Equilibrium constants

The equilibrium constants, temperatures and reaction enthalpies used in the chemical-equilibrium unit. Every K(T) on the benches is computed from these values via van ’t Hoff or thermodynamic data.


$\ce{H2(g) + I2(g) <=> 2HI(g)}$	Kc	54.3	430	−9.5	Δn = 0, so Kc = Kp
$\ce{N2O4(g) <=> 2NO2(g)}$	Kp	0.148	25	57.2	from NBS ΔfG° values
$\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$	Kp	5.44 × 10⁵	25	−91.8	K is huge, but the reaction is very slow at room temperature
$\ce{N2(g) + 3H2(g) <=> 2NH3(g)}$	Kp	4.90 × 10⁻⁵	450	−105.7	with heat capacities (NIST Shomate parameters)
$\ce{CaCO3(s) <=> CaO(s) + CO2(g)}$	Kp	0.229	800	169.4	Kp = p(CO₂)/bar; with heat capacities (ΔrH° at 800 °C)
$\ce{H2O(l) <=> H2O(g)}$	Kp	0.0317	25	44.0	K is the vapour pressure in bar
$\ce{Fe^3+(aq) + SCN^-(aq) <=> FeSCN^2+(aq)}$	K°	708	25	−6.3	at zero ionic strength
$\ce{[Co(H2O)6]^2+(aq) + 4Cl^-(aq) <=> [CoCl4]^2-(aq) + 6H2O(l)}$	K′	5.00 × 10⁻⁶	25	62.0	conditional constant (model value calibrated to the classroom demonstration)
$\ce{CH3COOH + C2H5OH <=> CH3COOC2H5 + H2O}$	K	4	—	—	about 4 (Berthelot and Péan de Saint-Gilles, 1862)
$\ce{2H2(g) + O2(g) <=> 2H2O(l)}$	K	1.22 × 10⁸³	25	−571.7	from ΔfG°(H₂O, l) = −237.13 kJ/mol
$\ce{N2(g) + O2(g) <=> 2NO(g)}$	K	4.72 × 10⁻³¹	25	180.5	from ΔfG°(NO) = 86.55 kJ/mol

Standard states: 1 mol/L for solutes, 1 bar for gases; K is treated as dimensionless. ΔrH° is the 25 °C value (except the NH₃ row at 450 °C).

Sources

D. D. Wagman et al., “The NBS tables of chemical thermodynamic properties”, J. Phys. Chem. Ref. Data 11, Suppl. 2 (1982).
NIST Chemistry WebBook, SRD 69 (Shomate parameters for N₂, H₂, NH₃).
K. de Berg, M. Maeder, S. Clifford, Inorg. Chim. Acta 466 (2017) (FeSCN²⁺); H. S. Frank, R. L. Oswalt, J. Am. Chem. Soc. 69, 1321 (1947).
T. W. Swaddle, L. Fabes, Can. J. Chem. 58, 1418 (1980) (cobalt chloride enthalpy).
W. M. Haynes (ed.), CRC Handbook of Chemistry and Physics, 97th ed. (2016) (vapour pressure of water).

Sources

Used in these lessons