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 → Analytical Chemistry ↓


Classical methods


Analytical chemistry is a fundamental aspect of chemistry that deals with the separation, identification, and quantification of chemical components. Classical methods, also known as wet chemistry, refer to traditional techniques in analytical chemistry that have been used for years. These methods are often taught at the undergraduate level and are foundational to understanding chemical analysis. Despite the advent of sophisticated equipment, classical methods remain relevant due to their simplicity, low cost, and effectiveness for certain types of analyses.

Gravimetric analysis

Gravimetric analysis is a technique that understands the composition of a compound by measuring its mass. The basic idea is to convert the analyte (the substance being analyzed) into a compound of known composition, isolate it, and weigh it to calculate its quantity in the original analyzed sample.

Principle of gravimetric analysis

The principle of gravimetric analysis involves converting the analyte into an insoluble form. This insoluble form is then filtered, washed, dried, and weighed. For example, to determine the amount of sulfate in a sample, barium sulfate BaSO4 can be precipitated by adding a solution of barium chloride BaCl2. By weighing the precipitated barium sulfate, the amount of sulfate in the sample can be calculated.

BaCl 2 + SO 4 2- → BaSO 4 (precipitate) + 2Cl -

Example calculation

1. Weigh the precipitate of BaSO4.
2. Calculate the moles of BaSO4 using its molecular weight.
3. Determine the moles of SO 4 2- present from the moles of BaSO 4.
4. Convert moles of SO4 2- to grams.

Visual example

1. Dissolve the sample 2. Add precipitating agent 3. Filter and dry the precipitate 4. Weigh the precipitate

Titrimetry (volumetric analysis)

Titrimetry is an analytical method in which the amount of a substance in a sample is determined by reacting it with a standard solution of known concentration, known as the titrant. The amount of titrant needed to complete the reaction provides information about the amount of the analyte in the sample.

Types of titration

  • Acid-base titration: Used to determine the concentration of an acid or base by neutralizing it with a base or acid. An indicator is often used to signal the end point.
  • Redox titrations: This involves the transfer of electrons between two substances; this includes reactions such as the titration of iron with potassium permanganate.
  • Complexometric titration: Used to determine the concentration of metal ions using complexing agents such as EDTA.
  • Precipitation titration: This involves the formation of a precipitate during the reaction, such as the titration of halides with silver nitrate.

Principle of titration

The underlying principle of titration involves a visible or measurable endpoint corresponding to the completion of the reaction between the titrant and the analyte. For acid-base titrations, the endpoint is often represented by a color change.

Titration process

1. Prepare the sample solution.
2. Add a few drops of indicator to the sample.
3. Titrate with the standard titrant until the colour of the indicator changes.
4. Measure the volume of titrant used.

Visual example of an acid-base titration setup

1. Acid in the flask 2. Indicator added 3. Titrant in the burette 4. Observe the color change

Example calculation for an acid-base titration

1. Note the initial and final burette readings to calculate the volume of titrant used.
2. To find the unknown concentration, use the equation M 1 V 1 = M 2 V 2, where
   M1, V1 = concentration and volume of the titrant,
   M 2, V 2 = concentration and volume of the unknown.

Colorimetry

Colorimetry is a technique for determining the concentration of a colored compound in a solution. The method is based on the Beer-Lambert law, which relates the absorption of light to the properties of the substance through which the light is traveling.

Theory of colorimetry

The intensity of the colour produced by the compound is measured according to the Beer-Lambert law:

A = εlc

where A is the absorbance, ε is the molar absorbance, l is the path length, and c is the concentration. By measuring the absorbance, the concentration of unknown samples can be determined.

Colorimetry method

1. Prepare a series of standard solutions of known concentrations.
2. Measure the absorption capacity of each standard.
3. Draw a calibration curve of absorbance versus concentration.
4. Measure the absorbance of the unknown solution.
5. Interpolate the concentration of the unknown from the calibration curve.

Visual example of a calibration curve

Concentration Absorption

Precipitation methods

Precipitation is a common technique used in classical analytical chemistry to separate a substance as a solid from a solution based on its solubility. The precipitated solid is then filtered out of the solution, washed, and weighed.

Theory of precipitation

This process involves the addition of a reagent that forms an insoluble compound with the analyte. The precipitate formed can be measured quantitatively to determine the amount of analyte in the original solution.

Example of precipitation reaction

AgNO 3 + NaCl → AgCl (precipitate) + NaNO 3

Procedure of rainfall analysis

1. Dissolve the sample in a suitable solvent.
2. Add a precipitating agent to form an insoluble compound.
3. Filter the precipitate, wash and dry.
4. Weigh the dried precipitate and calculate the amount of analyte.

Visual example of the precipitation process

1. Mix the solution 2. Formation of precipitate 3. Filter the precipitate 4. Dry and weigh

Calculating rainfall yield

The yield of the precipitated compound can be calculated using the initial concentration and the mass of the precipitate.

Conclusion

Classical methods in analytical chemistry provide a fundamental framework for understanding chemical analysis. These traditional techniques, such as gravimetric analysis, titrimetry, colorimetry, and precipitation methods, are essential learning components in undergraduate chemistry education. They help students understand the principles of chemical measurement and the underlying principles governing quantitative analysis. Despite advances in instrumental techniques, classical methods remain invaluable for their simplicity, accuracy, and cost-effectiveness in specific applications.


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Classical methods

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