This can be explained thermodynamically, taking entropy and free energy changes into account.
\(C(s)+O_{2}(g) \rightarrow C O_{2}(g)\) 
\(2 C(s)+O_{2}(g) \rightarrow 2 C O(g)\) 
Case i: Volume of CO₂ produced  = Volume of O₂ used.
\(\therefore \triangle S\) is very small and \(\triangle G\) does not change with temperature.
\(\therefore\) Plot of \(\triangle G\) Vs T is almost hrizontal
Case ii: Volume of CO produced = 2 x Volume of O₂ used
\(\therefore \triangle S\) is positive and hence \(\triangle G\) beccomes increasingly negative as the temperature increases.
\(\therefore\) Plot of \(\triangle G\) Vs T slopes downwards.
As can be seen from ΔG° Vs T plot (Ellingham diagram), lines for the reactions, C ——–> CO₂ and C ——–> CO cross at 983 K. Below 983 K, the reaction (a) is energetically more favourable but above 673 K, reaction (b) is favourable and preferred. Thus, below 673 K both C and CO can act as a reducing agent but since CO can be more easily oxidised to CO₂ than C to CO₂ , therefore, below 673 K, CO is more effective reducing agent than carbon.