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 → Hydrocarbons ↓


Hydrocarbons


Alkanes are a fundamental group of compounds in organic chemistry. They are saturated hydrocarbons, meaning they contain only single bonds between carbon atoms. Alkanes are the simplest type of hydrocarbon and can be found in a variety of natural sources such as crude oil and natural gas. In this lesson, we will discuss their structure, properties, reactions, and uses. Understanding alkanes is important for students to understand more complex organic chemistry concepts.

Chemical structure of alkenes

An alkane contains carbon and hydrogen atoms that are completely connected by single bonds. The general formula for an alkane is C n H 2n+2, where n is the number of carbon atoms. This formula shows that the alkanes are saturated, which means they are completely filled with hydrogen atoms.

Take methane as the simplest alkane:

CH 4

In this one carbon atom is bonded to four hydrogen atoms.

Linear and branched alkanes

Alkanes can be linear (straight-chain) or branched. Linear alkanes have carbon atoms linked in a single continuous chain. Branched alkanes have one or more carbon chains attached to the main chain.

Here's a simple example of a straight-chain alkane:

CH 3 -CH 2 -CH 2 -CH 2 -CH 3 (pentane)

and a branched alkene:

CH 3 , CH 3 -CH-CH 2 -CH 3 (2-methylbutane)

Naming of alkenes

The naming of alkanes follows a systematic method established by the International Union of Pure and Applied Chemistry (IUPAC). The names are derived from the number of carbon atoms in the longest continuous chain. The suffix -ane is always added to indicate that they are alkanes. Here are some examples:

  • CH 4 - methane
  • C 2 H 6 - ethane
  • C 3 H 8 - propane
  • C 4 H 10 - butane

For branched alkanes, the chain is numbered so that the number of side chains is as minimal as possible. For example:

CH 3 -CH-CH 2 -CH 3 , CH 3 Name: 2-Methylbutane

Properties of alkenes

Alkanes exhibit several unique properties due to their saturated nature. Let us look at some of these properties:

Physical properties

  • Boiling and melting points: Generally, as the number of carbon atoms increases, the boiling and melting points of alkanes also increase. This is due to stronger van der Waals forces in larger molecules.
  • Solubility: Alkanes are insoluble in water because they are nonpolar substances. However, they are soluble in nonpolar solvents such as hexane.
  • Density: Alkanes are less dense than water. As a result, when they mix with water, they form a separate layer on top.

Chemical properties

  • Combustion: Alkanes readily undergo combustion reactions, forming carbon dioxide, water and releasing energy. These reactions are the basis for using alkanes as fuel.
  • Substitution reactions: Although alkenes are generally unreactive, they can undergo substitution reactions under specific conditions. For example, when alkenes are exposed to halogens such as chlorine or bromine in the presence of light, a hydrogen atom is replaced by a halogen atom.

Reactions of alkenes

Despite being relatively inert, alkenes engage in some important chemical reactions:

Combustion

The combustion of alkanes is an exothermic reaction that releases heat. The general equation for the combustion of alkanes is as follows:

C n H 2n+2 + (1.5n + 0.5)O 2 → nCO 2 + (n+1)H 2 O

For example, the complete combustion of methane:

CH 4 + 2O 2 → CO 2 + 2H 2 O

Halogenation

One of the few reactions in which alkanes participate is halogenation, where they react with halogens. For methane, the reaction with chlorine can be represented as:

CH 4 + Cl 2 → CH 3 Cl + HCl

This reaction proceeds via a free radical mechanism initiated by UV light.

Uses of alkenes

Alkenes have many practical applications:

  • Fuel: Alkanes are widely used as fuel in cars, jets, and homes. Propane, butane, and natural gas are common examples.
  • Lubricants: Higher alkanes are used as lubricants and as an additive to improve the viscosity of motor oils.
  • Solvent: Some types of alkenes are used as solvents for reactions and extractions in the chemical industry.

Conclusion

Alkanes are the foundation of organic chemistry. They are the simplest hydrocarbons, characterized by a single carbon-carbon bond. Their relatively inertness makes them stable and useful as fuels and chemical feedstocks. Understanding alkanes, their properties, and reactions is the foundational knowledge needed to delve deeper into more complex organic compounds. Whether it's a small molecule like methane or larger chains used in a variety of applications, alkanes are an integral part of both the natural world and industrial processes. This text serves as an introduction to the fascinating world of alkanes in organic chemistry.


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Hydrocarbons

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