1
Carbonated water bottles are kept refrigerated before sale. If such a bottle is accidentally shaken and immediately opened, it releases excessive fizz. Using Le Chatelier’s Principle and Henry’s law, explain why refrigeration reduces fizzing.
▼
Lower temperature increases CO₂ solubility (Henry’s Law), keeping equilibrium shifted toward dissolved CO₂. Shaking increases pressure locally, forming nucleation sites. When opened, sudden pressure drop shifts equilibrium toward gaseous CO₂, causing fizz. Refrigeration slows this shift.
2
In deep-sea diving, breathing mixtures contain helium instead of nitrogen. Relate this to gas equilibrium concepts.
▼
Nitrogen dissolves in blood under high pressure (N₂(g) ⇌ N₂(aq)). Rapid ascent decreases pressure, shifting equilibrium backward, forming bubbles (decompression sickness). Helium’s low solubility reduces such equilibrium shift and bubble formation.
3
In the Haber process, a catalyst and high pressure are used together. Why can’t we just increase pressure indefinitely to obtain maximum ammonia yield?
▼
Increasing pressure favors NH₃ formation, but also increases energy cost and may deactivate the catalyst. Beyond an optimum point, rate improvement is minimal because equilibrium constant (Kₚ) is temperature dependent, not pressure dependent.
4
When a soda can is opened at high altitude, fizzing is less compared to sea level. Explain this using the concept of equilibrium partial pressures.
▼
At high altitude, atmospheric pressure is lower, so equilibrium CO₂(aq) ⇌ CO₂(g) was already shifted slightly right even before opening. Pressure difference on opening is smaller, hence less rapid CO₂ release.
5
Hemoglobin binds both O₂ and CO₂ in blood. During intense exercise, CO₂ level rises. Explain, using equilibrium, why oxygen delivery to muscles increases.
▼
Increased CO₂ shifts Hb + CO₂ ⇌ HbCO₂ equilibrium right, freeing more Hb from HbO₂ ⇌ Hb + O₂ equilibrium. Thus, O₂ release increases (Bohr effect), improving tissue oxygenation.
6
When an aqueous solution of weak acid is diluted, its pH increases. Explain the equilibrium change quantitatively.
▼
For weak acid HA ⇌ H⁺ + A⁻, ( K_a = \frac{[H^+][A^-]}{[HA]} ). Dilution decreases [HA] and [H⁺], but the ratio tends to restore Kₐ. However, [H⁺] ∝ √(Kₐ × c), so lowering concentration (c) reduces [H⁺], increasing pH.
7
In an automobile catalytic converter, harmful NO and CO gases are converted to N₂ and CO₂. Why does equilibrium not revert back under normal exhaust conditions?
▼
Reaction 2NO + 2CO ⇌ N₂ + 2CO₂ is exothermic. In the exhaust, gases move rapidly, so equilibrium isn’t re-established backward. Product gases are removed, continuously shifting equilibrium toward products (Le Chatelier’s principle).
8
Cloud formation in the atmosphere can be viewed as a physical equilibrium. Explain.
▼
Water vapor ⇌ liquid droplets equilibrium exists in air. Cooling (decrease in temperature) shifts equilibrium toward condensation, forming clouds. Heating shifts equilibrium toward evaporation (clear sky).
9
In soft drink bottling, CO₂ is dissolved at high pressure. Why can the amount of dissolved CO₂ not be increased indefinitely even if pressure is extremely high?
▼
At high pressures, deviation from ideal gas behavior occurs, and CO₂ starts forming carbonic acid (chemical equilibrium H₂O + CO₂ ⇌ H₂CO₃). Thus, physical dissolution equilibrium is limited by chemical conversion.
10
Industrial preparation of sulfur trioxide (SO₃) involves reversible oxidation of SO₂. Why is temperature control critical for maximum yield?
▼
2SO₂ + O₂ ⇌ 2SO₃ (ΔH < 0). Lower temperature favors SO₃ formation but reduces rate. Industrially, moderate temperature (~723 K) and catalyst are used to achieve dynamic balance between rate and yield.
11
In oceans, CO₂ absorption affects marine life. How does equilibrium explain ocean acidification?
▼
Dissolved CO₂ forms H₂CO₃ ⇌ H⁺ + HCO₃⁻. Excess atmospheric CO₂ shifts equilibrium right, increasing [H⁺], lowering pH. Acidic water dissolves CaCO₃ shells (CaCO₃ ⇌ Ca²⁺ + CO₃²⁻).
12
A patient with respiratory acidosis retains CO₂ in blood. Explain how this affects bicarbonate equilibrium and blood pH.
▼
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. Increased CO₂ shifts equilibrium right, increasing H⁺ concentration, lowering blood pH — causing acidosis.
13
In a closed system of N₂O₄ ⇌ 2NO₂, why does increasing temperature darken the color?
▼
The forward reaction (N₂O₄ → 2NO₂) is endothermic. Higher temperature shifts equilibrium toward NO₂ (brown gas), increasing color intensity.
14
Why does calcium carbonate not completely dissolve in acidic rainwater even though acid enhances solubility?
▼
CaCO₃(s) ⇌ Ca²⁺ + CO₃²⁻. H⁺ reacts with CO₃²⁻ forming H₂CO₃ → CO₂ + H₂O, shifting equilibrium right. But as CO₂ accumulates in the confined environment (soil pores), equilibrium shifts back, limiting dissolution.
15
In biological cells, enzyme-substrate interactions often reach equilibrium. Why does equilibrium favor product formation in living systems?
▼
Products are continuously removed or converted in subsequent reactions, shifting equilibrium toward product side — an open equilibrium system driven by metabolic flux.
16
While preparing ammonia in lab using NH₄Cl and Ca(OH)₂, why is the gas collected over dry water or displaced air and not over water?
▼
NH₃(g) dissolves in water forming NH₄OH ⇌ NH₄⁺ + OH⁻. Collecting over water would shift equilibrium toward dissolution, reducing gaseous yield.
17
In metallurgy, CO is used to reduce Fe₂O₃. Why does high temperature sometimes decrease efficiency of reduction?
▼
Fe₂O₃ + 3CO ⇌ 2Fe + 3CO₂ is exothermic. Very high temperature shifts equilibrium left, decreasing reduction efficiency, despite faster kinetics.
18
When ice melts at 0°C, adding salt lowers the freezing point. Explain equilibrium-wise.
▼
Salt decreases water’s chemical potential, shifting equilibrium H₂O(s) ⇌ H₂O(l) toward melting. Thus, equilibrium temperature (freezing point) decreases — basis of de-icing.
19
In an aquarium, adding more fish increases CO₂ concentration. Predict how this affects carbonate equilibrium and water pH.
▼
CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. More CO₂ shifts equilibrium right, increasing H⁺ concentration and lowering pH — making water acidic.
20
In the contact process, why is SO₃ not directly dissolved in water to form H₂SO₄?
▼
SO₃ + H₂O ⇌ H₂SO₄ is highly exothermic and forms mist, shifting equilibrium toward reactants. Instead, SO₃ is absorbed in H₂SO₄ to form oleum, which is then diluted to maintain equilibrium and control reaction.
