Ammonia will decompose into nitrogen and hydrogen at high temperature. An industrial chemist studying this reaction fills a 25.0 L tank with 8.3 mol of ammonia gas, and when the mixture has come to equilibrium measures the amount of hydrogen gas to be 2.5 mol. Calculate the concentration equilibrium constant for the decomposition of ammonia at the final temperature of the mixture. Round your answer to 2 significant digits. C

Suppose a 250. mL flask is filled with 0.60 mol of N₂ and 0.40 mol of O₂. The following reaction becomes possible: N₂(g) + O₂(g) → 2NO(g) The equilibrium constant K for this reaction is 9.69 at the temperature of the flask. Calculate the equilibrium molarity of N₂. Round your answer to two decimal places. X

Data Table Mg (s) MgO (s) Initial Temperature of HCI (aq) ("C) Final Temperature of solution ("C) Change in temperature ("C) 23.0 23.0 34.8 38.0 11.8 15.0 Mass of substance 0.13 0.97 (g) Concentration of HCI (aq) = 1.00mol/L Volume of HCI (ag) used in each trial = 0.050L

"Synthesis gas" is a mixture of carbon monoxide and water vapor. At high temperature synthesis gas will form carbon dioxide and hydrogen, and in fact this reaction is one of the ways hydrogen is made industrially. A chemical engineer studying this reaction fills a 125 L tank with 20. mol of carbon monoxide gas and 24. mol of water vapor. When the mixture has come to equilibrium he determines that it contains 7.0 mol of carbon monoxide gas, 11. mol of water vapor and 13. mol of carbon dioxide. The engineer then adds another 7.0 mol of carbon monoxide, and allows the mixture to come to equilibrium again. Calculate the moles of hydrogen after equilibrium is reached the second time. Round your answer to 2 significant digits. mol

Suppose a 250. mL flask is filled with 0.60 mol of SO₂ and 1.0 mol of SO3. This reaction becomes possible: 2SO₂(g) + O₂(g) — 2SO3(g) Complete the table below, so that it lists the initial molarity of each compound, the change in molarity of each compound due to the reaction, and the equilibrium molarity of each compound after the reaction has come to equilibrium. Use x to stand for the unknown change in the molarity of O₂. You can leave out the M symbol for molarity. initial change equilibrium SO₂ 0₂ 0 X 0 SO₂ 0 010 X Ś

Suppose a 250. mL flask is filled with 1.7 mol of NO2, 0.80 mol of CO and 0.70 mol of CO2. The following reaction becomes possible: NO2(g) + CO(g) = NO(g) + CO2(g) The equilibrium constant K for this reaction is 2.90 at the temperature of the flask. Calculate the equilibrium molarity of NO. Round your answer to two decimal places.

Suppose a 500. mL flask is filled with 1.5 mol of H₂ and 1.2 mol of Cl₂. The following reaction becomes possible: H₂(g) + Cl₂(g) → 2HCl(g) The equilibrium constant K for this reaction is 3.45 at the temperature of the flask. Calculate the equilibrium molarity of H₂. Round your answer to two decimal places. M Ś

Use this information to answer Questions 3, 4, and 5: The equilibrium constant (K) of the reaction below is K = 6.0 x 10-2, with initial concentrations as follows: [H2] = 1.0 x 102 M, [N2] = 4.0 M, and [NH3] = 1.0 x 10-4 M. N2(g) + 3H2(g) = 2NH3(g) Ton: N. x10-2 ent (O Pncca waythe chemical system would 5 the vara BI U = E T O Word(s) ------ T T; O Word(s) 5. If the concentration of the product NH3 was increased from 1.0x 10-4 M to 5.6 x 10-3 M, calculate the reaction quotient (Q) and determine which way the chemical system would shift by comparing the value of Q to K. B I U E T O Word(s)

Suppose a 250. mL flask is filled with 1.2 mol of N₂ and 1.4 mol of NO. The following reaction becomes possible: N₂(g) + O₂(g) → 2NO(g) The equilibrium constant K for this reaction is 0.481 at the temperature of the flask. Calculate the equilibrium molarity of NO. Round your answer to two decimal places. M X Ś