Internal Energy

ΔU represents the change in internal energy between reactants and products. A positive ΔU means the system gained energy from its surroundings, while a negative ΔU indicates the system lost energy. ΔU depends only on the initial and final internal energy states, not the path taken during the reaction.

According to the first law of thermodynamics, any change in a system's internal energy equals the sum of heat (q) and work (w) exchanged with surroundings. This relationship is expressed as ΔU = q + w. Energy is conserved; any change in the system must be balanced by an equal and opposite change in its surroundings.

Positive q indicates heat flows into the system from surroundings, while negative q means heat flows out. Positive w means work is done on the system, and negative w means the system does work on surroundings. When both q and w are positive, ΔU increases; when both are negative, ΔU decreases.

A system exchanges energy with surroundings through two mechanisms: heat transfer and work. Energy enters a system when it absorbs heat or when surroundings do work on it, such as bending a wire. Energy leaves when the system releases heat or performs work, like burning rocket fuel that lifts a space shuttle.

Internal energy is a state function because ΔU depends only on the initial and final states of the system, not on how the change occurs. Although absolute values of initial and final internal energy cannot be determined, the change between them can be calculated. This property makes internal energy useful for analyzing chemical processes regardless of reaction pathway.

When carbon dioxide transforms to elemental carbon and oxygen, the products have higher internal energy than the reactant, making ΔU positive. Energy transfers from surroundings into the system, increasing its internal energy. This endothermic process requires energy input to break chemical bonds.

The joule (J) is the SI unit for internal energy, heat, and work. All energy quantities in thermochemistry are expressed in joules or derived units like kilojoules (kJ). This standardized unit allows consistent measurement and comparison of energy changes across different chemical processes.