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 → Physical chemistry → Quantum Chemistry ↓


Atomic and molecular spectroscopy


Introduction

Atomic and molecular spectroscopy are fundamental tools in the fields of quantum chemistry and physical chemistry. These spectroscopic techniques provide information about the structure of atoms and molecules by examining how they interact with electromagnetic radiation. Studying these interactions helps us understand the energy levels within atoms and molecules, ultimately revealing much about their electronic configuration and properties.

Basic principles

Electromagnetic radiation

Spectroscopy involves the study of how electromagnetic radiation, including light, interacts with matter. Light can be thought of as both a wave and a particle (photon), where its properties are represented by its wavelength (λ) or frequency (ν). The energy of a photon can be calculated using Planck's equation:

e = hν = frac{hc}{λ}

Where:

  • E is the energy of the photon
  • h is Planck's constant (6.626 x 10^-34 Js)
  • ν is the frequency of the electromagnetic radiation
  • c is the speed of light (3.00 x 10^8 m/s)
  • λ is the wavelength of electromagnetic radiation

Atomic spectra

Atoms emit and absorb light at specific wavelengths, producing an atomic spectrum. These emissions or absorptions are caused by transitions of electrons between different energy levels or orbitals within an atom. Each element has a unique atomic spectrum, often called its fingerprint. This unique pattern of lines allows us to identify the presence of elements in different substances.

Example: hydrogen atom

The hydrogen atom is the simplest atom and has a characteristic spectrum, known as the Balmer series, which can be observed in the visible region. The wavelength can be calculated using the Rydberg formula:

frac{1}{λ} = R_H (frac{1}{n_1^2} - frac{1}{n_2^2})

Where:

  • λ is the wavelength of the emitted light
  • R_H is the Rydberg constant (1.097 x 10^7 m^-1)
  • n_1 and n_2 are integers representing the initial and final energy levels of the electron
n=3 n=4 n=5

The above figure shows the transition lines corresponding to the electrons falling to n=2 level in the hydrogen atom.

Molecular spectra

Unlike atomic spectra, molecular spectra arise from transitions between different vibrational and rotational energy levels in the molecule. Molecular spectroscopy is divided into different types depending on the part of the electromagnetic spectrum being examined, such as infrared (IR) spectroscopy and ultraviolet-visible (UV-Vis) spectroscopy.

Forms of spectroscopy

Infrared (IR) spectroscopy

IR spectroscopy is a technique that deals with the infrared region of the electromagnetic spectrum, which is light with longer wavelength and lower energy than visible light. It is mainly used to investigate vibrational and rotational transitions in molecules.

When molecules absorb infrared light, the energy causes changes in their vibrational states. Different bonds and functional groups absorb specific frequencies of IR radiation. By measuring these frequencies, chemists can extract information about molecular structure.

Example: IR spectrum of water

Absorption Wave number (cm^-1)

The diagram above shows a simplified representation of the IR spectrum for a water molecule. The peaks correspond to the stretching and bending vibrations of the O-H bonds.

Ultraviolet-visible (UV-Vis) spectroscopy

UV-Vis spectroscopy involves the absorption of ultraviolet or visible light by molecules that result in electronic transitions. This type of spectroscopy is particularly useful for studying conjugated systems and can indicate how much light a substance absorbs at different wavelengths.

UV–Vis spectroscopy can provide information about the electronic structure of molecules, since absorption of light in this range often leads to transitions of electrons in bimolecular orbitals, such as from the HOMO (highest occupied molecular orbital) to the LUMO (lowest unoccupied molecular orbital).

Example: UV-Vis spectrum of benzene

Absorption Wavelength (nm)

The figure shows a typical UV-Vis spectrum for benzene, with the peaks corresponding to electronic transitions between different molecular orbitals.

Nuclear magnetic resonance (NMR) spectroscopy

NMR spectroscopy is a powerful technique used to determine the local environment of hydrogen nuclei within molecules. It uses the magnetic properties of atomic nuclei to provide detailed information about molecular structure, dynamics, reaction states, and chemical environments.

In an NMR experiment, the sample is placed in a strong magnetic field, and when exposed to radiofrequency radiation, the nuclei resonate at a frequency specific to their chemical environment.

Example: NMR spectrum of ethanol

The NMR spectrum for ethanol includes signals from hydrogen in different chemical environments:

  • a singlet for the –OH proton
  • a triad for the -CH3 group
  • A quartet for the -CH2- group adjacent to the -OH group
Chemical Shift (ppm)

Applications of spectroscopy

Spectroscopy has many applications in various fields due to its ability to provide qualitative and quantitative information about the structure and composition of matter. These include:

  • Analytical Chemistry: Identifying unknown substances, determining concentrations.
  • Environmental science: Monitoring of pollution levels.
  • Pharmaceuticals: Analyzing drug composition and ensuring quality control.
  • Materials Science: Properties of materials and nanomaterials.
  • Astrophysics: Determining the structure of stars and galaxies.

Conclusion

Atomic and molecular spectroscopy are indispensable techniques in the field of chemistry for understanding how atoms and molecules interact with light. By studying these interactions, chemists can gain important information about the composition, structure, and dynamics of various substances. With advances in technology and a deeper understanding of quantum mechanics, the field of spectroscopy continues to grow, providing more and more detailed and accurate data that pave the way for new discoveries and applications in science and technology.


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Atomic and molecular spectroscopy

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