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 ↓


titrimeter


Titrimetry is an important technique among the classical methods of analytical chemistry, widely used in undergraduate chemistry education. This method involves determining the concentration of a solute by adding a solution of known concentration (titrant) to the solute until the reaction is complete. This article will explore the principles of titrimetry, types of titrations, indicators used, and practical examples to understand this essential chemical process.

Basic principles of titrimetry

At its core, titrimetry, also known as volumetric analysis, relies on stoichiometric reactions between the analyte (the substance being measured) and the titrant (a solution of known concentration). The key outcome is the equivalence point — the exact point at which the amount of titrant is sufficient to completely react with the analyte. This is different from the end point, which is the practical point we look for with an indicator.

The basic equation used in titrimetry is:

Ca × Va = Ct × Vt

Where:

  • Ca = concentration of analyte
  • Va = volume of the analyte
  • Ct = concentration of the titrant
  • Vt = volume of titrant

Types of titration

Different types of titrations are determined based on the nature of the chemical reactions taking place in the solution. The main types include acid-base, redox, complexometric, and precipitation titrations.

Acid-base titration

This type is one of the most common titrations, where an acid reacts with a base. The end point can be determined using an indicator such as phenolphthalein, which changes color at different pH levels.

Example: Titration of hydrochloric acid (HCl) with sodium hydroxide (NaOH).

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
Equivalence point

Redox titration

Redox (reduction-oxidation) titrations involve electron transfer between the titrant and the analyte. A common example of a redox titration is the titration of ferrous ions (Fe 2+) with potassium permanganate (KMnO 4).

MnO 4 - + 5Fe 2+ + 8H + → Mn 2+ + 5Fe 3+ + 4H 2 O
color change at end point

Complexometric titration

These involve the formation of a complex between the analyte and the titrant, often using a chelating ligand such as EDTA. These are particularly useful for the detection of metal ions.

Example: Determination of calcium ions using EDTA.

Ca 2+ + EDTA 4- → CaEDTA 2-
Endpoints with indicators

Precipitation titration

These titrations involve the formation of an insoluble precipitate. An example of this is the titration of silver nitrate (AgNO 3) with chloride ions (Cl -).

Ag + + Cl - → AgCl (solid)
Formation of precipitation

Indicator in titrimeter

Indicators are substances that change colour at the end point of a titration. The choice of indicator depends on the type of titration and the strength of the acid or base involved.

Phenolphthalein in acid-base titrations

Phenolphthalein is a commonly used indicator used in strong acid-strong base titrations. It is colorless in acidic solution and pink in alkaline solution.

Starch indicator in iodometric titration

In iodometric titration, starch is used to detect the end point. It forms a blue-black complex with iodine, which disappears when the end point is reached.

Eriochrome Black T in complexometric titrations

Eriochrome Black T is used in complexometric titrations, especially with EDTA. It changes from wine red to blue at the end point.

Practical application and example calculations

Let's look at a practical example that shows how titrimetry is applied to find the concentration of an unknown hydrochloric acid solution using sodium hydroxide:

  1. Prepare a NaOH solution with a known concentration of 0.1 M.
  2. Fill a burette with the NaOH solution.
  3. Measure 25.0 ml of HCl solution and add a few drops of phenolphthalein to it.
  4. Titrate the HCl with NaOH until a pink colour persists.

Suppose, 23.5 mL of NaOH was used. Calculate the concentration of HCl:

Let Ca = concentration of HCl and Va = 25.0 mL Ct = 0.1 M and Vt = 23.5 mL Using the formula, Ca × Va = Ct × Vt Ca × 25.0 = 0.1 × 23.5 Ca = (0.1 × 23.5) / 25.0 Ca = 0.094 M

Conclusion

Titrimetry is an essential analytical technique in chemistry, capable of accurately determining unknown concentrations. It forms the basis of many laboratory analyses, where understanding stoichiometry and chemical reactions is crucial. Mastering titrimetry can greatly enhance a chemist's ability to perform quantitative analysis efficiently and accurately, making it a cornerstone of undergraduate chemistry education.


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