Chapter 9: Chemical Equilibrium

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Chapter Overview & SLOs

What is chemical equilibrium? Chemical equilibrium is a state in a reversible chemical reaction where the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products remain constant over time.

What is the difference between forward and reverse reactions?

• Forward reaction: Proceeds from left to right (reactants → products). It is fastest at the beginning when reactant concentrations are highest.
• Reverse reaction: Proceeds from right to left (products → reactants). It speeds up as products accumulate in the system.
• At equilibrium, the rates of the forward and reverse reactions are equal.

What are reversible and irreversible reactions?

• Reversible reactions: Reactions that can proceed in both forward and reverse directions. They can establish equilibrium. Symbol: $\rightleftharpoons$.
• Irreversible reactions: Reactions that proceed only in one direction (forward) and go to completion. They never reach equilibrium. Example: Combustion of fuels, burning of paper, rusting of iron.

Examples of reversible reactions:

• $N_2 + 3H_2 \rightleftharpoons 2NH_3$ (Haber process for ammonia synthesis)
• $2BrCl \rightleftharpoons Br_2 + Cl_2$
• $H_2 + I_2 \rightleftharpoons 2HI$ (hydrogen iodide formation)
• $CoCl_2 \cdot 6H_2O \rightleftharpoons CoCl_2 + 6H_2O$ (cobalt chloride hydration equilibrium)

What is dynamic equilibrium? At equilibrium, the forward and reverse reactions continue to occur. They do NOT stop. This is called dynamic equilibrium because there is constant activity at the molecular level, but no net change in concentrations.

• Rate of forward reaction = Rate of reverse reaction
• Concentrations of reactants and products remain constant (not necessarily equal)
• Both reactants and products are present in the equilibrium mixture (e.g., $H_2$, $I_2$, and $HI$ in a sealed container).

What are the conditions for equilibrium? For a system to maintain equilibrium, the following must remain constant:

• Temperature: Constant temperature
• Pressure (for gases): Constant pressure
• Volume (for gases): Constant volume
• Concentrations: Constant concentrations of all reactants and products

What is Le Chatelier's principle? Le Chatelier's principle states that when a system at equilibrium is subjected to a change in conditions (temperature, pressure, or concentration), the equilibrium shifts in the direction that tends to counteract the change.

How does temperature affect equilibrium?

• Exothermic forward reaction ($\Delta H < 0$): Increasing temperature favors the reverse reaction (shifts left), decreasing temperature favors the forward reaction (shifts right).
• Endothermic forward reaction ($\Delta H > 0$): Increasing temperature favors the forward reaction (shifts right), decreasing temperature favors the reverse reaction (shifts left).

Example - Cobalt chloride equilibrium: Hydrated cobalt(II) chloride is pink. Anhydrous cobalt(II) chloride is blue.

• $$CoCl_2 \cdot 6H_2O \rightleftharpoons CoCl_2 + 6H_2O$$
• Heating: Shifts equilibrium to the right (produces blue anhydrous $CoCl_2$).
• Adding water: Shifts equilibrium to the left (produces pink hydrated $CoCl_2 \cdot 6H_2O$).

How does pressure/concentration affect equilibrium?

• Increasing concentration of reactants: Equilibrium shifts right (more products formed).
• Increasing concentration of products: Equilibrium shifts left (more reactants formed).
• Increasing pressure (for gases): Equilibrium shifts toward the side with fewer gas molecules.
• Decreasing pressure (for gases): Equilibrium shifts toward the side with more gas molecules.

How does a catalyst affect equilibrium? A catalyst speeds up both the forward and reverse reactions equally. It helps the system reach equilibrium faster but does NOT change the position of equilibrium or the equilibrium constant.

The Haber Process ($N_2 + 3H_2 \rightleftharpoons 2NH_3$):

• Forward reaction is exothermic ($\Delta H = -92$ kJ/mol).
• Iron catalyst is used to speed up the reaction.
• High pressure favors forward reaction (fewer gas molecules on product side: 4 moles reactants → 2 moles products).
• Optimum temperature is a compromise: lower temperature favors yield (exothermic) but slows reaction rate; higher temperature increases rate but lowers yield.

Key differences between reversible and irreversible reactions:

| Property | Reversible Reactions | Irreversible Reactions | |----------|---------------------|----------------------| | Direction | Both forward and reverse | Only forward | | Equilibrium | Can reach equilibrium | Never reach equilibrium | | Symbol | $\rightleftharpoons$ | $\rightarrow$ | | Completion | Does not go to completion | Goes to completion | | Example | Haber process, cobalt chloride | Combustion, rusting |

These notes are strictly aligned with the Student Learning Outcomes (SLOs) for the FBISE 2026 annual examination.

• How do we distinguish between forward and reverse reactions? Compare their directions (left-to-right reactants → products vs right-to-left products → reactants), their starting rates (forward fastest at beginning, reverse speeds up as products accumulate), and how they transform reactants into products or vice versa.
• What are the requirements for a system to be at equilibrium? Explain the necessity of constant temperature, constant pressure (for gases), constant volume (for gases), and the fact that concentrations of all chemical species must remain constant over time, with forward and reverse reactions occurring at equal rates.
• How do we represent reversible chemical processes? Master writing balanced chemical equations for systems like ammonia synthesis ($N_2 + 3H_2 \rightleftharpoons 2NH_3$), hydrogen iodide formation ($H_2 + I_2 \rightleftharpoons 2HI$), bromine chloride decomposition ($2BrCl \rightleftharpoons Br_2 + Cl_2$), and cobalt chloride hydration ($CoCl_2 \cdot 6H_2O \rightleftharpoons CoCl_2 + 6H_2O$).
• How does temperature influence chemical balance? Analyze how adding or removing heat can shift the position of equilibrium (Le Chatelier's principle), causing visible changes such as the color shift in cobalt chloride solutions (pink hydrated $CoCl_2 \cdot 6H_2O$ ↔ blue anhydrous $CoCl_2$).

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