Effect of Temperature on Equilibrium 🔥🧊

Temperature changes shift equilibrium! For example:

  • Reaction: \(\ce{[Co(H2O)6]^{3+} (aq) + 4Cl^{-} (aq) \rightleftharpoons [CoCl4]^{2-} (aq) + 6H2O(l)}\)
    Pink ⇄ Blue
  • At room temperature: mixture is blue 🔵 (more \(\ce{[CoCl4]^{2-}}\)).
  • When cooled: mixture turns pink 💖 (more \(\ce{[Co(H2O)6]^{3+}}\)).

Effect of Catalysts ⚡

Catalysts speed up reactions but don’t change equilibrium:

  • Lower activation energy for both forward/reverse reactions.
  • Example 1: Ammonia synthesis (\(\ce{N2 + 3H2 <=> 2NH3}\))
    Catalyst: Iron (Fe) | Conditions: 500°C, 200 atm.
  • Example 2: Sulphuric acid production (\(\ce{2SO2 + O2 <=> 2SO3}\))
    Catalyst: Pt or \(\ce{V2O5}\) | \(K_e = 1.7 \times 10^{26}\).
  • ⚠️ Catalysts can’t help if \(K\) is very small!

Ionic Equilibrium in Solution 💧⚡

Some solutions conduct electricity → electrolytes!

  • Non-electrolytes: Don’t conduct electricity (e.g., sugar solution).
  • Electrolytes: Conduct electricity when dissolved (e.g., salts, acids, bases):
    • 💪 Strong electrolytes: Fully ionized (e.g., NaCl → 100% \(\ce{Na+}\) + \(\ce{Cl-}\)).
    • 🕊️ Weak electrolytes: Partially ionized (e.g., acetic acid → <5% ions).
      Equilibrium exists between ions + unionized molecules: \(\ce{CH3COOH <=> H+ + CH3COO-}\)

Acids, Bases & Salts 🍋🧼

  • Acids:
    • Taste sour (Latin “acidus” = sour!)
    • Turn blue litmus → red 🔵→🔴
    • Release \(\ce{H2}\) gas with metals.
    • Examples: HCl (gastric juice), citric acid (lemons), acetic acid (vinegar).
  • Bases:
    • Taste bitter, feel soapy.
    • Turn red litmus → blue 🔴→🔵
    • Example: Washing soda (\(\ce{Na2CO3}\)).
  • Salts:
    • Formed when acid + base react.
    • Example: \(\ce{HCl + NaOH -> NaCl + H2O}\)
    • \(\ce{NaCl}\) dissolves in water due to water’s high dielectric constant (80), which weakens ionic bonds → ions move freely 💃.

Key NEET Concepts ⭐

  1. Le Chatelier’s Principle: How temp/pressure shift equilibrium (e.g., Co-complex color change, \(\ce{NH3}\) synthesis).
  2. Strong vs. Weak Electrolytes: Full vs. partial ionization (NaCl vs. \(\ce{CH3COOH}\)).
  3. Catalysts: Speed up reaction but don’t alter equilibrium position or \(K\).
  4. Salt Dissolution: Role of dielectric constant (why \(\ce{NaCl}\) dissociates in water).

Fun Fact 🧪

Michael Faraday (1791–1867) discovered electrolysis laws and classified electrolytes! He was super humble and gave awesome science lectures 🎤.