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


Chemical bonding is a fundamental concept in chemistry that describes how atoms come together to form molecules and compounds. At its core, chemical bonding involves the interactions between electrons in the outermost shells of atoms, causing them to stick together in various configurations. The study of chemical bonding allows us to understand the structures, properties, and behaviors of various substances.

Types of chemical bonds

There are several primary types of chemical bonds, each of which has different properties and characteristics. The three main types of chemical bonds are ionic bonds, covalent bonds, and metallic bonds.

Ionic bond

Ionic bonds form when electrons are transferred from one atom to another, resulting in the formation of ions. This type of bond is usually formed between metals and nonmetals. The metal atom loses electrons to become a positively charged ion, or cation, while the nonmetal gains electrons to become a negatively charged ion, or anion. Ionic bonds are formed by the attraction between opposite charges.

For example, in the formation of sodium chloride (NaCl), sodium (Na) loses an electron to form Na+, and chlorine (Cl) gains an electron to form Cl–. Ionic bonding is the electrostatic attraction between these two oppositely charged ions.

Na → Na+ + e-
Cl + e- → Cl-
Na+ + Cl- → NaCl

Visual example of ionic bond:

Na Chlorine Na+ Cl-

Covalent bonds

Covalent bonds form when two atoms share one or more electron pairs. This type of bond typically forms between nonmetal atoms. In covalent bonds, the shared electrons allow each atom to achieve a complete outer electron shell, creating a more stable configuration.

A classic example of a covalent bond is found in the water (H2O) molecule. Each hydrogen atom shares one of its electrons with an oxygen atom, and the oxygen atom shares one of its electrons with each hydrogen atom, resulting in two covalent bonds.

H• + •H → H:H (Covalent Bond in H2O)

Visual example of a covalent bond:

O H H

Metal bonding

Metallic bonds are found in metallic elements where atoms are held together in a lattice by a set of shared electrons. In metallic bonds, the electrons are delocalized, meaning they are not attached to any particular atom but move freely throughout the structure. This delocalization is responsible for many of the properties of metals, such as electrical conductivity, ductility, and lustre.

A modification of this electron-sea model can be seen in metals such as copper (Cu) and aluminum (Al), where metallic bonding holds the metal ions together through shared electrons.

Visual example of a metallic bond:

Bond polarity and electronegativities

The concept of electronegativity is important in understanding the nature of chemical bonds. Electronegativity refers to the ability of an atom to attract electrons towards itself in a chemical bond. The difference in electronegativities between two bonded atoms can determine the polarity of the bond.

A nonpolar covalent bond is formed when the bonded atoms have similar electronegativities, resulting in equal sharing of electrons. For example, in a molecular hydrogen (H2) molecule, both hydrogen atoms have similar electronegativities, resulting in a nonpolar bond.

H:H (Nonpolar Covalent Bond in H2)

In contrast, a polar covalent bond occurs when there is a significant difference in electronegativities between the bonded atoms. This results in an uneven distribution of electron density, with the more electronegative atom acquiring a partial negative charge, and the less electronegative atom acquiring a partial positive charge. A familiar example is the water molecule (H2O) where the oxygen atom is more electronegative than the hydrogen atoms.

Hδ+ - Oδ- - Hδ+ (Polar Covalent Bond in H2O)

Intermolecular forces

While chemical bonds hold atoms together within a molecule, intermolecular forces (IMFs) are the attractive forces that exist between molecules. These forces play an essential role in determining the physical properties of substances, such as their melting and boiling points.

Types of intermolecular forces

1. London dispersion forces: These are the weakest intermolecular forces and arise from temporary fluctuations in electron density, creating short-lived dipoles. All molecules exhibit London dispersion forces, but they are particularly important in nonpolar molecules.

2. Dipole-dipole interactions: occur in polar molecules where the positive and negative ends of different molecules attract each other. The strength of dipole-dipole interactions depends on the polarity of the molecules involved.

3. Hydrogen bond: A special type of dipole-dipole interaction that occurs when hydrogen is bound to highly electronegative atoms such as nitrogen, oxygen or fluorine. Hydrogen bonds are relatively strong among intermolecular forces and are important in maintaining the structure of biological molecules such as water and DNA.

Visual example of intermolecular forces:

H2O H2O

Lewis structures and VSEPR theory

Lewis structures are a method of modeling molecules by showing how atoms are connected to one another and what lone pairs of electrons might be present. They are useful for predicting the geometry of molecules when used in conjunction with valence shell electron pair repulsion (VSEPR) theory.

VSEPR theory explains molecular shapes on the basis that electron pairs, both bonded and lone pairs, repel each other. By minimizing this repulsion, the theory helps predict the arrangement of atoms in a molecule.

Example of Lewis Structure for Water:
H / O  H

According to VSEPR theory, the molecular shape of water is bent due to the two lone electron pairs present on oxygen, which push the O-H bonds downwards.

Conclusion

Understanding chemical bonds is the key to understanding the principles of chemistry. The interactions between atoms through ionic, covalent, and metallic bonds help define myriad compounds and materials. In addition, studying intermolecular forces deepens our understanding of the properties of matter, and using models such as Lewis structures and VSEPR theory allows us to look at molecular geometry. By combining these concepts, we can better understand the complexity of chemical structures and reactions.


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Chemical bond

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