Which of the following would shift the equilibrium in a combustion reaction towards the products?

In a combustion reaction, shifting the equilibrium towards the products can be influenced by various factors. Increasing the concentration of reactants is an effective way to drive the reaction forward. According to Le Chatelier's principle, when the concentration of reactants is increased, the system will adjust by favoring the formation of products to restore equilibrium.

In combustion reactions, the reactants are primarily fuels (such as hydrocarbons) and oxygen, and as their concentrations increase, more product (such as carbon dioxide and water) will be generated. This principle applies regardless of the specific details of the reaction, as it is a fundamental concept in understanding chemical equilibria.

While other options like decreasing temperature or increasing pressure may affect the position of equilibrium in certain contexts, they are not always straightforward in combustion reactions. Decreasing temperature generally favors exothermic reactions, but depending on the specific reaction conditions, it might not significantly drive combustion towards more products. Similarly, increasing pressure may not consistently shift the equilibrium towards the products in reactions with equal numbers of gas molecules on both sides.

Thus, the most clear-cut method to favor product formation in this scenario is by increasing the concentration of reactants.