If a vessel, containing an equilbrium mixture of these gases, is pressurized by adding helium gas, then the equilibrium position will shift to the right.
N2(g) + 3H2(g) - 2NH3(g)
If an equilbrium mixture of these gases is released into a vessel of larger volume, then the equilibrium position will shift to the left.
Fe3+(aq) + SCN-(aq) - FeSCN2+(aq)
If AgNO3 is dissolved in this solution, then the equilibrium position will shift to the left. (Note that AgSCN is insoluble.)
H2(g) + I2(g) - 2HI(g)
If hydrogen gas is added to an equilbrium mixture of these gases, then the equilibrium position will not shift.
H2(g) + I2(g) - 2HI(g) (Endothermic)
If the temperature is increased, then the equilibrium constant for this reaction equation will decrease.How do you do these?!?! Is the statement true or false, with respect to the specified reaction in each case?
use LeChatlier's principle - simply that placing stress/change on equilibrium will cause equilibrium to shift to accomodate/adjust for the stress.
In the First reaction adding more gas, increasing pressure would not affect the equilibrium because two moles of gas form two moles of gas... pressure would not be changed by either the forward or reverse reaction and helium is not involved in the reaction so the answer is FALSE
the second is TRUE, larger vessel would decrease pressure, the reverse reaction would produce more molecules to accomodate the lowered pressure.
three is TRUE. Addition of AgNO3 would precipitate out SCN- ions, reducing the amount of SCN-, equilibrium would shift to produce some more SCN- ion, shifting to the left
the HI(g) reaction would shift... the answer is FALSE - addition of the reactant hydrogen gas would push the equilibrium favoring the forward reaction (consumption of some of the added H2(g)
increased temperature would favor the forward reaction, causing the new equilibrium constant to be larger.... statement 5 is FALSE
For a great review of this subject, see the source cited
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