ChemistryEquilibriumClass 11

Chemical & ionic equilibrium / pH titration: interactive simulation

Chemical equilibrium is a dynamic state where the rate of the forward reaction equals the rate of the backward reaction. Any perturbation — temperature, pressure or concentration — causes the system to counteract the change according to Le Chatelier’s principle.

Useful for: JEE Main, JEE Advanced, NEET, CBSE Boards · 4 modes in one lab

What you can explore

  1. 1

    Le Chatelier's Box

    Haber process N₂ + 3H₂ ⇌ 2NH₃ with live temperature, pressure and concentration shifts.
  2. 2

    pH Titration & Indicators

    Titration curves with live color-changing indicators and equivalence point readouts.
  3. 3

    Buffer & Henderson Eq.

    Interactive buffer solution comparing pH resistance against unbuffered water.
  4. 4

    K_sp & Common Ion Effect

    AgCl solubility suppression and precipitation with added NaCl.

Try this

Explore full pH titration curves of strong and weak acids with live indicator transition bands, see how conjugate acid-base buffer mixtures resist pH shifts, and test common-ion suppression of solubility products.

  1. 1

    In Le Chatelier’s Box, raise temperature from 400 K to 700 K.

    What you should see

    Ammonia yield drops sharply because the synthesis of NH₃ is exothermic (ΔH = −92 kJ/mol).

  2. 2

    In pH Titration, compare Equivalence Point pH of HCl vs CH₃COOH titrated with NaOH.

    What you should see

    HCl has neutral equivalence at pH 7.00, while CH₃COOH has basic equivalence at pH 8.72 due to acetate hydrolysis.

  3. 3

    In Buffer sim, add 0.05 mol HCl to the buffer vs pure water.

    What you should see

    Buffer pH shifts only by ~0.4 units, while pure water pH plunges from 7.0 to 1.3.

Formulas behind it

Henderson-Hasselbalch equation (Acidic buffer)

pH=pKa+log⁡10[Salt][Acid]\text{pH} = \text{p}K_a + \log_{10} \frac{[\text{Salt}]}{[\text{Acid}]}

Solubility product & precipitation condition

Ksp=[Ay+]x[Bx−]y,Precipitates when Qsp>KspK_{\text{sp}} = [A^{y+}]^x [B^{x-}]^y, \quad \text{Precipitates when } Q_{\text{sp}} > K_{\text{sp}}

Van ’t Hoff temperature dependence of K

ln⁡(K2K1)=ΔH∘R(1T1−1T2)\ln\left(\frac{K_2}{K_1}\right) = \frac{\Delta H^\circ}{R}\left(\frac{1}{T_1} - \frac{1}{T_2}\right)

Salt hydrolysis pH (Weak acid + Strong base)

pH=7+12(pKa+log⁡10C)\text{pH} = 7 + \frac{1}{2}(\text{p}K_a + \log_{10} C)

Common mistakes

  • Assuming catalyst shifts the position of equilibrium. A catalyst speeds up forward and backward rates equally; K_c remains unchanged.
  • Using methyl orange for weak acid vs strong base titration — methyl orange changes color at pH 3.1–4.4, long before the equivalence pH of ~8.7.
  • Forgetting that solubility in presence of a common ion is calculated using K_sp / [common ion], not √(K_sp).
Stuck on a chemical & ionic equilibrium / ph titration question?Snap it and get a whiteboard video solution.

FAQ

Why does adding an inert gas at constant pressure shift equilibrium?

At constant pressure, total volume expands, decreasing partial pressures of all gases; the reaction shifts towards the side with more moles of gas.

What is buffer capacity?

The moles of strong acid or base needed to change the pH of 1 L buffer by 1 unit. It is maximal when [Salt] = [Acid], where pH = pKa.

Can K_sp change with concentration?

No. Like any equilibrium constant, K_sp is temperature-dependent only.

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