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 → Thermodynamics ↓


Spontaneity of reactions


In the world of chemistry, understanding the spontaneity of reactions is a fundamental concept that helps us predict whether a chemical reaction will occur on its own under certain conditions. The study of this concept falls under the branch of thermodynamics, which deals with the energy changes that accompany chemical and physical processes.

Understanding ease

A spontaneous process is a process that occurs naturally without any external intervention. Think of a ball rolling down a hill; it needs no push to move down, as gravity pulls it down spontaneously. In the field of chemistry, a spontaneous reaction is one that can proceed without any external energy input.

Thermodynamic criterion for spontaneity

To understand spontaneity in the context of thermodynamics, we consider two main factors: enthalpy and entropy. Both play an important role in determining whether a reaction or process is spontaneous or not.

Enthalpy (H)

Enthalpy refers to the heat content of a system at constant pressure. For a reaction to be spontaneous, the enthalpy change, represented as ΔH, can affect its spontaneity. Generally, exothermic reactions (ΔH < 0) release heat and are more likely to be spontaneous. Imagine burning wood; it releases heat, making the process spontaneous.

Entropy (S)

Entropy is a measure of the disorder or randomness in a system. The second law of thermodynamics states that the total entropy of an isolated system can never decrease over time. The change in entropy, ΔS, can also quantify spontaneity. Processes that increase the entropy of the universe are generally spontaneous. For example, melting ice increases the disorder in water and is, thus, spontaneous.

        H2O(s) → H2O(l) (spontaneous because ΔS > 0)

Gibbs free energy (G)

The spontaneity of a reaction is most accurately quantified using the Gibbs free energy or simply Gibbs energy. This thermodynamic potential combines both enthalpy and entropy into one value:

        ΔG = ΔH – TΔS

Here:

  • ΔG is the change in Gibbs free energy.
  • ΔH is the change in enthalpy.
  • ΔS is the change in entropy.
  • T is the absolute temperature in Kelvin.

The sign of ΔG tells us about the spontaneity of the reaction:

  • If ΔG < 0, then the reaction is spontaneous.
  • If ΔG > 0, then the reaction will occur spontaneously.
  • If ΔG = 0, then the system is at equilibrium.

Examples analysing spontaneity

Let us consider some examples to understand the concept of spontaneity better.

Example 1: Combustion of methane

        CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)

In the combustion of methane, ΔH < 0 because it is an exothermic reaction that releases heat. Generally, the entropy of the products is higher than that of the reactants due to the increase in disorder of the gas molecules. Therefore, ΔS > 0.

Overall, at ΔH < 0 and ΔS > 0, ΔG will be negative, indicating a spontaneous reaction.

Example 2: Freezing of water

        H2O(l) → H2O(s)

This process occurs spontaneously at temperatures below 0°C. Here, ΔH < 0 because heat is released to the surroundings. However, ΔS < 0 because the system becomes more organized.

At low temperatures, the effect of ΔH is greater than TΔS, making ΔG < 0. Thus, the reaction is spontaneous under these conditions.

Example 3: Dissolving salt in water

        NaCl(s) → Na+ (aq) + Cl- (aq)

In this process, ΔH can be slightly positive or negative depending on the type of salt. Dissolution increases the randomness, resulting in ΔS > 0.

An increase in entropy usually drives TΔS process, causing ΔG < 0. Therefore, disintegration is usually spontaneous.

Effect of temperature

Temperature plays an important role in determining whether a reaction is spontaneous or not. Since ΔG = ΔH - TΔS, TΔS term becomes more important as the temperature increases.

Consider a reaction with ΔH > 0 and ΔS > 0. At low temperatures, ΔH can dominate, causing ΔG > 0 and the reaction to be nonspontaneous. However, as the temperature increases, TΔS can overtake ΔH, resulting in ΔG < 0 and a spontaneous reaction.

Non-spontaneous processes

Not all chemical processes are spontaneous. Sometimes an external energy source is needed to make a reaction proceed. For example, electrolyzing water into hydrogen and oxygen requires an electric current, because ΔG > 0 for this process.

        2H2O(l) + electrical energy → 2H2(g) + O2(g)

Conclusion

Understanding the spontaneity of chemical reactions is crucial for predicting and using natural processes in chemistry. Through the concepts of enthalpy, entropy, and Gibbs free energy, we can explore and control the conditions under which reactions occur. Recognizing the importance of these parameters allows chemists to innovate in areas ranging from energy production to pharmaceuticals, where the driving forces of reactions directly influence technological progress.


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Enthalpy, entropy and Gibbs free energy
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Spontaneity of reactions

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