Chemical Equilibrium Dynamics: Arrhenius, Brønsted-Lowry, and Le Chatelier

Fundamental Chemical Principles: Acid-Base Theories

Arrhenius Theory: Ionic Dissociation

The Arrhenius theory postulates the existence of positive and negative ions in aqueous solutions of acids, bases, and salts (electrolytes) to explain their electrical conductivity.

Key Definitions in Arrhenius Theory

• Acid: A substance, electrically neutral, that in aqueous solution dissociates into $ ext{H}^+$ ions (protons) and negative ions.
• Base: A substance, electrically neutral, that in aqueous solution dissociates into $ ext{OH}^-$ ions (hydroxide or hydroxyl ions) and positive ions.

Neutralization Reaction (Arrhenius)

According to this theory, the neutralization reaction occurs between an acid and a base, yielding a salt and water.

Brønsted-Lowry Theory: Conjugate Acid-Base Pairs

This theory does not consider acids and bases in isolation, but rather as interrelated entities. Under the Brønsted-Lowry framework, reactions are defined as proton transfer reactions.

Definitions of Brønsted-Lowry Acids and Bases

• Acid: A substance (molecule or ion) capable of donating a proton ($ ext{H}^+$) to another substance.
• Base: A substance (molecule or ion) capable of accepting or capturing a proton ($ ext{H}^+$) from another substance.

The concepts of acid and base are complementary. An acid acts as a proton donor only in the presence of a substance capable of accepting it (the base). Conversely, the base can only accept a proton if it reacts with an acid that transfers it. The neutralization reaction is the transfer of a proton from the acid to the base. We consider this process as an equilibrium in which the substances formed can also transfer a proton between them (forming conjugate pairs).

Le Chatelier's Principle: Shifting Chemical Equilibrium

Le Chatelier's principle states that if a system in equilibrium experiences a change in one of the factors affecting it (such as temperature, pressure, or concentration), the system will evolve by shifting the equilibrium position in the direction that tends to counteract or minimize the variation.

Effects of Changing Reaction Conditions on Equilibrium

• By Increasing the Pressure: An increase in pressure will shift the equilibrium toward the side with fewer moles of gas. In the specific reaction example mentioned (where the number of moles of gaseous products is three and the number of moles of gaseous reactants is two), the equilibrium shifts toward the reactants (to the left).
• By Lowering the Temperature: Assuming the reaction is exothermic (releasing heat), a decrease in temperature shifts the equilibrium in the direction that produces more heat. This causes the equilibrium to shift to the right (toward the products).
• When Introducing a Catalyst: Introducing a catalyst increases the speed of both the direct and reverse reactions equally. It does not affect the final equilibrium state, but rather ensures that the equilibrium is reached faster.
• By Introducing More Amount of NO (Product): If the amount of a product, such as NO, is increased in the reaction system, the balance will shift to compensate for this increase. The system will consume the excess NO by reacting it with chlorine to produce NOCl, meaning the reaction shifts toward the reagents (to the left).