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 → States of matter ↓


Gas Laws


The study of gases is an important part of understanding the physical chemistry of substances. Gases are one of the three main states of matter, along with liquids and solids. Unlike solids and liquids, gases have no definite shape or volume. They expand to fill the container they are in. This characteristic makes them unique in many chemical contexts. The behavior of gases is described by a number of rules that relate the variables of pressure, volume, temperature, and the number of gas molecules (or moles). These laws are collectively known as the gas laws.

Concept of ideal gas

Before going into the specific gas laws, it is necessary to understand the concept of an ideal gas. An ideal gas is a theoretical gas composed of many randomly moving point particles that interact only when they collide. Real gases approximate the behavior of an ideal gas under a wide range of conditions, but the behavior of an ideal gas is an approximation that becomes inaccurate at high pressures and low temperatures where deviations from the ideal gas law occur due to interactions between gas molecules.

The ideal gas law is represented by the following equation:

PV = nRT

Where:

  • P is the pressure of the gas
  • V is the volume of the gas
  • n is the number of moles of gas
  • R is the ideal gas constant
  • T is the temperature in Kelvin

Individual gas laws

Boyle's law

Boyle's law relates the pressure of a gas to its volume at a constant temperature. It states that the pressure of a given mass of gas is inversely proportional to its volume, provided the temperature remains constant. In mathematical terms, it can be expressed as:

P₁V₁ = P₂V₂

Here, P₁ and V₁ are the initial pressure and volume, while P₂ and V₂ are the final pressure and volume. Let’s visualize this:

P₁, V₁ P₂ > P₁, V₂ < V₁

As you can see, when the volume decreases, the pressure increases while the temperature remains constant. Real-life examples of Boyle's law include breathing. When you breathe in, the diaphragm expands the lungs, increasing their volume and lowering the pressure inside them relative to the outside air pressure, allowing air to flow in.

Charles's law

Charles's law describes how gases expand when heated. It states that the volume of a gas is directly proportional to its temperature in Kelvin, provided the pressure remains constant. Mathematically, Charles's law is represented as:

V₁/T₁ = V₂/T₂

Here, V₁ and T₁ are the initial volume and temperature, while V₂ and T₂ are the final volume and temperature. A simple visual representation can be made:

V₁, T₁ V₂ > V₁, T₂ > T₁

In this illustration, as the temperature increases, so does the volume. A common example is a hot air balloon, where heating the air inside the balloon causes it to expand, which decreases the density and causes the balloon to rise because the lower density air is lighter than the colder air outside.

Avogadro's law

Avogadro's law states that the volume of a gas is directly proportional to the number of moles of gas, provided the temperature and pressure are constant. This law can be written as:

V₁/n₁ = V₂/n₂

Illustration of Avogadro's law:

V₁, n₁ V₂ > V₁, n₂ > n₁

In this example, as the number of gas particles (or moles) increases, so does the volume. This principle is often experienced when blowing up a balloon; the more air (moles of gas) you add, the bigger the balloon becomes.

Gay-Lussac's law

Gay-Lussac's law states that the pressure of a gas is directly proportional to its temperature in Kelvin, assuming its volume remains constant. It can be expressed as:

P₁/T₁ = P₂/T₂

Conceptual Visualization:

P₁, T₁ P₂ > P₁, T₂ > T₁

In this scenario, as the temperature of the gas increases, the pressure also increases if there is no change in volume. Think of a pressure cooker: as the temperature of the water inside increases, the steam generated builds up pressure inside the sealed vessel.

Combined gas law

The combined gas law unifies the laws of Boyle, Charles, and Gay-Lussac. This is helpful when more than one variable needs to be solved for. The combined law is formulated as follows:

(P₁V₁)/T₁ = (P₂V₂)/T₂

This equation shows how pressure, volume, and temperature are interrelated for a given amount of gas. If you know how two of these variables change, you can calculate how the third variable will change.

Understanding and application

Understanding these gas laws requires the practice and application of their principles. Solving problems in chemistry often involves identifying which law applies and using it to find unknown variables. It is also important to understand under what conditions these laws are true, especially when they predict real gas behavior under ideal conditions.

Modern real-world applications of the gas laws involve a variety of industrial and scientific processes. For example, in airbags, rapid chemical reactions produce nitrogen gas to inflate the bag in a crash. Calculations using the gas laws ensure that airbags inflate adequately and safely.

Likewise, engineers apply these rules when designing pressurized environments such as airplane cabins or underwater habitats. Gas laws are important for ensuring the safety and comfort of occupants by maintaining proper pressure and oxygen levels.

Limitations of gas laws

While the gas laws provide a fundamental understanding of gas behavior, they have their limitations. The ideal gas laws assume that gases are composed of particles in continuous random motion, with no intermolecular forces except during collisions. However, in reality, gases exhibit attractive and repulsive forces, which can affect their behavior under certain conditions. Deviations from ideal behavior occur particularly at high pressures and low temperatures.

More complex equations of state, such as the van der Waals equation, attempt to correct for these deviations by taking into account the volume occupied by the gas molecules and the forces between them. This is particularly important in chemical industries where precision and safety are paramount.

Conclusion

The gas laws are fundamental to the field of chemistry, providing essential information about the behavior of gases under various conditions. By understanding these laws, students and professionals can predict and manipulate gas behavior in practical applications. Even though we acknowledge their limitations, the gas laws remain fundamental concepts in both theoretical and applied chemistry, linking basic principles with complex real-world phenomena.


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Gas Laws

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