Graduate
  1. Physical Chemistry  1.1. Thermodynamics  1.1.1. Laws of Thermodynamics  1.1.2. Enthalpy and heat capacity  1.1.3. Entropy and free energy  1.1.4. Phase Equilibrium  1.1.5. Statistical thermodynamics  1.1.6. Thermodynamic cycle  1.1.7. Fugacity and activity  1.2. Quantum Chemistry  1.2.1. Principles of quantum mechanics  1.2.2. Schrödinger Equation  1.2.3. Particle in a box  1.2.4. Operators and eigenvalues  1.2.5. Atomic orbitals  1.2.6. Molecular orbital theory  1.2.7. Perturbation theory  1.2.8. Valence bond theory  1.2.9. Hybridization and chemical bonding  1.3. Chemical kinetics  1.3.1. Rate laws and reaction mechanisms  1.3.2. Collision theory  1.3.3. Transition state theory  1.3.4. Enzyme Kinetics  1.3.5. Chain Reactions and Polymerization  1.3.6. Photochemical reactions  1.4. Statistical mechanics  1.4.1. Partitioning functions  1.4.2. Molecular distribution functions  1.4.3. Boltzmann Distribution  1.4.4. Bose–Einstein and Fermi–Dirac statistics  1.5. Spectroscopy  1.5.1. Rotational Spectroscopy  1.5.2. Vibrational Spectroscopy  1.5.3. Electronic Spectroscopy  1.5.4. Nuclear magnetic resonance spectroscopy  1.5.5. Mass Spectrometry  1.5.6. Raman Spectroscopy  1.5.7. Electron paramagnetic resonance spectroscopy  1.6. Surface and colloidal chemistry  1.6.1. Absorption isotherm  1.6.2. Surface tension and wetting  1.6.3. Colloidal Stability  1.6.4. Catalysis  1.6.5. Emulsions and micelles  1.7. Electrochemistry  1.7.1. Nernst equation  1.7.2. Electrochemical cells  1.7.3. Conductivity and mobility  1.7.4. War  1.7.5. Fuel cells  1.7.6. Electrolysis
  2. Organic chemistry  2.1. Reaction mechanism  2.1.1. Nucleophilic substitution reactions  2.1.2. Electrophilic addition reactions  2.1.3. Elimination reactions  2.1.4. Rearrangement reactions  2.1.5. Radical reactions  2.1.6. Pericyclic reactions  2.2. Spectroscopy and structural determination  2.2.1. UV-Vis Spectroscopy  2.2.2. IR Spectroscopy  2.2.3. NMR Spectroscopy  2.2.4. Mass Spectrometry  2.2.5. X-ray crystallography  2.3. Stereoscopic  2.3.1. Chirality and optical activity  2.3.2. Structural analysis  2.3.3. Geometrical isomerism  2.3.4. Dynamic Stereochemistry  2.4. Organometallic Chemistry  2.4.1. Organolithium and organomagnesium reagents  2.4.2. Palladium-catalyzed cross-coupling reactions  2.4.3. Transition Metal Complexes  2.4.4. Metal–carbon bond  2.5. Polymer chemistry  2.5.1. Polymerization mechanism  2.5.2. Polymer Characteristics  2.5.3. Biodegradable Polymer  2.5.4. Conducting polymer  2.5.5. Supramolecular polymers  2.6. Medicinal chemistry  2.6.1. Drug design and development  2.6.2. Pharmacokinetics and pharmacodynamics  2.6.3. Structure-activity relationships  2.6.4. Molecular docking and drug screening
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Crystal field theory  3.1.2. Ligand field theory  3.1.3. Spectrochemical Series  3.1.4. Chelation and stability  3.2. Organometallic Chemistry  3.2.1. Metal Carbonyls  3.2.2. Catalysis by organometallic complexes  3.2.3. Metallocenes  3.3. Bio-inorganic chemistry  3.3.1. Metalloproteins and enzymes  3.3.2. The role of metals in biological systems  3.3.3. Metal ion transport and storage  3.4. Solid state chemistry  3.4.1. Crystal Structures  3.4.2. Band theory of solids  3.4.3. Superconductors  3.4.4. Defects in the crystal  3.5. Lanthanides and Actinides  3.5.1. Electronic configuration in lanthanides and actinides  3.5.2. Coordination chemistry of the lanthanides  3.5.3. Magnetic Properties of Lanthanides
  4. Analytical chemistry  4.1. Chromatography  4.1.1. Gas Chromatography  4.1.2. High Performance Liquid Chromatography  4.1.3. Thin Layer Chromatography  4.2. Spectroscopic Techniques  4.2.1. Atomic absorption spectroscopy  4.2.2. X-ray diffraction  4.2.3. Inductively coupled plasma spectroscopy  4.3. Electroanalytical methods  4.3.1. Potentiometry  4.3.2. Voltammetry  4.3.3. Colometry  4.4. Mass Spectrometry  4.4.1. Ionization Technique  4.4.2. Fragmentation Pattern  4.5. Chemometrics  4.5.1. Multidisciplinary analysis  4.5.2. Machine Learning in Chemistry
  5. Theoretical and Computational Chemistry  5.1. Molecular dynamics simulation  5.1.1. Force fields and energy minimization  5.1.2. Monte Carlo simulation in molecular dynamics simulations  5.2. Quantum chemical methods  5.2.1. Hartree–Fock theory  5.2.2. Density functional theory  5.2.3. Semi-empirical methods  5.3. Computational drug design  5.3.1. Molecular Docking  5.3.2. QSAR Modeling  5.3.3. Virtual Screening
  6. Biochemistry  6.1. Enzyme Kinetics  6.1.1. Michaelis-Menten kinetics  6.1.2. Inhibition mechanism  6.2. Metabolism and Bioenergetics  6.2.1. Glycolysis  6.2.2. Citric acid cycle  6.2.3. Electron transport chain  6.3. molecular Biology  6.3.1. DNA replication and repair  6.3.2. Protein synthesis  6.3.3. Gene Regulation  6.4. Structural Biochemistry  6.4.1. Protein Folding  6.4.2. Membrane Biophysics
  7. Environmental Chemistry  7.1. Atmospheric Chemistry  7.1.1. Greenhouse gases  7.1.2. Ozone depletion  7.1.3. Air pollutants and smoke  7.2. Water Chemistry  7.2.1. pH and water hardness  7.2.2. Wastewater treatment  7.2.3. Aquatic Chemistry  7.3. Soil Chemistry  7.3.1. Heavy metal contamination  7.3.2. Soil pH and Buffering  7.3.3. Nutrient cycling  7.4. Toxicology and Chemical Safety  7.4.1. Toxicokinetics  7.4.2. Risk Assessment in Toxicology and Chemical Safety in Environmental Chemistry  7.4.3. Effects of pollutants on the environment

Graduate → Analytical chemistry → Electroanalytical methods ↓


Potentiometry


Introduction

Potentiometry is a fundamental electroanalytical method widely used in analytical chemistry for the determination of ion concentrations in a solution. The method is based on the measurement of the potential difference between two electrodes without drawing any significant current. Potentiometry is valued for its sensitivity, simplicity, and applicability to a wide range of chemical systems.

Basic concepts

The basic principle behind potentiometry is the Nernst equation, which relates the potential difference across the electrodes to the concentration of ions in solution. The general form of the Nernst equation is:

E = E 0 + (RT/nF) * ln([Ox]/[Red])

Where:

  • E is the electrode potential.
  • E0 is the standard electrode potential.
  • R is the universal gas constant (8.314 J/(mol K)).
  • T is the temperature in Kelvin.
  • n is the number of electron moles transferred in the half-reaction.
  • F is the Faraday constant (96485 C/mol).
  • [Ox] and [Red] are the activities of the oxidized and reduced species, respectively.

Key components of potentiometry

Reference electrode

The reference electrode is important in potentiometry because it provides a stable potential against which the potential of the measuring electrode can be compared. Some common reference electrodes include:

  • Silver/Silver Chloride Electrode ((text{Ag/AgCl}))
  • Calomel Electrode ((text{Hg/Hg}_2text{Cl}_2}))

Indicator or working electrode

The indicator electrode, also called the working electrode, is sensitive to the activity or concentration of the analyte in solution. Common types are:

  • Metal electrodes such as platinum or gold for redox reactions.
  • Ion-selective electrodes (e.g., glass pH electrodes).

Understanding electrode cells in potentiometry

In potentiometry, an electrochemical cell consists of two half-cells connected to an external measurement circuit and an electrolyte solution. The potential difference between the two electrodes is established due to the ionic species in the solution. A simplified cell diagram can be represented as follows:

reference electrode Indicator Electrode Electrolyte Solution

Applications of potentiometry

Potentiometry is useful in a variety of fields. Some of the major applications include:

pH measurements

One of the most common uses of glass electrode potentiometry is for pH measurement. The potential difference is measured and correlated with the activity (concentration) of hydrogen ions. The relationship for this is expressed in the Nernst equation:

E = E 0 - (RT/F) * ln([H + ])

Ion-selective electrode (ISE)

ISE can be used to measure specific ions in solution, such as sodium, potassium or fluoride ions. These electrodes use membranes that provide selectivity to specific ions.

Titration

Potentiometry is often used in titrations (potentiometric titrations), where the potential change is measured as a function of the titrant volume. The end point can be accurately determined from the potential change.

Advantages and disadvantages

Potentiometry has many benefits, including:

  • Non-destructive: does not consume the analyzer for measurement.
  • Wide range of applications: from simple pH measurements to complex ion analysis.
  • High sensitivity: able to detect low concentrations of ions.
  • Simple equipment: Relatively simple and cost-effective equipment.

However, it also has its limitations:

  • Dependence on proper electrode functioning: Accuracy depends on the proper functioning of the electrodes.
  • Calibration requirements: Frequent calibration is required to ensure accuracy.
  • Electrode fouling: Can be affected by fouling, which affects performance.

Principles and mechanisms

The main principle involves the establishment of a known chemical environment using reference and working electrodes. The reference electrode maintains a constant potential against which the changing potential of the indicator electrode reflects the changes caused by the analyte.

Potential measurement

The measured cell potential is the difference between the potentials of the reference and indicator electrodes:

E cell = E indicator - E reference

Practical considerations

Calibration

Calibrating an ion-selective electrode (ISE) is important for accurate measurements. Standard solutions of known concentrations are used to create a calibration curve.

Maintenance

Regular maintenance of the electrodes, including their cleaning and proper storage, is essential to prevent calibration errors and ensure accuracy.

Case study: pH measurement in the laboratory

Consider a case where the laboratory needs to accurately measure the pH of a solution. Using a glass electrode, potentiometry can be used to accomplish this task. Here is a step-by-step method:

  1. Calibrate the pH meter with a standard buffer solution.
  2. Dip the electrode into the test solution.
  3. Allow the meter to settle and record the static potential value.
  4. Convert this potential to pH using the calibration curve established from the first step.

Future directions in potentiometry

Technological advances continue to enhance the capabilities of potentiometry. Future directions may be the following:

  • Miniaturization: Creation of compact, portable potentiometric instruments for on-site analysis.
  • Integration with digital technologies: Improved data processing and integration with software for analysis and storage.
  • Improved selectivity: development of more selective electrodes for complex matrices.

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Potentiometry

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