Equilibrium

Definition

Equilibrium occurs when the rate of the forward process equals the rate of the backward process, resulting in no net change in the system.

Explanation

In a physical equilibrium, only the state of matter changes (e.g., ice melting and water freezing). In a chemical equilibrium, the chemical identity changes as reactants turn into products and vice versa.

Key Points

• Occurs only in a closed system.
• Concentrations of reactants and products remain constant.
• The system is stable but remains active.

Examples

• Physical: A sealed bottle of soda (CO₂ gas vs dissolved CO₂).
• Chemical: H₂(g) + I₂(g) ⇌ 2HI(g).

Real-Life Use: Hemoglobin & Oxygen

In the lungs (high O₂), oxygen binds to hemoglobin. In muscles (low O₂), the equilibrium shifts, and hemoglobin releases oxygen. This balance ensures oxygen reaches all tissues.

Definition

Dynamic equilibrium means the reaction hasn't stopped; instead, the forward and backward reactions are occurring at the exact same speed.

Simple Explanation

Imagine two people on a treadmill: one pushes the belt forward and the other pulls it back with equal force. The belt doesn't move (macroscopically), but both people are working hard (microscopically).

Key Points

• Rate of Forward Reaction = Rate of Backward Reaction.
• Macroscopic properties (color, pressure, concentration) stay constant.
• Microscopic activity continues unabated.

Examples

• Water evaporating from a closed beaker at the same rate it condenses.
• Saturation of sugar in tea where dissolving and crystallization rates are equal.

Real-Life Use: Ozone Layer

O₃ is constantly broken down by UV rays and simultaneously reformed. This dynamic equilibrium keeps the ozone layer stable enough to protect Earth.

Definition

For a reversible reaction at constant temperature, the ratio of the product concentrations to reactant concentrations (each raised to their stoichiometric coefficients) is constant.

The "Scorecard" Logic

The Equilibrium Constant (K) tells us the "preference" of a reaction:

• Large K: Reaction prefers making products.
• Small K: Reaction prefers staying as reactants.

Key Types of Constants

• K_c: Calculated using molar concentrations.
• K_p: Calculated using partial pressures (for gases).
• Temperature: K only changes if temperature changes.

Numerical Illustration

For the reaction: N₂ + 3H₂ ⇌ 2NH₃

Kc = [NH3]2 / ([N2][H2]3)

Real-Life Use: Industrial Ammonia

In the Haber Process, engineers use K to predict yields and decide the exact pressure/temperature to make the process profitable.

Definition

If a system at equilibrium is disturbed (concentration, pressure, or temperature), the system shifts its position to counteract the disturbance.

The "Stubborn Teenager" Analogy

If you push the equilibrium one way, it will push back in the opposite direction to regain its balance.

Key Shifts

• Concentration: Add reactant $\rightarrow$ shift toward products.
• Pressure: Increase pressure $\rightarrow$ shift toward fewer gas moles.
• Temperature: Increase heat $\rightarrow$ shift toward endothermic direction.

Examples

• Adding N₂ to Haber process pushes the reaction to create more NH₃.
• Increasing pressure in N₂O₄ ⇌ 2NO₂ shifts it toward N₂O₄ (fewer moles).

Real-Life Use: Carbonated Drinks

High pressure in a sealed bottle keeps CO₂ dissolved. When you open the cap, pressure drops; the system shifts to "replace" that pressure by turning dissolved CO₂ back into gas (the "hiss" and bubbles).

Definition

The balance between unionized molecules and ions in a solution.

Electrolyte Types

• Strong Electrolytes: 100% ionization (e.g., HCl, NaOH).
• Weak Electrolytes: Partial ionization (e.g., CH₃COOH).

Key Concept: Degree of Ionization ($\alpha$)

The fraction of total molecules that actually dissociate into ions.

Examples

• Strong: HCl $\rightarrow$ H⁺ + Cl⁻
• Weak: CH₃COOH ⇌ CH₃COO⁻ + H⁺

Real-Life Use: Stomach Acid

Stomach HCl is a strong electrolyte. Because it ionizes completely, it provides high H⁺ concentration for breaking down proteins and killing bacteria.

Definitions

• pH Scale: Logarithmic scale specifying acidity/basicity.
• Buffer Solution: A solution that resists pH changes when small amounts of acid/base are added.

The "Chemical Sponge" Concept

A buffer acts like a sponge, soaking up extra H⁺ or OH⁻ ions to keep the pH stable.

Key Formulas & Values

• $\text{pH} = -\log[\text{H}^+]$
• pH 7 = Neutral | < 7 = Acidic | > 7 = Basic

Real-Life Use: Human Blood

Blood must stay between pH 7.35 and 7.45. The Bicarbonate Buffer System (H₂CO₃ / HCO₃⁻) neutralizes acid spikes, preventing fatal pH shifts.

Definitions

• $K_{sp}$: The equilibrium constant for a solid dissolving in water.
• Common Ion Effect: Decrease in solubility of an ionic compound when a salt containing a common ion is added.

Simple Explanation

The Common Ion Effect is like trying to enter a crowded room; if the room is already full of similar ions, the new salt can't dissolve and stays solid.

Precipitation Condition

If Ionic Product > $K_{sp}$, a precipitate forms.

Example

Adding NaCl to AgCl solution reduces AgCl solubility because of the common Cl⁻ ion.

Real-Life Use: Kidney Stones

When calcium or oxalate ions in urine exceed the $K_{sp}$ of Calcium Oxalate, the equilibrium shifts toward the solid, forming crystals (stones).

1. At equilibrium, the rate of forward and backward reactions are:
   Ans: (b) Equal
2. If $K_c$ for a reaction is very large, it means:
   Ans: (b) Products are favored
3. Increasing the pressure of a gaseous system shifts the equilibrium toward:
   Ans: (b) The side with fewer moles
4. Which of the following is a weak electrolyte?
   Ans: (c) $\text{CH}_3\text{COOH}$
5. A buffer solution consists of:
   Ans: (b) Weak acid + Its conjugate base
6. If $[\text{H}^+] = 1 \times 10^{-5}\text{ M}$, the pH is:
   Ans: (a) 5
7. The common ion effect generally ___________ the solubility of a salt.
   Ans: (b) Decreases
8. Which factor does NOT change the value of $K$?
   Ans: (d) Both b (Concentration) and c (Pressure)
9. $\text{pH} + \text{pOH}$ equals:
   Ans: (c) 14
10. A solution with pH 3 is how many times more acidic than pH 4?
    Ans: (b) 10 times

1. Q: State Le Chatelier's Principle. What happens if you add a catalyst?
   Ans: If a system at equilibrium is disturbed, the system shifts to counteract that change. A catalyst does NOT shift equilibrium; it only helps reach it faster.
2. Q: Distinguish between $K_p$ and $K_c$.
   Ans: $K_c$ uses molar concentrations ($\text{mol/L}$), while $K_p$ uses partial pressures ($\text{atm}$). Relation: $K_p = K_c(RT)^{\Delta n}$.
3. Q: Why is the pH of blood maintained at a constant value?
   Ans: Blood uses a bicarbonate buffer system (H₂CO₃ / HCO₃⁻) to neutralize acid/base spikes, preventing protein denaturation.
4. Q: What is the Common Ion Effect? Give an example.
   Ans: Suppression of dissociation of a weak electrolyte by adding a strong electrolyte with a common ion. Example: Adding $\text{CH}_3\text{COONa}$ to $\text{CH}_3\text{COOH}$.
5. Q: Define Solubility Product ($K_{sp}$). What is the condition for precipitation?
   Ans: Product of molar concentrations of ions in a saturated solution. Precipitation occurs when $\text{Q}_{sp} > K_{sp}$.
6. Q: Calculate the pH of a $0.01\text{ M HCl}$ solution.
   Ans: $\text{HCl}$ is strong, so $[\text{H}^+] = 10^{-2}\text{ M}$. $\text{pH} = -\log[10^{-2}] = 2$.
7. Q: How does temperature affect an endothermic reaction at equilibrium?
   Ans: Heat acts as a reactant. Increasing temperature shifts equilibrium to the right (towards products) and increases $K$.
8. Q: Difference between strong and weak electrolytes?
   Ans: Strong electrolytes ionize completely ($\approx 100\%$) with high conductivity. Weak electrolytes ionize partially with low conductivity.
9. Q: Explain "Degree of Ionization."
   Ans: Ratio of moles of electrolyte dissociated into ions to the total moles dissolved. Denoted by $\alpha$.
10. Q: Why does adding $\text{NaCl}$ to $\text{AgCl}$ cause precipitation?
    Ans: $\text{NaCl}$ adds $\text{Cl}^-$ (common ion), shifting equilibrium to the left ($\text{Ag}^+ + \text{Cl}^- \rightarrow \text{AgCl(s)}$) per Le Chatelier.

• Le Chatelier's "Rebel" Rule: Add $\rightarrow$ Remove | Remove $\rightarrow$ Add | Heat $\rightarrow$ Cool | Pressure $\rightarrow$ Small side.
• pH Scale: "Acid is Low (0-6), Base is High (8-14)".
• K subscripts: p for Pressure, c for Concentration.