TERMOQUIMICA Teoría 5 Primer principio termodinámica en reacciones químicas - Energía interna U

Introduction to Thermodynamics and Energy Conservation

Fundamental Laws of Chemical Reactions

• The introduction highlights two fundamental laws governing chemical reactions: the law of conservation of mass and the law of conservation of energy.

• The second law, concerning energy conservation, is derived from the first principle of thermodynamics, stating that energy cannot be created or destroyed but only transferred or transformed.

Closed Systems and Energy Transfer

• In a closed system where 50 joules of energy are lost, there must be an equivalent increase in the surrounding environment by 50 joules due to energy transfer.

• This principle leads to mathematical expressions relating system energy changes to heat (q) and work (W), which were discussed in previous sections.

Mathematical Expression of the First Principle

• The first principle can be mathematically expressed as the change in internal energy (ΔU) being equal to the sum of heat exchanged (q) and work done (W).

• ΔU represents variations in internal energy, emphasizing its relationship with heat and work exchanges between a system and its surroundings.

Understanding Internal Energy

Definition and Contributions

• Internal energy is defined as the total energy contained within a system due to its internal structure.

• For example, a gas consists of molecules in constant motion contributing kinetic energies from vibrations, rotations, and translations.

Kinetic Energy Factors

• The kinetic contributions arise from various particle movements—vibrational, rotational, or translational—indicating that particles are never truly at rest unless at absolute zero temperature.

• As temperature increases, so does molecular movement; thus higher temperatures correlate with greater internal energies due to increased kinetic activity.

Potential Energy Contributions

Types of Potential Energy

• Internal energy also includes potential energies arising from electrostatic interactions among particles—both repulsive forces between like charges (e.g., electrons) and attractive forces between opposite charges (e.g., protons and electrons).

Summary on Internal Energy

• Overall, internal energy encompasses both kinetic contributions from particle motion and potential contributions from inter-particle forces. This comprehensive understanding sets up for practical applications illustrated through examples.

Application Example: Gas Expansion

Illustrative Case Study

Thermodynamics: Internal Energy Variation

Understanding Energy Exchange in Thermodynamics

• The system exchanges heat and work, leading to a variation in internal energy. The calculation involves understanding the correct signs for heat and work involved in the process.

• When a gas expands and performs work of 40 joules, this work is considered negative since it is done by the system. Conversely, if work is done on the system, it would be positive. This sign convention is crucial for accurate calculations.

• The IUPAC sign convention states that when the system does work, it is negative; when work is done on the system, it is positive. This principle applies consistently across thermodynamic processes.

• For heat exchange, if the system releases heat (exothermic), it is assigned a negative value; if it absorbs heat (endothermic), it receives a positive value. This distinction helps clarify energy changes during reactions or processes.

• In this scenario, with -60 joules of heat released and -40 joules of work performed by the gas, the total change in internal energy amounts to -1 joule (ΔU = -60 + (-40)). This illustrates how both factors contribute to internal energy variations.

Application of First Law of Thermodynamics to Chemical Reactions

• The first law can also be applied to chemical reactions where there’s a change in internal energy from reactants to products (ΔU). Reactants have an initial internal energy (U_initial) while products have a final internal energy (U_final). These values are often complex and not directly calculable but can be determined experimentally through changes in free energy.

• The change in internal energy for a reaction can be expressed as ΔU = U_final - U_initial because internal energy depends only on initial and final states—this aligns with its classification as a state function within thermodynamics.

• Similarly, changes in free energy during reactions are calculated as ΔG = G_final - G_initial where G represents Gibbs free energy for products and reactants respectively; these values are more accessible than absolute energies due to experimental methods available for measurement.

Heat Transfer During Chemical Reactions

• The relationship between internal energy variation and exchanged heat/work can be summarized as ΔU = q + W where q represents heat exchanged and W represents work done during the reaction process—this equation forms the basis for analyzing thermodynamic systems effectively.