Understanding covalent bonding is paramount in grasping the architecture of many molecules found in nature. This section explores the essential features of covalent bonding, and why some elements are more inclined to form them than others.
Definition and Characteristics of a Covalent Bond
Covalent bonding arises from the mutual sharing of electrons between two atoms. Unlike ionic bonds, where electrons are transferred, covalent bonds involve the overlapping of atomic orbitals.
- Nature: Non-metal atoms are the primary participants in forming covalent bonds.
- Strength: These bonds are strong and require a substantial amount of energy to break.
- Directional Nature: Covalent bonds have a set direction, which determines the shape of the molecule.
- Electron Sharing: The number of electrons shared can vary, leading to single, double, or triple bonds.
Covalent bonding between two fluorine atoms.
Image courtesy of Jacek FH
The Octet Rule: Introduction and Significance
Atoms form bonds to attain a more stable electron configuration, often resembling that of noble gases. This tendency is termed the octet rule.
- Stability Pursuit: Atoms strive to have eight electrons in their outermost shell, achieving noble gas configuration and maximising stability.
- Significance in Covalent Bonding: Covalent bonding allows atoms to share electrons, helping them attain or come close to this octet configuration.
Methodology for Deducing Lewis Structures
Lewis structures, or electron dot diagrams, represent the arrangement of electrons in a molecule.
- Total Valence Electrons: Count the total number of valence electrons for all atoms in the molecule.
- Central Atom Selection: Choose the least electronegative atom (excluding Hydrogen) as the central atom.
- Placement of Electrons: Arrange atoms around the central atom and place electrons around atoms. Initially, fulfill the octet for external atoms, leaving the central atom for last.
- Multiple Bonds: If the central atom does not have a complete octet, consider multiple bonds.
Lewis structure, using BF3 as an example.
Image courtesy of Jerónimo Cueto
Limitations of the Octet Rule
While the octet rule works for many molecules, there are exceptions.
- Less than Octet: Some molecules like boron trifluoride (BF3) have fewer than eight electrons around an atom.
- Extended Octets: Elements in the third period and beyond can accommodate more than eight electrons due to d orbitals. For example, sulphur hexafluoride (SF6) has 12 electrons around sulphur.
Noble Gases and Their Reluctance to Form Covalent Bonds
Noble gases, with their complete outer electron shells, are intrinsically stable. Their reluctance to bond stems from this stability.
- Inherent Stability: They already possess a complete octet, reducing the necessity to bond with other atoms.
- High Ionisation Energy: Noble gases have high ionisation energies, making it energetically unfavourable to lose electrons.
- Exceptions: Under specific conditions, some noble gases like xenon can form compounds, such as xenon hexafluoride (XeF6).
Electron shell diagram for Argon, with a complete outer shell.
Image courtesy of Pumbaa (original work by Greg Robson)
Electron shell diagram for Xenon, with a complete outer shell.
Image courtesy of Pumbaa (original work by Greg Robson)
The fascinating world of covalent bonds reveals the underlying principles governing molecular formation. As we proceed, the depth of these principles and their nuances will unravel further, offering a comprehensive understanding of molecular interactions.
FAQ
D orbitals come into play with elements in the third period and beyond on the periodic table, allowing these elements to expand beyond the traditional octet. This phenomenon occurs because the d orbitals in these elements can participate in bonding, thereby accommodating more than eight electrons. For instance, in phosphorus pentachloride (PCl5), phosphorus forms five covalent bonds, resulting in ten electrons around it, which is possible because of the availability and involvement of the d orbitals in bonding.
Lewis structures are invaluable because they provide a visual representation of how atoms are bonded in a molecule and how electrons are arranged around them. They give insight into the electron distribution, indicating regions of electron density, which can help predict molecular reactivity. Furthermore, by showing shared or lone pairs of electrons, Lewis structures can guide predictions about molecular geometry using theories like VSEPR. Overall, they serve as foundational tools for chemists to depict and predict molecular properties.
Yes, besides the octet rule, the duet rule is another guideline, specifically for hydrogen. It dictates that hydrogen seeks to have two electrons in its valence shell, achieving the electron configuration of helium. Consequently, hydrogen typically forms one covalent bond, sharing its sole electron with another atom. Another principle is the 18-electron rule for transition metals, where these metals tend to accommodate up to 18 electrons (including valence electrons and those shared in bonds) to achieve a stable configuration.
Double and triple covalent bonds form when atoms share two or three pairs of electrons, respectively, rather than just one. This usually occurs when both atoms involved need more than one electron to complete their octet. The primary difference between these bonds and single covalent bonds is the number of shared electron pairs. Consequently, double and triple bonds are shorter and stronger than single bonds. For instance, the bond in O=O (oxygen molecule) is a double bond, making it stronger and shorter than the single bond in H-H (hydrogen molecule).
The bond length is the distance between the nuclei of two bonded atoms. When multiple covalent bonds (double or triple) form between two atoms, they share more electrons, leading to increased electron-electron repulsion. This repulsion pushes the nuclei closer together, decreasing bond length. Simultaneously, with more electrons being shared, the electrostatic attraction between the positively charged nuclei and shared electrons intensifies, making the bond stronger. Therefore, a triple bond is shorter and stronger than a double bond, which in turn is shorter and stronger than a single bond.
Practice Questions
Ionic bonds arise due to the transfer of electrons from one atom, usually a metal, to another, typically a non-metal. This transfer results in the formation of positively and negatively charged ions, which are held together by electrostatic forces of attraction. On the other hand, covalent bonds involve the mutual sharing of electrons between two non-metal atoms, with no formation of ions. Non-metal atoms have a high electronegativity, making them more inclined to accept electrons. However, when two non-metals come together, neither atom is willing to donate electrons entirely, leading to the sharing of electrons, hence forming covalent bonds.
Boron trifluoride (BF3) presents a classic example of a molecule that does not obey the octet rule. In BF3, boron has only six electrons around it instead of the expected eight. This deviation arises because boron has three valence electrons and forms bonds with three fluorine atoms by sharing these electrons. As a result, boron doesn't achieve a complete octet. This scenario highlights a limitation of the octet rule, indicating that not all atoms, especially those in the second period of the periodic table, need to attain a full octet to form stable compounds.
