: H has 2 (duet), O has 8 (octet). The structure is complete.

2. Draw the Lewis structure for Carbon Dioxide (CO₂).

   1. Count valence electrons: C (4) + O (6x2) = 16 valence electrons.

   2. Carbon is the central atom.

   3. Place single bonds between C and each O. (4 electrons used, 12 remaining).

   4. Place three lone pairs on each Oxygen to satisfy their octets. (12 electrons used, 0 remaining).

   5. Check Carbon: It only has 4 electrons. Move one lone pair from each Oxygen to form double bonds with Carbon.

   6. Final Structure: O=C=O with two lone pairs on each Oxygen.

3. Draw the Lewis structure for the Nitrate Ion (NO₃⁻).

   1. Count valence electrons: N (5) + O (6x3) + 1 (for the negative charge) = 24 valence electrons.

   2. Nitrogen is the central atom.

   3. Connect N to three O atoms with single bonds. (6 electrons used, 18 remaining).

   4. Fill octets for all three Oxygen atoms. (18 electrons used, 0 remaining).

   5. Check Nitrogen: It only has 6 electrons. Move one lone pair from one Oxygen to form a double bond.

   6. The structure requires brackets and a "-" sign to indicate the charge. Resonance structures exist for this ion.

Practice Questions

1. Draw the Lewis structure for Methane (CH₄).

2. Draw the Lewis structure for Ammonia (NH₃).

3. Draw the Lewis structure for Nitrogen gas (N₂).

4. Determine the Lewis structure for Phosphorus Trichloride (PCl₃).

5. Draw the Lewis structure for Sulfur Hexafluoride (SF₆) and explain why it violates the octet rule.

6. Draw the Lewis structure for the Cyanide ion (CN⁻).

7. Draw the Lewis structure for Formaldehyde (CH₂O).

8. Draw the Lewis structure for Carbon Monoxide (CO) and identify the formal charges.

9. Draw the Lewis structure for Xenon Tetrafluoride (XeF₄).

10. Draw the Lewis structure for the Ammonium ion (NH₄⁺).

Answers & Explanations

1. CH₄: Carbon (4) + 4xH (4) = 8 electrons. Carbon is central with four single bonds to Hydrogen atoms. No lone pairs remain. Carbon has an octet; each Hydrogen has a duet.

2. NH₃: Nitrogen (5) + 3xH (3) = 8 electrons. Nitrogen is central with three single bonds to Hydrogen. One lone pair remains on Nitrogen to complete its octet.

3. N₂: 2xN (10) = 10 electrons. Two Nitrogen atoms connected by a triple bond (6 electrons) with one lone pair on each Nitrogen atom (4 electrons) to satisfy the octet rule.

4. PCl₃: P (5) + 3xCl (21) = 26 electrons. Phosphorus is central with three single bonds to Chlorine. Each Chlorine gets 3 lone pairs, and Phosphorus gets 1 lone pair.

5. SF₆: S (6) + 6xF (42) = 48 electrons. Sulfur is central with six single bonds to Fluorine. Sulfur has 12 electrons, violating the octet rule (expanded octet), which is possible for elements in period 3 or below.

6. CN⁻: C (4) + N (5) + 1 = 10 electrons. Carbon and Nitrogen are connected by a triple bond. Each atom has one lone pair.

7. CH₂O: C (4) + 2xH (2) + O (6) = 12 electrons. Carbon is central, double-bonded to Oxygen and single-bonded to two Hydrogens. Oxygen has two lone pairs.

8. CO: C (4) + O (6) = 10 electrons. Carbon and Oxygen are connected by a triple bond. Each atom has one lone pair. Formal charge: C is -1, O is +1.

9. XeF₄: Xe (8) + 4xF (28) = 36 electrons. Xenon is central with four single bonds to Fluorine. Each Fluorine has 3 lone pairs. Xenon has 2 lone pairs (expanded octet).

10. NH₄⁺: N (5) + 4xH (4) - 1 = 8 electrons. Nitrogen is central with four single bonds to Hydrogen. The entire structure is placed in brackets with a "+" sign.

Quick Quiz

Frequently Asked Questions

What is the octet rule in Lewis structures?

The octet rule is a chemical rule of thumb that reflects the observation that atoms of main-group elements tend to combine in such a way that each atom has eight electrons in its valence shell. This configuration provides the same electronic stability as a noble gas.

How do you calculate the total number of valence electrons for a polyatomic ion?

To find the total valence electrons for an ion, sum the valence electrons of all atoms involved, then add electrons equal to the magnitude of a negative charge or subtract electrons equal to the magnitude of a positive charge. This total is used to build the Lewis structure.

Which atoms are exceptions to the octet rule?

Hydrogen and Helium follow the duet rule (2 electrons), while Boron and Beryllium often form electron-deficient compounds with fewer than 8 electrons. Additionally, elements in the third period and below, such as Sulfur or Phosphorus, can have expanded octets totaling more than 8 electrons.

What is a formal charge and why is it useful?

Formal charge is the hypothetical charge assigned to an atom in a molecule, calculated by subtracting the number of non-bonding electrons and half the bonding electrons from the atom's total valence electrons. It helps identify the most stable Lewis structure when multiple resonance forms are possible.

When should you use multiple bonds in a Lewis structure?

Multiple bonds, such as double or triple bonds, should be used when the central atom does not have a full octet after all valence electrons have been distributed as single bonds and lone pairs on terminal atoms. Moving lone pairs to form shared bonds satisfies the central atom's octet.

What are resonance structures in Lewis diagrams?

Resonance structures occur when more than one valid Lewis structure can be drawn for a single molecule or ion, differing only in the position of electrons. The actual molecule is a hybrid of these structures, often resulting in bond lengths that are intermediate between single and double bonds.

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