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 → Organic chemistry ↓


Structure and relationships


Organic chemistry is the study of the structure, properties, composition, reactions, and synthesis of carbon-containing compounds. The molecular structure and bonding patterns of organic molecules are important for understanding how these molecules behave, interact, and react with other compounds. In this lesson, we will explore these concepts in detail to provide a comprehensive understanding of the fundamentals of organic chemistry.

Covalent bond

At the core of organic chemistry is covalent bonding. Covalent bonds form when atoms share pairs of electrons. Carbon is unique because it has four valence electrons, which allows it to form stable bonds with other carbon atoms or other elements such as hydrogen, oxygen, nitrogen, and halogens.

Consider the simplest hydrocarbon, methane (CH4 ):

          H
          ,
        HCH
          ,
          H

In methane, the carbon atom forms four single covalent bonds with four hydrogen atoms. Each bond contains a pair of shared electrons between the carbon and hydrogen atom.

Hybridization

The concept of hybridization describes how carbon atoms can form four equivalent bonds. Hybridization is the process of mixing atomic orbitals to form new hybrid orbitals. The type of hybridization can affect the shape and geometry of molecules.

In methane (CH4), carbon undergoes sp3 hybridization, which means that one s orbital combines with three p orbitals to form four equivalent sp3 hybrid orbitals. These orbitals arrange themselves in a tetrahedral shape to minimize repulsion.

sp3

Hydrocarbon structures

Hydrocarbons are compounds that are composed exclusively of carbon and hydrogen. The structure of hydrocarbons can be classified into alkanes, alkenes, and alkynes based on the type of bonds present in them.

Hydrocarbons

Alkanes are saturated hydrocarbons that contain only single bonds. Their general formula is CnH2n+2. The structure varies from linear to branched forms.

Example of ethane (C2H6):

        HH
         ,
          C
         ,
        HH

       ,
       C
      ,
     HH

Alkene

Alkanes contain at least one carbon-carbon double bond. Their general formula is CnH2n. The presence of double bonds leads to sp2 hybridization, resulting in a planar structure.

Example of ethene (C2H4):

        H2C=CH2

Alkynes

Alkynes contain a carbon-carbon triple bond and their general formula is CnH2n-2. The carbons involved in the triple bond have sp hybridization, forming a linear structure.

Example of ethene (acetylene, C2H2):

        HC≡CH

Functional group

In addition to hydrocarbons, organic molecules often contain other atoms, which are arranged in specific groups called functional groups. These groups determine the chemical reactivity and properties of the molecules. Common functional groups include alcohols, ethers, aldehydes, ketones, acids, amines, and esters.

Alcohol

Alcohols contain -OH group bonded to a carbon atom. They are polar molecules and can engage in hydrogen bonding, which affects their physical properties such as boiling point.

Example of methanol (CH3OH):

          H
          ,
        HC-OH
          ,
          H

Ether

In ethers, one oxygen atom is bonded to two alkyl or aryl groups. Their general structure is RO-R'. They have characteristic polar properties, but they do not usually form hydrogen bonds with each other.

Example of diethyl ether (C2H5-OC2H5):

        H3C-CH2-O-CH2-CH3

Aldehyde and ketone

Both aldehydes and ketones contain a carbonyl group (C=O). In aldehydes, the carbonyl group is bonded to at least one hydrogen atom, while in ketones, it is bonded to two carbon atoms.

Example of formaldehyde (HCHO):

        HC=O
         ,
         H

Example of acetone (CH3COCH3):

        H3C-C=O
          ,
         CH3

Intermolecular forces

The physical properties of organic compounds, such as boiling point, melting point, and solubility, are greatly affected by intermolecular forces. These forces are interactions between molecules and can include hydrogen bonding, van der Waals forces, and dipole-dipole interactions.

Hydrogen bonding

Hydrogen bonding occurs when hydrogen bonds to highly electronegative atoms such as nitrogen, oxygen or fluorine. This results in a strong permanent dipole and significant intermolecular forces, which can affect the properties of a substance.

Van der Waals force

Van der Waals forces are weak interactions caused by temporary dipoles that occur when electron clouds overlap in adjacent molecules. They increase with greater surface area and molecular size, affecting boiling and melting points.

The larger the molecule or the greater its surface area, the stronger these forces can be. This is why, in general, long-chain alkanes have higher boiling points than short-chain alkanes.

Dipole–dipole interaction

Dipole-dipole interactions occur between polar molecules where positive and negative dipoles attract each other. These forces are generally stronger than van der Waals forces but weaker than hydrogen bonds.

Resonance and aromaticity

Some molecules may have several valid Lewis structures called resonance structures. The actual structure is a hybrid of these structures and is lower in energy than any single form.

Benzene ( C6H6) is an example, and is also an aromatic compound.

Aromaticity involves cyclic, planar structures with conjugated pi bonds above and below the plane of the ring, which obey Hückel's rule, which states that aromatic compounds should have (4n + 2) pi electrons.

Stereoscopic

Stereochemistry refers to the spatial arrangement of atoms within molecules, which is important in biological activities and reactions. Two common stereochemical concepts are chirality and geometric isomerism.

To right

A chiral molecule is one that cannot be superimposed on its mirror image. These molecules usually have at least one carbon atom with four different groups attached to it, known as stereocenters.

Example of a chiral molecule (2-butanol):

        CH3-CH(OH)-CH2-CH3

Geometrical isomerism

Geometric isomerism occurs in compounds with double bonds or in cyclic compounds, where the spatial arrangement around the bond or ring may differ. Isomers are labeled as "cis" or "trans" depending on the position of the substituents.

Conclusion

Structure and bonding in organic chemistry shape the behavior and reactivity of organic molecules. Understanding covalent bonds, hybridization, functional groups, and intermolecular forces forms the basis for analyzing organic reactions and molecular properties. With a strong understanding of these topics, chemists can design and synthesize a wide range of organic compounds, improve materials, pharmaceuticals, and understand biological systems.


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Structure and relationships

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