Frequently Asked Question

Energy changes are fundamental to both physical and chemical processes. Here's a breakdown:
Physical Processes:
Phase Changes: Energy is absorbed (endothermic) during melting, boiling, and sublimation, as molecules transition to states with greater freedom of motion. Conversely, energy is released (exothermic) during freezing, condensation, and deposition.
Dissolving: The process of dissolving can be either endothermic or exothermic depending on the specific solute and solvent.
Heating/Cooling: Adding heat to a substance increases its kinetic energy, raising its temperature. Removing heat decreases kinetic energy and lowers temperature.
Chemical Processes:
Chemical Reactions: Reactions involve the breaking and forming of chemical bonds. Energy is required to break bonds (endothermic), and energy is released when bonds are formed (exothermic).
Enthalpy Change (ΔH): A measure of the heat absorbed or released in a reaction at constant pressure. A positive ΔH indicates an endothermic reaction, while a negative ΔH indicates an exothermic reaction.
Key Concepts:
Exothermic: Releases energy to the surroundings (feels hot)
Endothermic: Absorbs energy from the surroundings (feels cold)
Thermochemistry: The study of energy changes in chemical reactions.
Hess's Law: Allows calculation of enthalpy changes for reactions that cannot be directly measured.
Entropy: A measure of the disorder or randomness of a system. Reactions that increase entropy are favored.
Examples:
Burning wood is exothermic (releases heat and light)
Ice melting is endothermic (absorbs heat from the surroundings)
Photosynthesis is endothermic (plants absorb light energy to create glucose)