Introduction to organic chemistry, history and hybrid orbital functions

 Organic chemistry can be seen as a field driven by curiosity about organic compounds found in living organisms. 

Most organic compounds, which are characterized by containing carbon, exhibit a wide range of applications. For instance, they are integral to the composition of hair (proteins), DNA, as well as various types of food and medicine.

In essence, organic chemistry deals with substances derived from plants and animals, or those found in them.

Organic chemistry has its roots in the mid-1700s, when it was recognized that compounds obtained from plants and animals exhibit distinct properties compared to inorganic compounds. These organic compounds typically have lower melting points, making them challenging to decompose, separate, and purify.

In 1816, Michel Eugène Chevreul created soap by reacting animal fat with alkali, resulting in the formation of fatty acids—an organic compound. This experiment served to disprove the prevailing concept of vitalism, which attributed special life force properties to organic compounds. Chevreul's work demonstrated that organic compounds could be synthesized from simpler, inorganic precursors, marking a significant advancement in the understanding of organic chemistry.

In 1828, Friedrich Wöhler achieved a groundbreaking milestone by synthesizing urea—an organic compound—from inorganic starting materials, specifically ammonium cyanate. This landmark experiment definitively refuted the theory of vitalism, which held that organic compounds could only be produced by living organisms. Wöhler's synthesis of urea from inorganic salts demonstrated that organic compounds could be created through purely chemical processes, revolutionizing the field of organic chemistry.




Carbon indeed holds paramount importance in organic chemistry. Carbon constitutes the backbone of most organic compounds, making up over 90% of their composition. It typically resides in Group 4A of the periodic table, also known as Group 14 or the "carbon group." This positioning reflects carbon's ability to form four bonds with other atoms, allowing for the creation of diverse and complex molecular structures—a fundamental characteristic underlying the vast array of organic compounds observed in nature.




Indeed, carbon's ability to form four covalent bonds due to its four valence electrons contributes to the high molecular weight of compounds with a high carbon content. Additionally, organic compounds commonly include elements such as hydrogen (H), nitrogen (N), oxygen (O), fluorine (F), phosphorus (P), sulfur (S), chlorine (Cl), bromine (Br), and iodine (I). Familiarizing oneself with these elements can facilitate the memorization and understanding of organic chemistry concepts. Before delving into organic chemistry, it's beneficial to gain a basic understanding of atomic structure and terminology. Some key terms to acquaint oneself with include the inductive effect, resonance effect, steric effect, stereochemistry (including stereo-specific, chemo-selective, stereo-selective, and regio-selective reactions), and light-transmissive selection. Understanding these concepts at a fundamental level is crucial for navigating the complexities of organic chemistry. Fortunately, the structure of atoms is relatively straightforward, providing a solid foundation for further exploration in the field.





We studied when we were young. 





The nucleus, composed of protons and neutrons, has a size of about 10^-15m. The atomic number represents the number of protons in the nucleus, while isotopes vary in the number of neutrons. For practical purposes, we often use the average atomic weight.
The size of an atom is approximately 2*10^-10m or 200 picometers (pm), commonly expressed as 2Å (Angstrom).
Understanding orbital functions, such as s, p, d, and f, can be challenging. The square of the wave function (ψ^2) in these orbitals represents the probability of finding an electron at a given location, which is an important concept to grasp.



In addition, if it is S, it is round, if it is p, it is a dumbbell type, if d, it is a clover model. 
As you enter the electronic layout, you only need to know two principles and one rule: three. 
1. It is a stacking principle. 
You can observe the stacking order in orbitals: 1s, 2s, 2p, 3s, 3p, 4s, 3d.
Pauli's exclusion principle dictates that within a single orbital, only two electrons can exist, with each having opposite spins.
Hund's Rule states that electrons tend to fill orbitals of the same energy level singly before pairing up, maximizing total electron spin.

3. Hunt's Rule



In a stable atom, the stacking principle dictates that electrons fill orbitals in the order of 1s, 2s, and then 2p. Within the 2p orbitals (designated as 2px, 2py, and 2pz), electrons must first occupy each orbital individually before pairing up, with each electron possessing opposite spins.

Additionally, understanding Van't Hoff and Le Bel's three-dimensional tetrahedral structure, the octet rule, electrostatic attraction, and covalent bonding is beneficial. These concepts contribute to comprehending the spatial arrangement of molecules, adherence to the octet rule, the attractive forces between charged particles, and the sharing of electron pairs in covalent bonds.

Furthermore, familiarity with non-covalent electron pairs and Lewis dot structures aids in discerning differences in bonding. Lewis dot structures illustrate the distribution of valence electrons around atoms, while line bond structural formulas depict the connectivity of atoms and the arrangement of bonds between them.



When delving into valence bonding theory (molecular orbital theory), it's crucial to grasp the concept of cylindrical overlapping, which is exemplified by sigma bonding in hydrogen—an essential example of overlapping orbital functions.

One of the most significant hybrid orbital functions is illustrated by Sp3 hybridization in the structure of CH4.




Carbon has 1s^2, 2s^2, 2p^2 orbital functions, resulting in the same C-H bonds in CH4 undergoing hybridization.

The + part of the 2s orbital overlaps with the + and - parts distributed between the nodes in the 2p orbitals. Adding to the + part increases its size, while offsetting from the - part decreases it, resulting in an asymmetric structure.

Hence, it's evident that the SP3 hybrid orbital function is larger than the S or P orbital function due to its more efficient overlap with other atomic orbitals.

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