Le Chatelier and Energy in a Reversible Complex-Ion Reaction
A student prepares a pink aqueous solution of cobalt(II) chloride. The equilibrium is [Co(H2O)6]2+(aq) + 4 Cl−(aq) ⇌ [CoCl4]2−(aq) + 6 H2O(l), with the hydrated ion pink and the chloride complex blue. The student divides the solution among three test tubes. Tube 1 is warmed to 60 °C and turns blue. Tube 2 is cooled to 5 °C and stays pink. Tube 3 is left at 25 °C; adding a few drops of concentrated HCl turns it blue, and then adding silver nitrate solution produces a white precipitate and the mixture turns pink again. A classmate says, 'Heating always makes reactions faster, so the blue color appears because the particles collide more often.' The student must decide the sign of ΔH for the forward reaction, predict and explain the effect of adding water to the blue equilibrium mixture, and evaluate the classmate's explanation using Le Chatelier's principle rather than collision rate alone.
1. Determine the sign of ΔH for the forward reaction as written. Justify using the temperature observations.
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The forward reaction produces the blue complex, [CoCl4]2−. When Tube 1 is warmed, the mixture turns blue, so the equilibrium shifts toward the products.
Increasing temperature favors the endothermic direction. Since heating shifts this equilibrium to the right, the forward reaction must be endothermic:
ΔH > 0 for the forward reaction as written.
Cooling Tube 2 does not produce a strong blue color, which is consistent with the reverse reaction being exothermic. Cooling favors the exothermic direction, so the equilibrium shifts left toward the pink hydrated ion.
2. Predict and explain the effect of adding a small amount of water to the blue equilibrium mixture. Then explain why adding AgNO3 turns the mixture pink.
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Adding water: water appears on the product side, and adding it also dilutes the dissolved ions. Both effects favor the reverse reaction. The left side has more dissolved particles (1 complex ion + 4 Cl− ions = 5 particles) than the right side (1 complex ion). Dilution favors the side with more dissolved particles, which is the left side. The equilibrium shifts left, converting blue [CoCl4]2− back to pink [Co(H2O)6]2+. The mixture becomes pinker.
Adding AgNO3: silver ions react with chloride ions to form solid AgCl: Ag+(aq) + Cl−(aq) → AgCl(s)
This removes Cl− from the equilibrium. According to Le Chatelier's principle, the system shifts left to replace some of the removed Cl−. That consumes blue [CoCl4]2− and produces pink [Co(H2O)6]2+, so the mixture turns pink. The white precipitate is AgCl.
3. Evaluate the classmate's statement: 'Heating always makes reactions faster, so the blue color appears because the particles collide more often.'
Read the worked answer
The statement mixes up rate with equilibrium position. Heating does increase collision frequency and usually increases both the forward and reverse reaction rates. But a color change at equilibrium is not explained by 'more collisions' alone; it is explained by a change in the equilibrium constant.
Because the forward reaction is endothermic, raising the temperature increases the rate of the forward reaction more than the rate of the reverse reaction. The system is temporarily out of balance, and the net reaction shifts right until a new equilibrium is reached. At the new equilibrium, there is more blue [CoCl4]2− and less pink [Co(H2O)6]2+.
If the forward reaction were exothermic, heating would still increase collision rates, but the equilibrium would shift left and the color would change in the opposite direction. So the classmate's explanation is incomplete: kinetics explains how fast equilibrium is approached, while ΔH and Le Chatelier's principle explain which direction the equilibrium shifts.