We Are Presented with the Following System at Equilibrium

CHEM 504 Name:

Equilibrium Review

We are presented with the following system at equilibrium :

N2(g) + 3 H2(g)  2 NH3(g) + 102 kJ

What changes, among the following, will favour a forward reaction?

1.An increase in the NH3 concentration

2.A decrease in temperature

3.An increase in pressure brought on by a decrease in volume

4.Adding a catalyst

A) / 1 and 2 / C) / 2 and 3
B) / 1 and 4 / D) / 3 and 4

Which of the factors listed below would positively affect the production of carbon dioxide (CO2) in the equilibrium system?

CO(g) + NO2(g) CO2(g) + NO(g) + 227 kJ

1.Increase the pressure on the system.

2.Increase the temperature.

3.Decrease the temperature.

4.Add carbon monoxide (CO).

5.Add a catalyst.

A) / 1, 2 and 4 / C) / 2 and 5
B) / 1, 3 and 5 / D) / 3 and 4

The following table shows four systems at equilibrium and the change made to each system.

System at equilibrium / Change made
1.A(g) + 2B(g)  C(g) / The pressure was increased.
2.3D(g) + 2E(g)  2F(g) / A catalyst was added.
3.2E(g) + G(g)  E2G(g) / The quantity of G(g) was reduced.
4.A(g) + C(g)  3D(g) / The volume was increased.

In which system does the change favour the reverse reaction only?

A) / 1 / C) / 3
B) / 2 / D) / 4

The following chemical equations represent the ionization of four acids.

Which of the following is the strongest acid?

A) / HF(aq)  H+(aq) + F-(aq)Ka = 3.5  10-4
B) / HNO2(aq)  H+(aq) + NO2-(aq)Ka = 4.6  10-4
C) / NH4+(aq)  H+(aq) + NH3(aq)Ka = 5.6  10-10
D) / HCO3-(aq)  H+(aq) + CO32-(aq)Ka = 5.6  10-11

In studying the equilibrium of a chemical system, a student observed that the behaviour of the chromate ion, CrO42(aq), depends on the acidity of the system. When the acidity of the system changes, the colour of the solution can vary from yellow to orange. This reaction is illustrated by the following equation :

2 CrO42(aq) + 2 H+(aq)  Cr2O72(aq) + H2O(l)

Which expression should be used to find the equilibrium constant, Kc, of this system?

A) / / C) /
B) / / D) /

Consider the following equilibrium reaction:

A(g) + B(g) C(g) + D(g)

The equilibrium constants, at two different temperatures, are provided in the table below.

Temperature (°C) / Equilibrium constant (K)
450 / 1.8  102
550 / 8.4  101

Is the forward reaction endothermic or exothermic?

A) / Endothermic, because K increases with temperature and therefore favours the formation of products.
B) / Exothermic, because K increases with temperature and therefore favours the formation of products.
C) / Endothermic, because K increases with temperature and therefore favours the reduction of products.
D) / Exothermic, because K increases with temperature and therefore favours the reduction of reactants.

/ When the sulfur dioxide emitted by the metallurgical industry and internal combustion engines in automobiles combines with oxygen in the air, the reaction is the following :
2 SO2(g) + O2(g)  2 SO3(g) + 193 kJ
Seeing how this reaction is a reversible one, and the planet Earth a possible equilibrium system, your classmate states that the formation of sulfur trioxide would be favored by these conditions : a cold winter day with high atmospheric pressure.

Considering Le Chatelier’s principle, do you believe these two conditions will shift equilibrium favoring sulfur trioxide production?

Below is an incomplete data chart showing certain characteristics of two acids at 25°C.

Acids at Equilibrium / Initial Concentration (mol/L) / Percentage Dissociation (%) / pH / Ka
H2CO3  H+ + HCO3-
H2S  H+ + HS- / 0.1
----- / -----
0.095 / 3.7
----- / -----
9.1  10-8

Which acid is the stronger of the two?

One step in the manufacture of sulfuric acid consists of reacting sulfur dioxide with oxygen gas to form sulfur trioxide according to the following reversible reaction :

2 SO2(g) + O2(g)  2 SO3(g) + 205 kJ

When this system is at equilibrium, what will be the effect of increasing the temperature on the concentration of each substance in this reaction and on the value of the equilibrium constant?

Ammonia is produced in a 100-L container. When the system reaches equilibrium, 20 moles of hydrogen gas, 10 moles of nitrogen gas and 25 moles of ammonia are present.

The equation for this reaction is :

N2(g) + 3 H2(g)  2 NH3(g) + Energy

Calculate the equilibrium constant, Ke, for this reaction.

The following equation represents the formation of hydrogen iodide, HI(g), from its elements :

H2(g) + I2(g)  2 HI(g) + 11 kJ

How will a temperature increase affect the value of the equilibrium constant for this system?

Heat is used to decompose calcium carbonate, CaCO3(s), according to the following chemical equation :

CaCO3(s)  CaO(s) + CO2(g)

At a given temperature, a closed system contains CaCO3(s) in equilibrium with CaO(s) and CO2(g). While the temperature is kept constant, the volume of this system is reduced. What happens to the quantity of each substance in this system?

Vinegar is a weak acid commonly used in preparing food. In fact, vinegar is a highly diluted solution of acetic acid, HCH3COO. The ionization of acetic acid is represented by the following equation :

HCH3COO(aq)  H+(aq) + CH3COO(aq) Ka = 1.8  105 at 25C

At a temperature of 25C, the concentration of a solution of vinegar is 5 % m/V, namely 5 grams of HCH3COO per 100 mL of solution. What is the molar concentration of H+(aq) in this solution?

The initial concentration of an acid (HX) is 2.25 M. Calculate the dissociation constant (Ka) for this acid if the pH is 3.4.

Carbonic acid, H2CO3, is a weak acid. The dissociation of carbonic acid and the ionization constant, Ka, are shown below.

H2CO3(aq)  H+(aq) + HCO3(aq)Ka = 4.3  107

A chemistry student places 3.1  102 grams of carbonic acid into 5.0  102 mL of distilled water. What is the pH of this solution?