Q 5(a) (i) Using the information contained in the table below, calculate the free energy of formation, ?Gf?, of H?O(g) at 298 K, and Keq for the same reaction at 350 K: Compound | ?H?f (kJ mol?¹) | S? (J mol?¹ K?¹) --- | --- | --- H?O(g) | -241.818 | 188.925 H?(g) | 0 | 130.684 O?(g) | 0 | 205.138 (ii) For the reaction in 5(a)(i), sketch (approximate) graphs of ln(Keq) as a function of the inverse of temperature (1/T), AND of ?Gf? as a function of temperature (T). Briefly describe the features of both graphs.
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For the reaction, ΔH° and ΔS° can be calculated as follows: ΔH° = Σ ΔH°f(products) - Σ ΔH°f(reactants) = 2*(-241.818 kJ/mol) - [2*0 kJ/mol + 0 kJ/mol] = -483.636 kJ ΔS° = Σ S°(products) - Σ S°(reactants) = 2*188.925 J/mol.K - [2*130.684 J/mol.K + Show more…
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If a chemical system at equilibrium is disturbed, the rates of the forward and reverse reactions are temporarily unequal. This is known as ______. the Law of Mass Action The Law of Chemical Equilibrium the Haber-Bosch Process the Unequal Rate Principle Le Chatelier's Principle The following reaction occurs in a closed system. CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g) ΔH = +210 kJ Using Le Chatelier's Principle, check all boxes that apply when pressure is increased. More products will be formed. Keq will decrease. The amount of products and reactants will remain constant. Keq will increase. More reactants will be formed. Keq will not change.
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Calculate $\Delta G^{\circ}$ for each of the following reactions from the equilibrium constant at the temperature given. (a) $\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}(g) \quad \mathrm{T}=2000^{\circ} \mathrm{C} \quad K_{p}=4.1 \times 10^{-4}$ (b) $\mathrm{H}_{2}(g)+\mathrm{I}_{2}(g) \longrightarrow 2 \mathrm{HI}(g) \quad \mathrm{T}=400^{\circ} \mathrm{C} \quad K_{p}=50.0$ (c) $\mathrm{CO}_{2}(g)+\mathrm{H}_{2}(g) \longrightarrow \mathrm{CO}(g)+\mathrm{H}_{2} \mathrm{O}(g) \quad \mathrm{T}=980^{\circ} \mathrm{C} \quad K_{p}=1.67$ (d) $\mathrm{CaCO}_{3}(s) \longrightarrow \mathrm{CaO}(s)+\mathrm{CO}_{2}(g) \quad \mathrm{T}=900^{\circ} \mathrm{C} \quad K_{p}=1.04$ (e) $\mathrm{HF}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{F}^{-}(a q) \quad \mathrm{T}=25^{\circ} \mathrm{C} \quad K_{p}=7.2 \times 10^{-4}$ (f) $\operatorname{AgBr}(s) \longrightarrow \operatorname{Ag}^{+}(a q)+\operatorname{Br}^{-}(a q) \quad \mathrm{T}=25^{\circ} \mathrm{C} \quad K_{p}=3.3 \times 10^{-13}$
Calculate $\Delta G^{\circ}$ for each of the following reactions from the equilibrium constant at the temperature given. (a) $\mathrm{Cl}_{2}(g)+\mathrm{Br}_{2}(g) \longrightarrow 2 \mathrm{BrCl}(g) \quad \mathrm{T}=25^{\circ} \mathrm{C} \quad K_{p}=4.7 \times 10^{-2}$ (b) $2 \operatorname{SO}_{2}(g)+\mathrm{O}_{2}(g)=2 \operatorname{SO}_{3}(g) \quad \mathrm{T}=500^{\circ} \mathrm{C} \quad K_{p}=48.2$ (c) $\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{H}_{2} \mathrm{O}(g) \quad \mathrm{T}=60^{\circ} \mathrm{C} \quad K_{p}=0.196 \mathrm{atm}$ (d) $\operatorname{CoO}(s)+\mathrm{CO}(g) \rightleftharpoons \mathrm{Co}(s)+\mathrm{CO}_{2}(g) \quad \mathrm{T}=550^{\circ} \mathrm{C} \quad K_{p}=4.90 \times 10^{2}$ (e) $\mathrm{CH}_{3} \mathrm{NH}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{CH}_{3} \mathrm{NH}_{3}^{+}(a q)+\mathrm{OH}^{-}(a q) \quad \mathrm{T}=25^{\circ} \mathrm{C} \quad K_{p}=4.4 \times 10^{-4}$ (f) $\mathrm{PbI}_{2}(s) \longrightarrow \mathrm{Pb}^{2+}(a q)+2 \mathrm{I}^{-}(a q) \quad \mathrm{T}=25^{\circ} \mathrm{C} \quad K_{p}=8.7 \times 10^{-9}$
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