4.2.1. Fundamentals of Covalent Bonding

A covalent bond forms when two (or more) nonmetal atoms share one or more pairs of electrons in order to reach stable electron configurations.

A covalent bond is a type of chemical bond where atoms share electrons. This arrangement allows the atoms to achieve a more stable electron configuration, often resembling that of noble gases, which have full valence shells. For example, two hydrogen atoms (each has one electron) can share their electrons to form H₂, achieving a stable duet configuration.

Think of sharing a dessert at a party. When two friends (nonmetals) decide to share a piece of cake (electrons), both can enjoy the treat together instead of trying to get their own, just as atoms share electrons to feel 'full' or stable.

Types of covalent bonds:

• Single bond: one shared pair (2 electrons) → represented by a single dash (–).
• Double bond: two shared pairs (4 electrons) → represented by two dashes (=).
• Triple bond: three shared pairs (6 electrons) → represented by three dashes (≡).

Covalent bonds can vary by the number of electron pairs that are shared between atoms:

1. Single bonds involve one pair of shared electrons, like in H₂ or Cl₂.
2. Double bonds involve two pairs of electrons, as seen in molecules like O₂ (oxygen).
3. Triple bonds are found in nitrogen (N₂), where three pairs of electrons are shared. The more bonds formed, the stronger the connection between the atoms, affecting the molecule's properties.

Imagine the strength of a friendship: a single bond is like a casual acquaintance, a double bond is a close friend who you know well, and a triple bond is like your best friend, with a very strong connection that is hard to break.

Lewis structures (electron-dot diagrams):

• Count total valence electrons from all atoms.
• Place the least electronegative atom (except hydrogen) at the center; surround others.
• Connect atoms by single bonds; subtract two electrons per bond from the total.
• Distribute remaining electrons as lone pairs to satisfy the octet rule (or duet for H).
• Form double or triple bonds if necessary to ensure each atom (other than hydrogen) has eight electrons.
• Check formal charges to ensure the most stable (lowest magnitude) distribution; put negative charges on more electronegative atoms if needed.

Lewis structures help visualize the bonding in molecules. The steps include:

1. Count all the valence electrons from the atoms involved.
2. Identify a central atom, usually the least electronegative.
3. Connect the central atom to surrounding atoms with single bonds, which use up some valence electrons.
4. Distribute any additional electrons to fulfill the octet rule for each atom, forming bonds as needed.
5. Finally, assign formal charges to ensure a stable configuration, making adjustments if necessary.

Picture planning a party. You first count how many guests (valence electrons) can fit (octet rule) and decide who should sit where (central atom). You connect guest chairs together (single bonds) and ensure everyone has enough space to sit comfortably before making any adjustments (double or triple bonds).

Example 1: Lewis structure of carbon dioxide (CO₂)

• Total valence electrons: C (4) + 2 × O (2 × 6) = 4 + 12 = 16 electrons.
• Place C in the center, two O atoms on either side.
• Connect C–O with single bonds: two bonds use 4 electrons, leaving 12 electrons to distribute.
• Distribute 12 electrons as lone pairs: give each O three lone pairs (3 × 2 × 2 = 12 electrons). Now each O has an octet (6 electrons in lone pairs + 2 in C–O bond). Carbon has only 4 electrons (2 bonds), short of an octet.
• Convert lone pairs on O to bonding pairs: move one lone pair from each O to form a second bond to C → two C=O double bonds. Now C has 8 electrons (4 bonds); each O has 8 electrons (4 from two bonds, 4 from two lone pairs).
• Check formal charges: For each O: valence = 6; nonbonding = 4; bonding = 4 → FC = 6 – 4 – ½(4) = 0. For C: valence = 4; nonbonding = 0; bonding = 8 → FC = 4 – 0 – ½(8) = 0.
• Final Lewis structure: O=C=O, with two lone pairs on each O.

In this example, we build the Lewis structure for CO₂ step-by-step:

1. First, we calculate the total number of electrons we have to work with (16).
2. Next, we place carbon in the center because it is less electronegative than oxygen.
3. We connect each oxygen to carbon with single bonds, using up some electrons.
4. Following that, we assign lone pairs to oxygen atoms to complete their octets and adjust to form double bonds, ensuring each atom has 8 electrons for stability.
5. Finally, we verify that the formal charges are zero, indicating a stable structure with neutral overall charge.

Imagine you’re trying to balance a team project. You need to ensure every member (atom) contributes equally (has enough electrons) to complete the task successfully (achieve a stable configuration). By redistributing tasks to build strong connections, you achieve perfect teamwork just like how atoms share electrons to form CO₂.

Bond order and bond length:

• Bond order = number of shared electron pairs (e.g., single = 1, double = 2, triple = 3).
• As bond order increases, bond length decreases (atoms are pulled closer) and bond energy (strength) increases.

Bond order describes the number of shared pairs of electrons between atoms. In general, a higher bond order (more shared electron pairs) results in shorter and stronger bonds. For instance, a triple bond (N≡N in N₂) is much shorter and stronger than a single bond (H–H). As atoms share more electrons and form stronger bonds, they tend to pull closer together, which decreases the bond length and increases the bond energy.

Think of it like a friendship network: the more friends (shared electron pairs) someone has, the tighter their group becomes (shorter bond length) and the more supportive the network is (stronger bond energy). When individuals invest more time and effort into the friendships, the connections become stronger and more meaningful.

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