Here's how an equilibrium reaction responds to the addition of extra reactant:

Le Chatelier's Principle

The behavior of an equilibrium reaction when you add more reactant is explained by Le Chatelier's Principle. It states:

* If a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress.

The Shift

When you add more reactant to an equilibrium reaction, the system is stressed. To relieve this stress, the reaction will shift to the right, meaning it will favor the forward reaction.

* Forward Reaction: The reaction that consumes the reactants to produce products.

Why the Shift Occurs

* Increased Reactant Concentration: Adding more reactant increases its concentration.

* Increased Collision Rate: Higher reactant concentration means more collisions between reactant molecules.

* Increased Rate of Forward Reaction: More collisions lead to a faster rate of the forward reaction, consuming the added reactant.

* Shift in Equilibrium: The system shifts to the right to consume the excess reactant and reach a new equilibrium.

Example

Consider the following reversible reaction:

```

N2(g) + 3H2(g) ⇌ 2NH3(g)

```

If you add more nitrogen gas (N2), the equilibrium will shift to the right, producing more ammonia (NH3) and consuming some of the added nitrogen.

Key Points

* The equilibrium shifts to the right, favoring the forward reaction.

* The system will eventually reach a new equilibrium with a higher concentration of products and a lower concentration of the added reactant.

* The exact shift in equilibrium depends on the specific reaction and the amount of reactant added.

Let me know if you would like a more detailed explanation or have any further questions!