## 4 Results and Discussion  
The stability of a hydrate phase will ultimately depend on all independent thermodynamic variables of the systems. These include temperature, pressure, and concentrations of the components in all of the co-existing phases. It is possible to plot various projections of hydrate stability onto a sub-set of independent thermodynamic variables, with temperature and pressure projections being the most typical example, as illustrated by stability the limits of \(\text{CH}_{4}\) and \(\text{CO}_{2}\) hydrates in Figure 1. As noted previously, the phase transition undergone by \(\text{CO}_{2}\) will significantly affect the temperature-pressure projection of its hydrate existence region. Given that temperatures and pressures are independent thermodynamic variables, it is the hydrated Gibbs free energy that will determine the phase stability. The free energies of hydrates formed along the corresponding hydrate stability limits of Figure 1 are plotted in Figure 4. \(\text{CO}_{2}\) will form a more stable hydrate  
Figure 3: Concentration profile for \(\text{CO}_{2}\) across water (\(R=0\))/hydrate (\(R=12\)) interface [39; 40; 41]. \begin{table}
\begin{tabular}{c c c c c c} \hline \hline
**i** & **Parameter** & **I** & **Parameter** & **i** & **Parameter** \\ \hline
1 & 0.979242 & 4 & 171.673 & 7 & –9649.96 \\ \cline{2-6}
2 & 15.5427 & 5 & 6.76975 & 8 & 14,779.7 \\ \cline{2-6}
3 & \(-88.5112\) & 6 & 1939.55 & 9 & –7496.15 \\ \hline \hline \end{tabular}
\end{table} 
Table 3: Parameters for Equation (10). than CH4 for the whole range of thermodynamic conditions covered in Figure 1. In order to illustrate impact of a typical hydrate inhibitor, we also present the curves of Gibbs free energy for hydrate formation from the aqueous phase containing various mole fractions of methanol. The more methanol is added, the less stable the formed hydrate will be. Since Figure 5 may not be easily interpreted, the understanding of its plots should benefit from an additional reading of the two related 2D figures. Figure 5 shows the co-variation of pressure and temperature, while Figure 6 below shows enthalpy changes along the temperature projection of Figure 1.  
What is even more important in the context of this work are the differences in enthalpies of hydrate formation along the hydrate formation curves presented in Figure 6. One of the reasons for presenting the data for varying concentrations of methanol in the aqueous phase in Figure 4 is the fact that methanol also will act as a surfactant where the water/hydrate-former phase is concerned. This can be used to stimulate the CO2/CH4  
Figure 4: Calculated free energies of CO2 hydrates (solid) and CH4 hydrates (dashed). For each set of curves, the lowest curve is the free energy of hydrate formed from pure water, with the curves above corresponding to the following methanol content: 2 mol %, 4 mol %, 6 mol %, 8 mol %, 10 mol %, and finally 12 mol % for the topmost curves of CO2 and CH4, respectively. Figure 5: Enthalpy changes associated with hydrate formation along the hydrate P-T equilibrium curves for CO2 hydrate (solid) and CH4 hydrate (dashed). Units are kJ/mol hydrate former. exchange process since the methanol-enriched interface will hinder the formation of a sealing hydrate film between water and hydrate-former phases. In accordance with Figure 5, the heat of hydrate formation will be about 10 kJ per mol of hydrate former larger for the CO2 hydrate compared to methane hydrate. Heat released when new CO2 hydrate is created from the injected mixture of CO2 and surfactant will be transported through the aqueous phase in front of the CO2 plume. A portion of heat will be lost to the heating of the bedrock; an increased concentration of ions due to pore water being consumed by the forming CO2 hydrate is another important side effect of this process. The latter phenomenon is where Figure 5 comes into play more directly, since CH4 will dissolve more vigorously when the ion activity goes up.