Q.Prepare concept maps of chemical equilibrium.

For a general reversible reaction $aA+bB\rightleftharpoons cC+dD$, setting the forward rate ($k_f[A]^a[B]^b$) equal to the reverse rate ($k_r[C]^c[D]^d$) at equilibrium gives a constant ratio called the equilibrium constant: $K_c=\dfrac{[C]^c[D]^d}{[A]^a[B]^b}$, using equilibrium concentrations (mol dm$^{-3}$). For gaseous reactions it is often more convenient to use partial pressures instead: $K_p=\dfrac{P_C^{\,c}P_D^{\,d}}{P_A^{\,a}P_B^{\,d}}$. When concentrations used in this same ratio are NOT necessarily equilibrium values, the ratio is called the reaction quotient $Q_c$; comparing $Q_c$ with $K_c$ predicts which way a reaction must shift to reach equilibrium (forward if $Q_c<K_c$, reverse if $Q_c>K_c$, already at equilibrium if $Q_c=K_c$). The reverse reaction's equilibrium constant is the reciprocal of the forward one, $K'_c=1/K_c$. The unit of $K_c$ is $(\text{mol dm}^{-3})^{\Delta n}$, where $\Delta n$ is the difference between the total moles of species in the numerator and denominator. In a heterogeneous equilibrium (more than one phase present, e.g. a solid subliming to a gas), only gaseous and dissolved species are included in the expression, since a pure solid or pure liquid's concentration (activity) is constant and gets absorbed into $K_c$ itself. Key characteristics: $K_c$/$K_p$ i …