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 → Acids and bases ↓


pH and pOH


In general chemistry, understanding the concepts of pH and pOH is fundamental to understanding how acids and bases behave in solution. These concepts help us measure the acidity or alkalinity of a solution, and they are associated with the concentration of hydrogen ions (H +) and hydroxide ions (OH -).

Nature of acids and bases

Before diving into the concepts of pH and pOH, it is important to understand what acids and bases are. Acids are substances that release hydrogen ions (H +) when dissolved in water, while bases release hydroxide ions (OH -). These ions are essential in defining the acidic or alkaline nature of a solution.

Introduction to pH

The term pH means "potential of hydrogen" or "power of hydrogen." It is a numerical scale used to specify the acidity or alkalinity of an aqueous solution. The pH scale runs from 0 to 14, where:

  • pH < 7: The solution is acidic.
  • pH = 7: The solution is neutral.
  • pH > 7: The solution is basic (or alkaline).

The pH value of a solution is calculated using the following formula:

    pH = -log 10 [H +]

Here, [H +] represents the concentration of hydrogen ions in the solution, usually measured in moles per liter (M).

Example of calculating pH

Suppose you have a solution in which the concentration of hydrogen ions is 1 x 10 -3 M To find the pH, you would calculate:

        pH = -log 10 (1 x 10 -3) = 3

Therefore, the pH of this solution is 3, which indicates that it is acidic.

Visualization of pH scale

The pH scale can be visualized as follows:

0 1 2 3 4 5 6 7 8 Acidic Basic neutral

Introduction to pOH

While pH focuses on the concentration of hydrogen ions, pOH is in contrast a measure of the concentration of hydroxide ions in a solution. Like pH, pOH is calculated using the negative logarithm:

    pOH = -log 10 [OH -]

where [OH -] represents the concentration of hydroxide ions in the solution.

Example calculation of pOH

Imagine that you have a solution in which the concentration of hydroxide ions is 1 x 10 -4 M To find the pOH, you would calculate:

        pOH = -log 10 (1 x 10 -4) = 4

Therefore, the pOH of this solution is 4.

Relation between pH and pOH

One of the fundamental relationships in aqueous solutions is the relationship between pH and pOH, which is defined by the ionic product K w of water. At 25°C, K w is always:

    K w = [H +][OH -] = 1 x 10 -14

Taking the negative logarithm of both sides gives an important equation:

    pH + pOH = 14

This equation shows that if you know either pH or pOH, you can easily determine the other. For example, if a solution has a pH of 5, its pOH can be calculated as follows:

    pOH = 14 - pH = 14 - 5 = 9

Working with the pH and pOH scale

To better understand how the pH and pOH scales work, consider a neutral solution such as water at 25°C. In such a solution:

    [H +] = [OH -] = 1 x 10 -7 M
    pH = -log(1 x 10 -7) = 7
    pOH = -log(1 x 10 -7) = 7

Here, it is clear that the solution is neutral as both pH and pOH are equal, and the equation pH + pOH = 14 is true.

Example 1: Acidic solution

Consider a solution with [H +] = 1 x 10 -2 M The pH is:

        pH = -log(1 x 10 -2) = 2

The corresponding pOH is:

        pOH = 14 - pH = 14 - 2 = 12

This solution is acidic because its pH value is less than 7.

Example 2: Basic solution

Consider a solution with [OH -] = 1 x 10 -3 M The pOH is:

        pOH = -log(1 x 10 -3) = 3

The corresponding pH value is:

        pH = 14 - pOH = 14 - 3 = 11

This solution is alkaline because its pH value is more than 7.

Limitations and considerations

When discussing pH and pOH, it is important to acknowledge some limitations and considerations:

  • Temperature dependence: A pKw value of 14 is typical for 25°C. Changes in temperature can affect this value.
  • Non-aqueous solutions: The pH and pOH scale are primarily applied to aqueous solutions and cannot be directly applied to non-aqueous environments.
  • High concentrations: At very high or very low concentrations of H + and OH -, the log scale assumption may need adjustment.

Usefulness of pH and pOH

Understanding pH and pOH is important in a variety of areas:

  • Biology: Enzymatic and physiological activities have optimum pH ranges.
  • Agriculture: Soil pH affects the availability of nutrients for plants.
  • Environmental science: pH changes caused by pollution and acid rain can threaten the health of aquatic ecosystems.
  • Medical science: Blood pH can affect metabolic processes.

Conclusion

In short, pH and pOH are integral concepts in chemistry that help understand the nature of acids and bases. They are directly related through the water dissociation constant K w. Through various examples, we can appreciate the practical implications of these measures in scientific research and everyday life. By mastering these concepts, students can understand the nuances of chemical equilibrium and the behavior of solutions.


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Arrhenius, Brønsted–Lowry, and Lewis definitions
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Acid-base titration

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pH and pOH

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