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 → Atomic Structure ↓


Isotopes and their applications


Understanding isotopes

When discussing atomic structure, an essential concept to understand is isotopes. An atom has a nucleus containing protons and neutrons, surrounded by electrons that orbit in different energy levels. The number of protons in the nucleus of an atom defines the chemical element and is known as the atomic number. For example, all atoms that contain six protons are identified as carbon atoms.

Despite having the same number of protons, atoms of the same element may have different numbers of neutrons. These forms are known as isotopes. Thus, isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, leading to different atomic masses.

Visualization of carbon isotopes

Consider the most common isotopes of carbon:

  • Carbon-12: 6 protons and 6 neutrons
  • Carbon-13: 6 protons and 7 neutrons
  • Carbon-14: 6 protons and 8 neutrons
Carbon-126p, 6nCarbon-136p, 7nCarbon-146p, 8n

How are isotopes represented

Isotopes are usually represented in two main ways: by using the element name followed by a hyphen and mass number (e.g., carbon-12) or by using m Z notation, where m is the mass number and Z is the atomic symbol. For example, carbon-12 can also be written as 12 C.

The mass number is the sum of the protons and neutrons in the nucleus. For carbon-12, it is expressed as:

       Mass number (A) = Number of protons (Z) + Number of neutrons (N)
       A = 6 + 6 = 12

Stability and radioactivity

Not all isotopes are stable. The nuclei of some isotopes are unstable and undergo radioactive decay, a process through which they change into other elements. Such isotopes are called radioisotopes. The stability of an isotope depends on its neutron-to-proton ratio. Generally, there is a certain optimum ratio where the force binding the nuclei together is strongest. In unstable isotopes, the nucleus may release particles or energy to reach a more stable state.

Example of radioactive decay: carbon-14

Carbon-14 is an example of a radioactive isotope. It decays over time to form nitrogen-14 via beta decay, where a neutron is converted into a proton and an electron is emitted:

       14 C → 14 N + e − + ν e

Applications of isotopes

Isotopes have wide applications in various fields. They play an important role in medicine, archaeology, agriculture and chemistry, etc.

Medicine

One of its major medical applications is in diagnosis and treatment. Radioisotopes such as technetium-99m are used in imaging to diagnose a variety of conditions. Positron emission tomography (PET) scans use isotopes such as fluorine-18 to observe metabolic processes. In cancer therapy, isotopes such as cobalt-60 emit gamma rays to target and kill cancer cells.

Archaeology and carbon dating

One of the best-known applications of isotopes is carbon dating, which is used to determine the age of ancient organic material. It is based on the decay of carbon-14 into nitrogen-14. By comparing the amount of carbon-14 remaining in a sample to the expected original amount, scientists can estimate how many years have passed since the organism died. The half-life of carbon-14, about 5730 years, is important in this type of dating.

Agriculture

In agriculture, isotopes are used to improve food safety and pest control. Tracers can follow the absorption of fertilizers in plants. In addition, isotopes help study plant growth and photosynthesis. In pest control, the sterile insect technique (SIT) uses radiation to sterilize male insects, thereby controlling population growth.

Chemistry

In chemistry, isotopes play an important role in research and experimentation. Deuterium (heavy hydrogen with one neutron) is used in nuclear magnetic resonance (NMR) spectroscopy because of its distinctive magnetic properties. Isotopic labeling with stable isotopes helps trace chemical pathways and reaction mechanisms.

Division and natural abundance

Isotopes occur in varying natural quantities. For example, carbon-12 makes up about 98.89% of the carbon found in nature, while carbon-13 and carbon-14 make up about 1.11% and trace amounts, respectively.

Isotopic fractionation occurs when processes cause a measurable change in these natural ratios. It occurs in both physical and chemical processes and is important in fields such as geochemistry. For example, lighter isotopes react faster at lower temperatures, causing a smaller change in isotope ratios.

Visualization of natural isotopic abundances

98.89%C-121.11%C-13

Closing thoughts

The study of isotopes extends far beyond basic atomic structure concepts. By understanding isotopes, one learns more about the complexities of atomic interactions and element variations that govern natural processes. Through their applications, isotopes underscore their importance in advancing technology and scientific understanding.

In short, isotopes are the keys that unlock profound exploration possibilities in science, medicine, archaeology and a variety of industries, and reveal the subtle complexity and diversity of nature.


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