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 → Chemical reactions ↓


Thermochemical reactions


In the study of chemistry, reactions are fundamental phenomena that occur when substances interact. Thermochemical reactions are a special category of these reactions where heat exchange plays an important role. Understanding these reactions involves discovering how energy changes occur simultaneously with chemical changes. In this detailed lesson, we will delve deep into the principles, concepts, and examples of thermochemical reactions.

1. Introduction to thermochemical reactions

Thermochemistry is the branch of chemistry that deals with energy changes during chemical reactions. The term "thermochemical" refers to the combination of thermal (heat-related) phenomena and chemical changes. Therefore, a thermochemical reaction is a chemical reaction that involves the absorption or release of heat.

This study is important for understanding how energy is conserved in chemical processes, how to predict the feasibility of reactions, and in practical applications, such as designing industrial processes that use energy efficiently.

2. Basic concepts in thermochemical reactions

2.1 Energy and enthalpy

Energy transformations are at the core of thermochemical reactions. The primary focus is on changes in a property called enthalpy, denoted by H Enthalpy is a measure of the total energy of a system, including its internal energy as well as the energy required to make space for it by displacing its environment.

During a chemical reaction, the enthalpy change (ΔH) indicates the heat absorbed or released:

  • If ΔH > 0, then the reaction is endothermic (heat absorbing).
  • If ΔH < 0, then the reaction is exothermic (heat is released).

2.2 Heat capacity and specific heat

Heat capacity is the amount of heat required to change the temperature of a substance by one degree Celsius. Specific heat is the heat capacity per unit mass. These concepts are important for understanding how substances absorb and release heat during reactions.

2.3 Calorimetry

Calorimetry is the method of measuring the heat exchanged in chemical reactions. A calorimeter is an instrument for this purpose. The data obtained from calorimetry is necessary for calculating enthalpy changes.

3. Enthalpy change and thermochemical equation

A thermochemical equation is a chemical equation that involves an enthalpy change. The enthalpy change is expressed in kilojoules per mole (kJ/mol).

C(s) + O2 (g) → CO2 (g) ΔH = -393.5 kJ/mol

In the above example, the formation of carbon dioxide from carbon and oxygen releases 393.5 kJ of energy per mole, which indicates an exothermic reaction.

4. Types of thermochemical reactions

4.1 Exothermic reactions

Exothermic reactions occur with the release of heat. These reactions are spontaneous because they can increase the disorder by dissipating energy. Common examples include combustion reactions, such as the burning of methane:

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g) ΔH = -890.1 kJ/mol

4.2 Endothermic reactions

Endothermic reactions absorb heat from their surroundings. These reactions are less likely to occur spontaneously. A textbook example of this is the decomposition of calcium carbonate:

CaCO3 (s) → CaO(s) + CO2 (g) ΔH = +178 kJ/mol

5. Applications and implications of thermochemical reactions

5.1 Real-world applications

Thermochemical reactions are important in a variety of industries. For example, in the energy sector, understanding these reactions is crucial for the development of fuels and the improvement of energy storage systems. The production of ammonia via the Haber process is another important example that relies heavily on thermochemistry.

5.2 Environmental considerations

Environmental implications of thermochemical reactions include their role in global warming. Burning fossil fuels, which is an exothermic process, releases large amounts of carbon dioxide, a greenhouse gas.

6. Thermodynamics and thermochemical reactions

6.1 Laws of thermodynamics

Thermochemistry is deeply connected to the laws of thermodynamics, which describe how energy is evolved in physical and chemical processes. The first law, also known as the law of conservation of energy, states that energy can neither be created nor destroyed, which explains why the enthalpy change of a reaction can be measured.

6.2 Hess's law

Hess's law states that the total enthalpy change for a chemical reaction is the same no matter what path is taken, as long as the initial and final conditions are the same. This principle allows chemists to calculate enthalpy changes for reactions where direct measurement is not possible.

7. Visual representation of thermochemical reactions

7.1 Reaction pathways and energy diagrams

Energy diagrams provide a visual representation of energy changes during a reaction. In these diagrams, the y-axis represents energy, and the x-axis symbolizes reaction progress. Exothermic and endothermic reactions exhibit different characteristics:

Reactants Products activation energy exothermic reaction

In an exothermic reaction, the products have a lower energy level than the reactants.

Reactants Products endothermic reaction

In an endothermic reaction, the energy level of the products is higher than that of the reactants.

8. Calculations related to thermochemical reactions

8.1 Standard enthalpy change

The standard enthalpy change of a reaction refers to the enthalpy change when all reactants and products are in their standard states. It is typically tabulated and used as a reference point for calculating other enthalpy changes.

8.2 Bond enthalpy

Bond enthalpy, also known as bond dissociation energy, is the energy required to break one mole of a bond in a gaseous substance. It is used to estimate the enthalpy change of a reaction by considering the bonds broken and formed.

ΔH = Σ(bond enthalpy of bonds broken) - Σ(bond enthalpy of bonds formed)

Conclusion

Thermochemical reactions form the backbone of many chemical processes, providing information about energy flow during reactions. Understanding these reactions enables chemists to manipulate and control chemical processes for a variety of applications, from industrial manufacturing to the creation of sustainable energy sources. This comprehensive approach to learning about thermochemical reactions underscores the interplay between chemistry and energy in shaping our world.


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Thermochemical reactions

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