In the haber process for the preparation of NH3 which is given by the equilibruim?
N_2_(g)+3H_2_(g) <---> 2NH_3_(g)...... delta H= -92.38kJ
State the conditions of concentrations and temperture that will favor production of NH_3_?
Please help! would it be high tempertures? or am I wrong? please explain why cause I'm lost!
First, most everyone here recognizes that subscripts are a problem so we write it as
N2(g) + 3H2(g( ==> 2NH3(g) + heat
and keep in mind that numbers before the element are coefficients and numbers after the element are subscripts.
This problem is one on Le Chatelier's Principle.
To favor the production of NH3, we want the reaction to go to the right as much as possible. So we increase (N2) and/or (H2), decrease (NH3) as it is produced, and keep temperature low. Increasing pressure favors production of NH3 also.
Note: Remember Le Chatelier's Principle says that a reaction at equilibrium will try to UNDO what we do to it. So increasing temperature, as you suggest, would try to use up the added heat so the reaction would shift to the left (look at my equation above to see that) and that will DECREASE NH3. See also that adding N2 and/or H2 will try to use up the added N2 and H2 and that means more NH3 production. For pressure, the reaction shifts to the side with fewer mole when increasing P. There are 4 moles of gas on the left and 2 on the right so increasing P moves to more NH3 by shifting to the right. A rather long explanation and somewhat repetitious in spots but I hope this helps.
how to find reagents need to dissolve a chemical in a system? If I have a equilibrium equation would I use the chemical on the product side or reactant side to dissolve a compound.
You may be referring to the SbCl3 post which I answered just a second ago.
In the Haber process, the formation of ammonia (NH3) from nitrogen gas (N2) and hydrogen gas (H2) is an exothermic reaction, meaning it releases heat. The equation for the reaction is:
N2(g) + 3H2(g) ↔ 2NH3(g)
To determine the conditions of concentrations and temperature that favor the production of NH3, we need to consider Le Chatelier's principle. This principle states that when a system at equilibrium is subjected to a change in conditions, it will adjust to minimize the effect of that change. In this case, the change in conditions refers to the concentrations of the reactants and the temperature.
1. Concentration of Reactants:
According to Le Chatelier's principle, if the concentration of a reactant is increased, the equilibrium will shift in the direction that minimizes the effect of the change. In this case, if the concentrations of N2 or H2 are increased, the equilibrium will shift to the right, favoring the production of NH3. Therefore, higher concentrations of N2 and H2 will favor the production of NH3.
2. Temperature:
Changes in temperature can also affect the equilibrium position. In an exothermic reaction like the Haber process, increasing the temperature decreases the yield of the product (NH3). This is because the reaction will try to consume the additional heat energy by shifting in the direction that consumes heat. Therefore, lower temperatures favor the production of NH3.
Considering both concentration and temperature, increasing the concentrations of N2 and H2 while keeping the temperature low would be the conditions that favor the production of NH3. This is because higher reactant concentrations increase the likelihood of successful collisions between molecules, and a lower temperature improves the yield of ammonia.
To summarize, the conditions that favor the production of NH3 in the Haber process are high concentrations of N2 and H2, and a low temperature.