Le Chatelier: system at equilibrium responds to oppose applied change. Temperature: endothermic reaction ($\Delta H > 0$): increase T → forward shift → higher yield of products. Decrease T → reverse shift. For exothermic: opposite. Concentration: add reactant → forward shift. Remove product → forward shift (used in industry to drive reactions). Pressure: only matters if $\Delta n_g \neq 0$. Increase P → shifts toward fewer gas moles. Decrease P → shifts toward more gas moles. If $\Delta n_g = 0$ (e.g., H₂+I₂⇌2HI, N₂+O₂⇌2NO): pressure has NO effect on equilibrium composition. Catalyst: no change in position. Only reaches equilibrium faster. Industrial applications: Haber (N₂+3H₂→2NH₃, $\Delta n_g=-2$): high P (more NH₃) but low T not practical (too slow). Compromise: 400-500°C with Fe catalyst. Contact (2SO₂+O₂→2SO₃, $\Delta n_g=-1$): high P helps, moderate T 450°C with V₂O₅.

Van't Hoff equation: $\dfrac{d\ln K}{dT} = \dfrac{\Delta H^\circ}{RT^2}$. Integrating: $\ln\dfrac{K_2}{K_1} = \dfrac{\Delta H^\circ}{R}\left(\dfrac{1}{T_1} - \dfrac{1}{T_2}\right)$. For endothermic ($\Delta H > 0$): increase T → K increases. For exothermic ($\Delta H < 0$): increase T → K decreases. $N_2 + O_2 \rightleftharpoons 2NO$, $\Delta H = +180.5$ kJ: At 25°C: $K \approx 10^{-30}$ (extremely unfavourable!). At 2000°C: $K \approx 0.1$ (significant). This is why NO only forms at very high temperatures (lightning, combustion engines, high-T industrial processes). The large endothermic $\Delta H$ means $K$ increases dramatically with temperature.

When $\Delta n_g = 0$: $K_p = K_c$ (independent of pressure). Pressure changes don't shift equilibrium. Mole fractions of all species remain unchanged when pressure changes. Examples: $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$: $\Delta n_g = 2-2 = 0$. $N_2(g) + O_2(g) \rightleftharpoons 2NO(g)$: $\Delta n_g = 2-2 = 0$. $CO(g) + H_2O(g) \rightleftharpoons CO_2(g) + H_2(g)$: $\Delta n_g = 2-2 = 0$. For these reactions: only temperature changes K and equilibrium yield. All other factors (pressure, concentration, catalyst, inert gas at constant V) leave the equilibrium position unchanged. Adding inert gas at constant V: no effect (partial pressures unchanged). Increasing total pressure by compression: no shift (equal gases both sides, mole fractions unchanged).

$N_2 + O_2 \rightleftharpoons 2NO$, $\Delta H = +180.5$ kJ. At room temperature: $K \approx 10^{-30}$ → essentially no NO in equilibrium. Thermodynamics says NO should not exist at room temperature! Yet NO is present in atmosphere (~0.3 ppb). Reason: kinetics — at room temperature, rate of decomposition and formation are both negligibly slow (high activation energy). Once NO is formed (e.g., in lightning or combustion), it does not decompose quickly at room temperature. This is a kinetically trapped non-equilibrium state. Similarly, diamonds are thermodynamically unstable relative to graphite at room temperature, but conversion is kinetically impossible. Key lesson: thermodynamic stability (K) does not predict practical existence — kinetics matters!

NO (nitric oxide, nitrogen monoxide): colourless, paramagnetic (1 unpaired electron, odd electron molecule). Bond order 2.5 ([He²σ²σ*²σ²π⁴π*¹]). Biologically: produced by nitric oxide synthase (NOS) from arginine. Important vasodilator (relaxes smooth muscle in blood vessel walls). Signalling molecule. Discovery of NO as biological messenger: Nobel Prize 1998 (Furchgott, Ignarro, Murad). Clinical: nitroglycerin and nitrites used for angina → release NO → vasodilation → reduce cardiac workload. Environmental: NO + O₂ → NO₂. 3NO₂ + H₂O → 2HNO₃ + NO (acid rain). NO + O₃ → NO₂ + O₂ (ozone depletion in stratosphere). Catalytic converter: 2NO + 2CO → N₂ + 2CO₂ (Pt/Rh catalyst).

$\Delta G = \Delta H - T\Delta S$. Spontaneous when $\Delta G < 0$. For $N_2 + O_2 \to 2NO$: $\Delta H = +180.5$ kJ (unfavourable). $\Delta S = +24.7$ J/mol/K (slightly positive, 2 mol gas from 2 mol gas → small entropy gain). $\Delta G = +180.5 - T(0.0247)$ kJ. $\Delta G = 0$ when $T = 180.5/0.0247 = 7309$ K. Only above 7309 K is $\Delta G < 0$ (spontaneous forward reaction). This explains why NO formation requires extremely high temperatures. At 2000°C (2273 K): $\Delta G = 180.5 - 2273(0.0247) = 180.5 - 56.1 = +124.4$ kJ (still non-spontaneous, but K is much larger than at 25°C because T increases K for endothermic reactions).

Industrial chemists face a dilemma: thermodynamics (equilibrium) vs kinetics (rate). Haber process: $N_2+3H_2\to2NH_3$, $\Delta H=-92$ kJ. Thermodynamics: low T → high K (more NH₃). Kinetics: high T → faster rate. Compromise: 400-500°C with Fe catalyst. Contact process: $2SO_2+O_2\to2SO_3$, $\Delta H=-196$ kJ. Similar dilemma. Compromise: 450°C with V₂O₅. Ethanol from ethylene: $C_2H_4+H_2O\to C_2H_5OH$, $\Delta H=-46$ kJ. Low T favoured thermodynamically. H₃PO₄ catalyst. Compromise: 300°C, 70 atm. The catalyst is crucial in all these: it increases rate without changing equilibrium → allows lower operating temperature (better yield) with reasonable production rate. Without catalyst: would need impractically high temperatures for acceptable rate.

The atmosphere as a chemical system: not at equilibrium (maintained by solar energy input). Major reactions: Ozone formation: $O_2 + h\nu \to 2O^\bullet$; $O^\bullet + O_2 + M \to O_3 + M^*$. UV-B and UV-C absorbed by O₃ in stratosphere. Ozone depletion: CFCs release Cl atoms under UV: $Cl^\bullet + O_3 \to ClO^\bullet + O_2$ (chain reaction). Montreal Protocol (1987) banned CFCs. Acid rain: SO₂ + H₂O → H₂SO₃; NO₂ + H₂O → HNO₃. pH of acid rain < 5.6 (natural CO₂ gives pH 5.6). Damages forests, limestone buildings (CaCO₃ + H₂SO₄ → CaSO₄ + H₂O + CO₂), aquatic ecosystems. Greenhouse effect: CO₂, CH₄, N₂O, H₂O absorb outgoing IR radiation → warming. Pre-industrial CO₂: 280 ppm. Current: 420 ppm (2023). 1.5°C of warming above pre-industrial expected by 2030s.