How To Draw Lewis Structures For Covalent Compounds

Drawing Lewis structures is a fundamental skill in chemistry, offering a visual representation of the bonding between atoms in a molecule, particularly covalent compounds where atoms share electrons to achieve stability. These structures help predict molecular geometry, polarity, and reactivity, serving as a cornerstone for understanding chemical behavior.

Understanding the Basics: Why Lewis Structures Matter

Lewis structures, also known as electron dot diagrams, depict the valence electrons of atoms within a molecule. Valence electrons are the outermost electrons involved in chemical bonding. By representing these electrons as dots around the atomic symbol, we can visualize how atoms share them to form covalent bonds. Understanding the octet rule, which states that atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons (except for hydrogen, which aims for two), is crucial for accurately drawing Lewis structures. This visual representation allows chemists to predict molecular properties and understand reaction mechanisms, providing a powerful tool for studying and predicting chemical behavior.

Step-by-Step Guide to Drawing Lewis Structures

Creating Lewis structures involves a systematic approach. This detailed walkthrough ensures accuracy and a solid understanding of the underlying principles.

1. Determine the Total Number of Valence Electrons

Begin by identifying each atom in the molecule and determining the number of valence electrons it possesses. This information can be readily obtained from the element's group number in the periodic table.

• Group 1 elements (alkali metals) have 1 valence electron.
• Group 2 elements (alkaline earth metals) have 2 valence electrons.
• Group 13 elements have 3 valence electrons.
• Group 14 elements have 4 valence electrons.
• Group 15 elements have 5 valence electrons.
• Group 16 elements have 6 valence electrons.
• Group 17 elements (halogens) have 7 valence electrons.
• Group 18 elements (noble gases) have 8 valence electrons (except helium, which has 2).

Once you know the number of valence electrons for each atom, add them together to find the total number of valence electrons available for the molecule. For polyatomic ions, add electrons for negative charges and subtract electrons for positive charges.

Example: Carbon Dioxide (CO₂)

• Carbon (C) is in Group 14 and has 4 valence electrons.
• Oxygen (O) is in Group 16 and has 6 valence electrons.
• Total valence electrons = 4 (from C) + 2 * 6 (from O) = 16 valence electrons

2. Draw the Skeletal Structure

The skeletal structure shows how atoms are connected in the molecule. The least electronegative atom usually occupies the central position. Electronegativity generally increases across a period and decreases down a group in the periodic table. Hydrogen is always a terminal atom (an atom that is only bonded to one other atom).

Tips for Drawing the Skeletal Structure:

• Place the least electronegative atom in the center (except for hydrogen).
• Connect the other atoms to the central atom with single bonds.
• Symmetrical arrangements are often correct.

Example: Carbon Dioxide (CO₂)

• Carbon is less electronegative than oxygen, so it goes in the center: O C O

3. Distribute Electrons to Form Single Bonds

Place a pair of electrons (represented as a line) between each pair of atoms to form a single bond. Each single bond represents two shared electrons.

Example: Carbon Dioxide (CO₂)

• O-C-O (Each dash represents two electrons, so 4 electrons have been used.)

4. Distribute the Remaining Electrons as Lone Pairs

Distribute the remaining valence electrons as lone pairs around the atoms, starting with the most electronegative atoms (the terminal atoms), to satisfy the octet rule (or duet rule for hydrogen).

Example: Carbon Dioxide (CO₂)

• We started with 16 valence electrons and have used 4 for the single bonds, leaving 12 electrons. Place these as lone pairs around the oxygen atoms:
  • :O-C-O:
  • Place three lone pairs (6 electrons) on each oxygen atom to satisfy the octet rule: :Ö-C-Ö:

5. Form Multiple Bonds if Necessary

If any atoms (especially the central atom) do not have an octet, form multiple bonds (double or triple bonds) by sharing lone pairs from the surrounding atoms.

Example: Carbon Dioxide (CO₂)

• In the current structure, each oxygen atom has an octet, but the carbon atom only has four electrons around it. To give carbon an octet, move one lone pair from each oxygen atom to form double bonds:
  • Ö=C=Ö

Now, each atom has an octet: the carbon atom has four shared pairs (8 electrons), and each oxygen atom has two lone pairs and two shared pairs (8 electrons).

6. Check for Formal Charges (Optional but Recommended)

Formal charge helps determine the most plausible Lewis structure when multiple structures are possible. It is calculated using the formula:

Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 * Bonding Electrons)

The best Lewis structure minimizes formal charges, with the most electronegative atoms having negative formal charges and the least electronegative atoms having positive formal charges.

Example: Carbon Dioxide (CO₂)

• For Carbon: Formal Charge = 4 - 0 - (1/2 * 8) = 0
• For Oxygen: Formal Charge = 6 - 4 - (1/2 * 4) = 0
• Since all formal charges are zero, this is a very stable and plausible Lewis structure.

Examples of Drawing Lewis Structures

Let's walk through a few more examples to solidify your understanding.

Example 1: Water (H₂O)

1. Total Valence Electrons:
   • Hydrogen (H) has 1 valence electron.
   • Oxygen (O) has 6 valence electrons.
   • Total = 2 * 1 (from H) + 6 (from O) = 8 valence electrons
2. Skeletal Structure:
   • Oxygen is the central atom: H O H
3. Distribute Electrons to Form Single Bonds:
   • H-O-H (4 electrons used)
4. Distribute Remaining Electrons as Lone Pairs:
   • We have 8 - 4 = 4 electrons remaining. Place two lone pairs on the oxygen atom to complete its octet: H-Ö-H
5. Check Octets:
   • Hydrogen has a duet (2 electrons), and oxygen has an octet (8 electrons).
6. Formal Charges:
   • For Hydrogen: Formal Charge = 1 - 0 - (1/2 * 2) = 0
   • For Oxygen: Formal Charge = 6 - 4 - (1/2 * 4) = 0
   • The Lewis structure for water is H-Ö-H, with two lone pairs on the oxygen atom.

Example 2: Ammonia (NH₃)

1. Total Valence Electrons:
   • Nitrogen (N) has 5 valence electrons.
   • Hydrogen (H) has 1 valence electron.
   • Total = 5 (from N) + 3 * 1 (from H) = 8 valence electrons
2. Skeletal Structure:
   • Nitrogen is the central atom: H N H H
3. Distribute Electrons to Form Single Bonds:
   • H-N-H (6 electrons used) | H
4. Distribute Remaining Electrons as Lone Pairs:
   • We have 8 - 6 = 2 electrons remaining. Place one lone pair on the nitrogen atom: H-N-H | H
5. Check Octets:
   • Each hydrogen has a duet (2 electrons), and nitrogen has an octet (8 electrons).
6. Formal Charges:
   • For Hydrogen: Formal Charge = 1 - 0 - (1/2 * 2) = 0
   • For Nitrogen: Formal Charge = 5 - 2 - (1/2 * 6) = 0
   • The Lewis structure for ammonia is: H-N-H | H With one lone pair on the nitrogen atom.

Example 3: Sulfate Ion (SO₄²⁻)

1. Total Valence Electrons:
   • Sulfur (S) has 6 valence electrons.
   • Oxygen (O) has 6 valence electrons.
   • Add 2 electrons for the 2- charge.
   • Total = 6 (from S) + 4 * 6 (from O) + 2 = 32 valence electrons
2. Skeletal Structure:
   • Sulfur is the central atom: O S O O O
3. Distribute Electrons to Form Single Bonds:
   • O-S-O (8 electrons used) | | O O
4. Distribute Remaining Electrons as Lone Pairs:
   • We have 32 - 8 = 24 electrons remaining. Place three lone pairs on each oxygen atom to complete their octets: :Ö-S-Ö: | | :Ö Ö:
5. Check Octets:
   • Each oxygen and sulfur atom has an octet.
6. Formal Charges:
   • For Sulfur: Formal Charge = 6 - 0 - (1/2 * 8) = +2
   • For Oxygen (each): Formal Charge = 6 - 6 - (1/2 * 2) = -1
   • To minimize formal charges, we can form double bonds. Two of the oxygen atoms can form double bonds with the sulfur atom: Ö=S=Ö | | :Ö Ö:
   • New Formal Charges:
     • For Sulfur: Formal Charge = 6 - 0 - (1/2 * 12) = 0
     • For Oxygen (double-bonded): Formal Charge = 6 - 4 - (1/2 * 4) = 0
     • For Oxygen (single-bonded): Formal Charge = 6 - 6 - (1/2 * 2) = -1
   • The overall charge of the ion is 2-, so the Lewis structure is: [Ö=S=Ö]²⁻ | | :Ö Ö:

Common Mistakes to Avoid

• Incorrectly Counting Valence Electrons: Always double-check the number of valence electrons for each atom.
• Forgetting to Account for Charge in Ions: Remember to add or subtract electrons for polyatomic ions.
• Violating the Octet Rule: Ensure that most atoms (except hydrogen) have eight electrons around them.
• Placing Too Many Electrons: The total number of electrons in the Lewis structure should match the total number of valence electrons calculated.
• Ignoring Formal Charges: Use formal charges to evaluate the stability of different possible Lewis structures.

Advanced Concepts: Resonance Structures and Exceptions to the Octet Rule

Resonance Structures

Some molecules and ions cannot be accurately represented by a single Lewis structure. In these cases, we use resonance structures, which are multiple Lewis structures that, when considered together, provide a more accurate representation of the molecule. Resonance structures differ only in the arrangement of electrons, not in the arrangement of atoms. The true structure is a hybrid or average of all the resonance structures.

Example: Ozone (O₃)

1. Total Valence Electrons: 3 * 6 = 18 valence electrons
2. Skeletal Structure: O O O
3. Possible Lewis Structures:
   • O=O-O: and O-O=O Each oxygen atom has an octet in both structures, but the double bond can be in two different positions.
4. Resonance Hybrid: The true structure of ozone is a hybrid of these two resonance structures, with the electrons delocalized over all three oxygen atoms.

Exceptions to the Octet Rule

While the octet rule is a useful guideline, there are exceptions:

• Incomplete Octets: Some atoms, like beryllium (Be) and boron (B), can be stable with fewer than eight electrons around them. For example, in boron trifluoride (BF₃), boron has only six electrons.
• Expanded Octets: Atoms in the third period and beyond (like sulfur, phosphorus, and chlorine) can accommodate more than eight electrons around them due to the availability of d-orbitals. For example, in sulfur hexafluoride (SF₆), sulfur has twelve electrons.
• Odd Number of Electrons: Molecules with an odd number of valence electrons are called free radicals and cannot satisfy the octet rule for all atoms. For example, nitrogen monoxide (NO) has 11 valence electrons.

Software and Tools for Drawing Lewis Structures

While drawing Lewis structures manually is essential for understanding the underlying principles, several software tools and websites can assist in visualizing and verifying your structures:

• ChemDraw: A professional chemical drawing tool widely used in research and industry.
• MarvinSketch: A free chemical drawing tool that offers basic Lewis structure capabilities.
• Online Lewis Structure Generators: Many websites provide tools that generate Lewis structures based on the chemical formula you input.

Conclusion

Drawing Lewis structures is a vital skill for anyone studying chemistry. By following the step-by-step guide, understanding the octet rule, and recognizing common exceptions, you can accurately represent the bonding in covalent compounds. These structures provide valuable insights into molecular properties, reactivity, and overall chemical behavior, forming a cornerstone for further exploration in the world of chemistry.