Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Chemical bond ↓


Covalent bonds


Covalent bonding is a cornerstone concept in the field of chemistry. This type of bond is the fundamental way atoms come together to form molecules, and it is essential to the structure of organic compounds. In this lesson, we will delve deeply into the intricacies of covalent bonds. We will provide a detailed explanation suitable for undergraduate students, using both textual and visual examples.

What is a covalent bond?

Covalent bonding is a type of chemical bond in which two atoms share one or more electron pairs. By sharing electrons, each atom achieves a more stable electronic configuration, often emulating the electron structure of the nearest noble gas. This process is the opposite of ionic bonding, where electrons are transferred from one atom to another.

Why do atoms form covalent bonds?

Atoms form covalent bonds to reach a stable electronic arrangement. Most atoms are unstable in their pure form because their outer electron shells are not filled. In covalent bonding, atoms share electrons to fill these outer shells, leading to lower energy states and greater stability. The rule of thumb is that atoms are usually most stable when their outer shells mimic those of the noble gases, which have full valence shells.

Octave rule

The octet rule is a general guideline suggesting that atoms are most stable when they have eight electrons in their valence shell. Many elements attempt to achieve this octet configuration through the formation of covalent bonds. While this rule applies primarily to elements in the second row of the periodic table, there are many exceptions among the heavier elements and transition metals.

Types of covalent bonds

Single bonds

A single covalent bond is formed when two atoms share a pair of electrons. An example of a single bond is found in the molecule of hydrogen (H2). Each hydrogen atom contributes one electron, sharing it with another to form a bond.

H - H

Double bonds

In a double bond, two pairs of electrons are shared between the same two atoms. A typical example of a molecule with a double bond is oxygen (O2), where each oxygen atom shares its two electrons.

O = O

Triple bond

In a triple bond, three electron pairs are shared. The nitrogen molecule (N2) is a classic example, in which each nitrogen atom shares three electrons with the other.

N ≡ N

Polar and nonpolar covalent bonds

When electrons are shared equally between two atoms in a bond, it is called a nonpolar covalent bond. This usually occurs between atoms of the same element or different elements with similar electronegativities.

In polar covalent bonds, the electrons are shared unequally, leading to a slight separation of charges in the molecule. For example, in the water molecule (H2O), oxygen is more electronegative than hydrogen, so it attracts the shared electrons more strongly.

Visualization of covalent bonds

Covalent bonds can be visualized through several representations, including Lewis structures, which show the arrangement of electrons in the molecule. Lewis structures help us understand how atoms share electrons and what configurations result.

H - O - H (Water: H₂O)

Another notable representation is the structural formula, which shows the arrangement of atoms without the electron dots.

For example:

H  C = C / H (Ethylene: C₂H₄)

Strength and length of covalent bond

The strength and length of a covalent bond depend on a variety of factors, including the types of atoms involved and the number of shared electron pairs. Generally, the more electrons shared between atoms, the stronger and shorter the bond. Therefore, triple bonds are stronger and shorter than double bonds, which in turn are stronger and shorter than single bonds.

Exceptions to the octave rule

While the octet rule is a helpful guideline, many exceptions exist. For example, elements such as hydrogen and helium are stable with only two electrons in their valence shell. Also, molecules with an odd number of electrons, such as nitric oxide (NO), may not obey the octet rule for all atoms.

Complex ions and molecules with central atoms from period 3 and beyond can hold more than eight electrons due to the availability of d orbitals; phosphorus pentachloride (PCl5) and sulfur hexafluoride (SF6) are good examples.

Resonance structures

Sometimes, more than one valid Lewis structure can be formed for a molecule. In such cases, the actual structure is a resonance hybrid of several structures. The classic example of this is the benzene molecule (C6H6), which can be represented as:

C₆H₆: ____ /  | | | | ____/

The concept of resonance stabilizes the molecule by displacing the electrons to multiple locations, resulting in the lowering of the overall energy of the molecule.

Coordinate covalent bond

Coordinate covalent bonds, also called dative bonds, occur when one atom provides both electrons for bonding. A typical example of this is the ammonium ion (NH4+), where a nitrogen atom donates a lone pair to bond with a proton (H+).

Covalent compounds

Covalent bonds often result in the formation of covalent compounds or molecular compounds. These compounds have different characteristics than ionic compounds. They generally have low melting and boiling points, are often gases or liquids at room temperature, and do not conduct electricity in solution.

Conclusion

Covalent bonding is important in the formation of many substances, from the water we drink to the large molecules that make up our cells. Understanding the nature of covalent bonds - how they form, their types, and their exceptions - provides deep insight into the chemical world.

Throughout this detailed exploration of covalent bonds, we have used simple language and illustrative examples to clarify the topic, to ensure that it is accessible to undergraduate chemistry students who wish to understand the fundamentals of chemical bonding.


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Covalent bonds

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