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 → Solutions and Mixtures ↓


Henry's Law and Raoult's Law


The study of solutions and mixtures in chemistry is fundamental to understanding how substances interact with one another. Two important concepts essential in the study of solutions are Henry's law and Raoult's law. These laws describe how the pressure of gases in a solution and the vapor pressure of the solvent are affected by the presence of the solute. Let's look at each law in detail.

Henry's law

Henry's law deals with the solubility of gases in liquids. It states that the amount of a gas dissolved in a liquid is directly proportional to the partial pressure of that gas above the liquid at a constant temperature. The mathematical representation of Henry's law is given by this equation:

C = k H * P

Where:

  • C is the concentration of the dissolved gas.
  • k H is the Henry's law constant, which varies for different gases and solvents.
  • P is the partial pressure of the gas above the liquid.

Example of Henry's law

Consider carbonated beverages such as soda. In a sealed soda bottle, carbon dioxide (CO 2 ) is dissolved in the liquid. The pressure is high inside the bottle, which keeps the CO 2 dissolved. When you open the bottle, the pressure immediately drops, and the CO 2 begins to escape as bubbles because the gas is less soluble at the new lower pressure.

liquid CO2 in the gas CO 2 dissolved

Here, the gas molecules are in equilibrium between the liquid and the gas phase above it, and Henry's law helps us understand this equilibrium.

Raoult's law

Raoult's law applies mainly to ideal solutions, which are solutions where the interactions between different molecules are similar to the interactions between identical molecules. It states that the partial vapor pressure of each component in a solution is equal to the vapor pressure of the pure component multiplied by its mole fraction in the solution. It can be expressed mathematically as:

P solution = X A * P A 0 + X B * P B 0

Where:

  • P solution is the total vapor pressure of the solution.
  • X A and X B are the mole fractions of components A and B, respectively.
  • P A 0 and P B 0 are the vapor pressures of the pure components A and B.

Example of Raoult's law

Consider a binary solution composed of volatile liquids such as benzene and toluene. In this case, Raoult's law can predict how each component contributes to the total vapor pressure of the solution. If the vapor pressure of benzene is 100 mmHg and the vapor pressure of toluene is 40 mmHg, the partial pressures will vary depending on their mole fractions in the mixture.

Solution: benzene + toluene benzene vapor Toluene Vapor

In this solution, Raoult's law helps us determine how much each liquid contributes to the total pressure and what the composition of the vapor phase above the liquid will be.

Relevance and applications

Understanding these laws is important for many applications and industries. For example, distillation processes rely heavily on Raoult's law to separate components based on differences in volatility. Similarly, Henry's law is important for understanding phenomena such as gas exchange in the atmosphere or aqueous environments.

Industrial applications

1. **Beverages:** Henry's Law is directly applicable in carbonated beverage production. The solubility of CO2 decreases as the pressure decreases, leading to the bubbling that occurs when opening a beverage.

2. **Environmental Science:** Henry's Law is essential for predicting how gases will dissolve in the oceans and affect marine life, especially in terms of oxygen availability and pH levels.

Research and development

Research in chemistry often involves investigating how different substances interact under different conditions, and Henry's and Raoult's laws are used to understand and predict the behaviour of solutions.

Conceptual clarification

While both laws are fundamental, it is important to note their limitations. Henry's law applies strictly only when the gas does not react chemically with the solvent. Raoult's law applies most accurately to ideal solutions; real solutions may exhibit deviations due to differences in molecular interactions. Such deviations can be described using activity coefficients or through modified versions of the laws.

Conclusion

In short, Henry's Law and Raoult's Law are the backbone of solution chemistry, helping scientists and engineers predict the behavior of mixtures and develop efficient methods for industrial applications. As you delve deeper into chemistry, you will find that these laws are integral to understanding and manipulating chemical processes.


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Henry's Law and Raoult's Law

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