$$ \mathrm{C}(\mathrm{~S})+\mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}_2(\mathrm{~g})+400 \mathrm{~kJ} $$   1 g mole $$ \begin{aligned} & \mathrm{C}(\mathrm{~s})+\frac{1}{2} \mathrm{O}_2(\mathrm{~g}) \rightarrow \mathrm{CO}(\mathrm{~g})+100 \mathrm{~kJ} \\ & 0.6 \times 1000=600 \mathrm{gm} \\ & 600 \times \frac{60}{100} \text { (Pure Carbon) } \\ & \quad=360 \mathrm{gm}=\frac{360}{12}=30 \text { mole (Pure Carbon) } \end{aligned} $$ Carbon converted into $\mathrm{CO}_2=\left(30-30 \times \frac{60}{100}\right)=12$ mole and carbon converted in $\mathrm{CO}=30 \times \frac{60}{100}=18$ mole Energy generated during II equation $=18 \times 100=1800 \mathrm{~kJ}$ Energy generated during $\mathrm{l}^{\text {st }}$ reaction. $=12 \times 400=4800$ Total $=1800+4800=6600 \mathrm{~kJ}$