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Systems that have gaseous reactants and products behave a little bit differently when an experimenter works to force the system out of equilibrium in order to obtain a product. The mechanism of stress that is used is "pressure". [If you need to review this concept a little bit it can be found in Lesson 7.4 - Changing Reaction Rates. Guided notes might help support you, or even reviewing Unit 6 Rates of Reaction Lesson 2 a little bit. --- I am certain you can work through this!] Added pressure, or force exerted over area, can increase the number of collisions among particles in a closed system. Hopefully this makes sense. The system works to balance numbers of 'moles' of reactant(s) and product(s) throughout. We can find the number of moles by BALANCING the EQUATION. Here is our example from the Haber-Bosch process we used in Unit 6 portfolio work. It explains 'why' the reaction is more efficient at lower temperatures. N2(g) + 3H2(g) ⇌ 2NH3 (g) + heat energy nitrogen gas + hydrogen gas react reversibly to form ammonia gas + heat Please use the image above to help you count moles of reactants and products to determine which direction the shift will take the experimenter under the conditions applied. The equation is balanced for you. (6 points) Forward Reaction (shift right) Reverse Reaction (shift left) The experimenter ADDS heat to the closed system. The experimenter counts 2 moles of nitrogen and 6 moles of hydrogen gas in the closed system. The experiment counts 8 moles of ammonia in the closed system. The experimenter counts 1 mole of ammonia in the closed system. The experimenter ADDS the catalyst of iron to break the triple covalent bonds in nitrogen gas. [Check Lesson 7.4 for help if needed!] The experimenter REMOVES heat from the closed system. Skip to navigation they all nedd one

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To analyze how the equilibrium of the Haber-Bosch process shifts under different conditions, we can apply Le Chatelier's Principle. This principle states that if a stress is applied to a system at equilibrium, the system will adjust to counteract that stress and reestablish equilibrium.

The balanced chemical equation for the Haber-Bosch process is:

\[ \text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) + \text{heat} \]

Counting Moles:

  • Reactants:

    • Nitrogen (N₂) = 1 mole
    • Hydrogen (H₂) = 3 moles
    • Total = 4 moles of reactants

  • Products:

    • Ammonia (NH₃) = 2 moles
    • Total = 2 moles of products

Analyzing Each Condition:

  1. The experimenter ADDS heat to the closed system.

    • Here, we consider heat as a product since the reaction is exothermic.
    • According to Le Chatelier’s Principle, the system will respond to this stress by favoring the endothermic direction, which is the reverse reaction (shift left).
    • Shift Direction: Reverse Reaction (shift left)

  2. The experimenter counts 2 moles of nitrogen and 6 moles of hydrogen gas in the closed system.

    • Total moles of reactants = 2 (N₂) + 6 (H₂) = 8 moles.
    • Total moles of products = 8 moles of NH₃.
    • The system has more moles of reactants than products. However, since we already established that there is an imbalance favoring the reactants, the system will shift to the right to form more products until equilibrium is reestablished.
    • Shift Direction: Forward Reaction (shift right)

  3. The experiment counts 1 mole of ammonia in the closed system.

    • If there is only 1 mole of ammonia, there will be fewer products compared to reactants, causing the system to favor the forward reaction to produce more ammonia.
    • Shift Direction: Forward Reaction (shift right)

  4. The experimenter ADDS the catalyst of iron to break the triple covalent bonds in nitrogen gas.

    • Adding a catalyst speeds up both the forward and reverse reactions equally but does not alter the position of equilibrium.
    • The catalyst is used to help the system reach equilibrium faster but does not change the equilibrium concentrations of the reactants and products.
    • Shift Direction: No shift.

  5. The experimenter REMOVES heat from the closed system.

    • Again, since heat is a product, removing heat would shift the equilibrium toward the exothermic direction to produce more heat, favoring the forward reaction (toward the production of ammonia).
    • Shift Direction: Forward Reaction (shift right)

In summary:

  • Adding heat: Shift left (reverse reaction)
  • High amounts of reactants: Shift right (forward reaction)
  • Low amounts of products: Shift right (forward reaction)
  • Adding a catalyst: No shift
  • Removing heat: Shift right (forward reaction)

If you have any additional questions about the concepts or specific scenarios, feel free to ask!