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If I understand it correctly, products will be favored at equilibrium if you increase the amount of reactant when the reaction is at equilibrium in order to minimize the effect of adding more reactant; the forward reaction is favored, producing more product. Initial concentrations of reactants or products do not necessarily tell you whether or not the reactant/product is favored at equilibrium.
Something that helps me figure out which side of the equation is favored is to visualize the equation as a tub of water. If you add more water to the left side (adding more of a reactant), then the water will move to the right side (favor the products) in order to reach equilibrium. The opposite is true if you add more water to the left side (adding more product) which will cause water to move to the left side (favoring the reactants). If you take water away from the left side (remove reactants), then water from the right side will move to the left in order to fill in the hole (favoring more reactants) and vice versa.
In regards to amount/number of moles, I like to visualize Le Chatelier's Principle like a seesaw. Let's say the left of the seesaw are the reactants and the right of the seesaw are the products. Increasing the amount of reactant leaves the seesaw tilted, having more weight on the left side. In order to achieve an equilibrium (seesaw is neutral and a horizontal line), more products are formed and is favored in the forward reaction.
Like the replies above, it is very helpful to view these changes visually. Similarly, it's helpful to think of adding reactants or products to a reaction at equilibrium as tipping a scale or seesaw. In order to react to this increase, either the forward or reverse reaction is favored to even everything out.
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