In which direction (Left, right, or neither direction) would the reaction below have to shift to regain equilibrium, if the reaction was disturbed from the equilibrium by the remove of some SO3 gas? Explain.  2 SO3 (g)  ----->     2 SO2 (g) + O2 (g)             DeltaH= +198 kJ             <------

Consider the following exothermic reaction: A(g) + B(g) + 2C(g) Fill in the missing spots with either "products", "reactants", or "neither". 1) If the system is at equilibrium and one removes C, then the reaction will shift towards 2) If the system is at equilibrium and one adds C, then the reaction will shift towards 3) If the system is at equilibrium and one removes A, then the reaction will shift towards 4) If the system is at equilibrium and one raised the pressure, then the reaction will shift towards 5) If the system is at equilibrium and one raises the temperature, then the reaction will shift towards

Consider the following exothermic reaction system at equilibrium: N2(g) + 3H2(g) ⇒ 2NH3(g) Predict how the following changes affect the number of moles of each component of the system after equilibrium is reestablished by completing the table below: Note: Use I for increase, D for decrease, NC for no change Add N2(g) Remove H2(g) Add NH3(9) Add Ne(g) Increase the temperature (constant P) Decrease the volume (constant T) Add a catalyst Submit NH3 N₂ H2

5. Consider the following reaction and Kc determined at a particular temperature: NH4HS(s) NH3(g) + H2S(g) Ko = 8.95 x 10-2 A sealed 250.0 mL vessel contains 0.100 g NH4HS(s), 0.150 g NH3(g), and 0.230 g H2S(g). In which direction will the reaction proceed in order to reach equilibrium? Answer towards products or towards reactants. Supporting calculations are required. 6. Consider the following reaction: Co(g) + H2O(g) → CO2(g) + H2(g) Ko = 0.58 at 1273 K A sealed 4.00 L vessel initially contains 1.00 mole CO2(g), 1.00 mole H2(g), 0.600 moles CO(g), and 0.600 moles of H2O(g). Calculate the moles of H2(g) present at equilibrium. Reference: General Chemistry 8th Ed: Ebbing and Gammon: 2005 by Houghton Mifflin. New York. page 654, print 68

(2) Nitrogen oxide is a pollutant in the lower atmosphere that irritates the eyes and lungs and leads to the formation of acid rain. Nitrogen oxide forms naturally in the atmosphere according to the endothermic reaction: N2 (g) + O2 (g) = 2 NO (g) Kp = 4.1 × 10-31 at 298 K Use the ideal gas law to calculate the concentrations of nitrogen and oxygen present in air at a pressure of 1.0 atm and a temperature of 298 K. Assume that nitrogen composes 78% of air by volume and that oxygen composes 21% of air. Find the “natural” equilibrium concentration of NO in air in units of molecules/cm³.

For the reaction 2 A( g)ΔB( g), the equilibrium constant is Kp = 0.76. A reaction mixture initially contains 2.0 atm of each gas (PA = 2.0 atm and PB = 2.0 atm). Which statement is true of the reaction mixture?a) The reaction mixture is at equilibrium.b) The reaction mixture will proceed toward products.c) The reaction mixture will proceed toward reactants.d) It is not possible to determine from the information given the future direction of the reaction mixture.

Consider the following reaction at equilibrium. What effect will increasing the temperature have on the system? FeзО4(s) + CO(g) — 3 FeO(s) + СO2(9) AH°= +35.9 kJ The equilibrium constant will decrease. No effect will be observed. The concentration of CO2 (g) will increase.. The concentration of CO (g) will increase. The reaction will shift to the right in the direction of products. о

choose ONE of each of the red selections. for example you choose between endothermic or exothermic and fill out the other options to make the statement true

Calculate the pressure of N2 needed to reverse the reaction (last row of the table. 0.38 is wrong). thank you!

Consider the following equilibrium: N₂O4 (g) → 2NO₂ (g) AG = 5.4 kJ Now suppose a reaction vessel is filled with 0.399 atm of dinitrogen tetroxide (N₂O4) at 588. °C. Answer the following questions about this system: Under these conditions, will the pressure of N₂O4 tend to rise or fall? Is it possible to reverse this tendency by adding NO₂? In other words, if you said the pressure of N₂O4 will tend to rise, can that be changed to a tendency to fall by adding NO₂? Similarly, if you said the pressure of N₂O4 will tend to fall, can that be changed to a tendency to rise by adding NO₂? If you said the tendency can be reversed in the second question, calculate the minimum pressure of NO₂ needed to reverse it. Round your answer to 2 significant digits. O O rise fall yes no atm x10 X 5