Lewis Dot Structure Examples With Answers

The Lewis dot structure, also known as electron dot structure, is a visual representation of the valence electrons in an atom or molecule. It helps us understand how atoms bond together and predict the properties of molecules. By illustrating the arrangement of electrons around atoms, Lewis dot structures provide a simple yet powerful tool for understanding chemical bonding and reactivity.

Understanding the Basics of Lewis Dot Structures

Before diving into examples, let's clarify some fundamental concepts. The valence electrons, which are the electrons in the outermost shell of an atom, are the ones involved in chemical bonding. The Lewis dot structure focuses solely on these valence electrons, representing them as dots around the element's symbol.

Here's a quick rundown of key principles:

Element Symbol: Represents the nucleus and core electrons of the atom.
Dots: Each dot represents one valence electron.
Arrangement of Dots: Dots are placed around the element symbol, starting with one dot on each side (top, bottom, left, right) before pairing them up. This follows Hund's rule, which states that electrons will individually occupy each orbital within a subshell before doubling up in any one orbital.
Octet Rule (and Exceptions): Atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons (an octet), similar to the noble gases. Hydrogen (H) aims for two electrons (a duet). Beryllium (Be) can be stable with four electrons, and Boron (B) can be stable with six electrons. Some elements, like Sulfur (S) and Phosphorus (P), can exceed the octet rule.
Bonds: A shared pair of electrons between two atoms represents a covalent bond, usually depicted as a line. A double bond consists of two shared pairs (two lines), and a triple bond consists of three shared pairs (three lines).

Step-by-Step Guide to Drawing Lewis Dot Structures

Follow these steps to draw accurate Lewis dot structures:

Determine the Total Number of Valence Electrons: Sum the valence electrons of all atoms in the molecule or ion. You can quickly determine the number of valence electrons by looking at the element's group number on the periodic table. For example, oxygen (O) is in group 16 (or 6A), so it has six valence electrons.
Write the Skeletal Structure: This shows how the atoms are connected. The least electronegative element usually goes in the center (except for hydrogen, which is always on the periphery). If carbon is present, it's almost always central. For simple molecules, a central atom is surrounded by other atoms.
Distribute the Electrons as Single Bonds: Place a single bond (a line, representing two electrons) between the central atom and each surrounding atom.
Complete the Octets of the Outer Atoms: Distribute the remaining electrons as lone pairs around the outer atoms (except hydrogen, which only needs two electrons) to satisfy the octet rule.
Place Remaining Electrons on the Central Atom: If there are any electrons left after completing the octets of the outer atoms, place them as lone pairs on the central atom.
Form Multiple Bonds if Necessary: If the central atom does not have an octet, form double or triple bonds by sharing lone pairs from the outer atoms until the central atom achieves an octet.
Consider Formal Charges: Calculate the formal charge on each atom to determine the most plausible Lewis structure. The formal charge is calculated as: Valence Electrons - Non-bonding Electrons - (1/2 Bonding Electrons). The structure with the smallest formal charges on all atoms is generally the most stable. Ideally, the formal charges should be zero, but this isn't always possible. The sum of the formal charges must equal the overall charge of the ion or molecule.
Resonance Structures: If multiple valid Lewis structures can be drawn for a molecule, differing only in the placement of multiple bonds and lone pairs, these are called resonance structures. The actual structure of the molecule is a hybrid of all resonance structures.

Lewis Dot Structure Examples with Answers and Explanations

Let's work through several examples, ranging from simple to more complex, to solidify your understanding.

1. Water (H₂O)

Step 1: Valence Electrons: Hydrogen (H) has 1 valence electron each (x2 = 2), and Oxygen (O) has 6. Total: 2 + 6 = 8 valence electrons.
Step 2: Skeletal Structure: Oxygen is the central atom, with two hydrogen atoms bonded to it: H O H
Step 3: Single Bonds: Draw single bonds between oxygen and each hydrogen: H-O-H. This uses 4 electrons (2 bonds x 2 electrons/bond).
Step 4: Complete Octets: Each hydrogen already has 2 electrons (a duet). We have 4 electrons remaining (8 total - 4 used). Place these as two lone pairs on the oxygen atom.
Step 5: Central Atom Octet: The oxygen atom now has 8 electrons (2 from each bond to hydrogen, and 4 from the two lone pairs).
Final Lewis Structure:
```
   H
   |
H-O:
   ¨
```
(The two dots above and below the O represent the lone pairs)

2. Carbon Dioxide (CO₂)

Step 1: Valence Electrons: Carbon (C) has 4 valence electrons, and Oxygen (O) has 6 each (x2 = 12). Total: 4 + 12 = 16 valence electrons.
Step 2: Skeletal Structure: Carbon is the central atom, with two oxygen atoms bonded to it: O C O
Step 3: Single Bonds: Draw single bonds between carbon and each oxygen: O-C-O. This uses 4 electrons.
Step 4: Complete Octets (Outer Atoms): Place lone pairs around each oxygen atom to give them an octet. Each oxygen needs 6 more electrons (16 total - 4 used = 12 remaining; 6 electrons per oxygen).
Step 5: Central Atom Octet: The carbon atom currently has only 4 electrons (2 from each bond). It needs 4 more to achieve an octet.
Step 6: Form Multiple Bonds: Share one lone pair from each oxygen atom to form a double bond with the carbon atom. This gives carbon an octet (4 from the two double bonds) and each oxygen still has an octet (2 from the double bond and 6 from the lone pairs).
Final Lewis Structure:
```
O=C=O
¨ ¨
```
(Two lone pairs on each oxygen)

3. Ammonia (NH₃)

Step 1: Valence Electrons: Nitrogen (N) has 5 valence electrons, and Hydrogen (H) has 1 each (x3 = 3). Total: 5 + 3 = 8 valence electrons.
Step 2: Skeletal Structure: Nitrogen is the central atom, bonded to three hydrogen atoms.
Step 3: Single Bonds: Draw single bonds between nitrogen and each hydrogen. This uses 6 electrons.
Step 4: Complete Octets (Outer Atoms): Each hydrogen has its duet already.
Step 5: Central Atom Octet: We have 2 electrons remaining (8 total - 6 used). Place these as a lone pair on the nitrogen atom. This gives nitrogen an octet.
Final Lewis Structure:
```
   H
   |
H-N:
   |
   H
```
(The two dots above N represent the lone pair)

4. Methane (CH₄)

Step 1: Valence Electrons: Carbon (C) has 4 valence electrons, and Hydrogen (H) has 1 each (x4 = 4). Total: 4 + 4 = 8 valence electrons.
Step 2: Skeletal Structure: Carbon is the central atom, bonded to four hydrogen atoms.
Step 3: Single Bonds: Draw single bonds between carbon and each hydrogen. This uses all 8 electrons.
Step 4: Complete Octets (Outer Atoms): Each hydrogen has its duet.
Step 5: Central Atom Octet: The carbon atom has an octet (2 electrons from each of the four bonds).
Final Lewis Structure:
```
   H
   |
H-C-H
   |
   H
```

5. Sulfur Dioxide (SO₂)

Step 1: Valence Electrons: Sulfur (S) has 6 valence electrons, and Oxygen (O) has 6 each (x2 = 12). Total: 6 + 12 = 18 valence electrons.
Step 2: Skeletal Structure: Sulfur is the central atom, bonded to two oxygen atoms.
Step 3: Single Bonds: Draw single bonds between sulfur and each oxygen. This uses 4 electrons.
Step 4: Complete Octets (Outer Atoms): Place lone pairs around each oxygen atom to give them an octet. This requires 12 electrons (6 per oxygen). 18 total - 4 used - 12 for oxygen octets = 2 electrons remaining.
Step 5: Central Atom Octet: Place the remaining 2 electrons as a lone pair on the sulfur atom.
Step 6: Form Multiple Bonds: The sulfur atom only has 6 electrons around it (2 from each single bond and 2 from the lone pair). Form a double bond by sharing a lone pair from one of the oxygen atoms.
Step 7: Resonance Structures: Notice that we could have formed the double bond with either oxygen atom. This means that sulfur dioxide has two resonance structures. The actual molecule is a hybrid of these two.
Final Lewis Structures (Resonance Structures):
```
  O=S-O:  <-->  :O-S=O
  ¨ ¨     ¨ ¨
```
(Two lone pairs on the singly bonded oxygen, one lone pair on the doubly bonded oxygen, and one lone pair on the sulfur)

6. Nitrate Ion (NO₃⁻)

Step 1: Valence Electrons: Nitrogen (N) has 5 valence electrons, Oxygen (O) has 6 each (x3 = 18). Add 1 electron for the negative charge. Total: 5 + 18 + 1 = 24 valence electrons.
Step 2: Skeletal Structure: Nitrogen is the central atom, bonded to three oxygen atoms.
Step 3: Single Bonds: Draw single bonds between nitrogen and each oxygen. This uses 6 electrons.
Step 4: Complete Octets (Outer Atoms): Place lone pairs around each oxygen atom to give them an octet. This requires 18 electrons (6 per oxygen). 24 total - 6 used - 18 for oxygen octets = 0 electrons remaining.
Step 5: Central Atom Octet: The nitrogen atom only has 6 electrons around it.
Step 6: Form Multiple Bonds: Form a double bond by sharing a lone pair from one of the oxygen atoms.
Step 7: Resonance Structures: Similar to sulfur dioxide, the double bond can be with any of the three oxygen atoms, resulting in three resonance structures. Remember to enclose the entire structure in brackets and indicate the charge.

Final Lewis Structures (Resonance Structures):

   O=N-O:         :O-N=O         :O-N-O:
   ¨ | ¨    <-->   ¨ | ¨    <-->    ¨ | ¨
     O:              O:              O=
     ¨               ¨                ¨
  [          ]⁻    [          ]⁻    [          ]⁻

7. Sulfate Ion (SO₄²⁻)

Step 1: Valence Electrons: Sulfur (S) has 6 valence electrons, Oxygen (O) has 6 each (x4 = 24). Add 2 electrons for the negative charge. Total: 6 + 24 + 2 = 32 valence electrons.
Step 2: Skeletal Structure: Sulfur is the central atom, bonded to four oxygen atoms.
Step 3: Single Bonds: Draw single bonds between sulfur and each oxygen. This uses 8 electrons.
Step 4: Complete Octets (Outer Atoms): Place lone pairs around each oxygen atom to give them an octet. This requires 24 electrons (6 per oxygen). 32 total - 8 used - 24 for oxygen octets = 0 electrons remaining.
Step 5: Central Atom Octet: The sulfur atom has 8 electrons, satisfying the octet rule. However, sulfur is in the third row and can accommodate more than 8 electrons (expanded octet). While the structure with single bonds and formal charges is valid, minimizing formal charges often leads to a more accurate representation.
Step 6: Minimize Formal Charges (Optional, but Recommended): Form double bonds with two of the oxygen atoms. This reduces the formal charges on the sulfur and those two oxygen atoms to zero.
Final Lewis Structure (with minimized formal charges):
```
    O=         O:
    ||         ||
:O-S-O:   or  :O-S=O
    ||         ||
    O=         O:
  [          ]²⁻  [          ]²⁻
```
This representation minimizes the formal charges, leading to a more stable structure. Note that resonance structures are also possible with this arrangement.

8. Ozone (O₃)

Step 1: Valence Electrons: Oxygen (O) has 6 valence electrons each (x3 = 18). Total: 18 valence electrons.
Step 2: Skeletal Structure: O-O-O
Step 3: Single Bonds: Draw single bonds: O-O-O. This uses 4 electrons.
Step 4: Complete Octets: Add lone pairs to the terminal oxygens to give them octets. This uses 12 electrons (6 per oxygen). 18 - 4 - 12 = 2 electrons remaining.
Step 5: Place Remaining Electrons: Place the remaining two electrons as a lone pair on the central oxygen.
Step 6: Form Multiple Bonds: The central oxygen only has 6 electrons. Form a double bond with one of the terminal oxygens.
Step 7: Resonance Structures: The double bond can be with either of the terminal oxygens.
Final Lewis Structures (Resonance Structures):
```
 O=O-O:   <-->  :O-O=O
 ¨ ¨        ¨ ¨
```

9. Carbon Monoxide (CO)

Step 1: Valence Electrons: Carbon (C) has 4 valence electrons, and Oxygen (O) has 6. Total: 10 valence electrons.
Step 2: Skeletal Structure: C-O
Step 3: Single Bond: Draw a single bond: C-O. This uses 2 electrons.
Step 4: Complete Octets (Attempt): Add lone pairs. To give each atom an octet, we'd need to add 6 electrons to the carbon and 6 electrons to the oxygen, requiring 12 electrons total. We only have 8 remaining (10 total - 2 used).
Step 5: Form Multiple Bonds: We need to form multiple bonds to satisfy the octet rule. Form a triple bond. This uses 6 electrons.
Step 6: Add Remaining Lone Pairs: We have 4 electrons remaining. Add two electrons (one lone pair) to each atom.
Final Lewis Structure:
```
:C≡O:
```

10. Hydrogen Cyanide (HCN)

Step 1: Valence Electrons: Hydrogen (H) has 1, Carbon (C) has 4, and Nitrogen (N) has 5. Total: 1 + 4 + 5 = 10 valence electrons.
Step 2: Skeletal Structure: H-C-N (Hydrogen is always terminal).
Step 3: Single Bonds: Draw single bonds: H-C-N. This uses 4 electrons.
Step 4: Complete Octets (Attempt): Hydrogen has its duet. To complete the octets of carbon and nitrogen with the remaining 6 electrons, we'd need to add 4 to carbon and 2 to nitrogen. Carbon would still be short of an octet.
Step 5: Form Multiple Bonds: Form a triple bond between carbon and nitrogen.
Step 6: Add Remaining Lone Pairs: We have 6 electrons - (3 bonds * 2 electrons/bond) = 0 electrons remaining on C and N, place one lone pair to N to fulfill octet rule.
Final Lewis Structure:
```
H-C≡N:
```

Common Mistakes to Avoid

Forgetting to Count All Valence Electrons: This is the most common mistake. Double-check your periodic table and remember to account for the charge in ions.
Violating the Octet Rule Unnecessarily: While some elements can exceed the octet rule, stick to it whenever possible.
Incorrectly Placing Lone Pairs: Make sure lone pairs are placed on atoms to fulfill the octet rule (or duet for hydrogen).
Ignoring Formal Charges: Use formal charges to evaluate the plausibility of different Lewis structures.
Not Recognizing Resonance Structures: If multiple valid structures exist, consider resonance.
Confusing Lone Pairs and Bonding Pairs: Remember that a line represents a shared pair of electrons (a bond), while dots represent lone pairs.

Advanced Considerations

Expanded Octets: Elements in the third period and beyond (like S, P, Cl, Br, I) can sometimes accommodate more than eight electrons due to the availability of d orbitals. This is common in compounds like SF₆ and PCl₅.
Odd-Electron Species (Free Radicals): Molecules with an odd number of valence electrons, such as nitric oxide (NO), cannot satisfy the octet rule for all atoms. These are called free radicals and are often highly reactive.
Hypervalent Molecules: Molecules with a central atom bonded to more atoms than normally allowed by the octet rule are called hypervalent. Common examples include XeF₄ and IF₅. These structures require expanding the octet of the central atom.

Conclusion

Mastering Lewis dot structures is a crucial step in understanding chemical bonding and molecular properties. By following the steps outlined above and practicing with various examples, you can confidently draw Lewis structures for a wide range of molecules and ions. Remember to pay attention to valence electrons, the octet rule (and its exceptions), formal charges, and resonance structures. This knowledge will provide a strong foundation for further exploration in chemistry.