CHEMICAL EQUILIBRIUM

I. INTRODUCTION

A. Principle

If you are in a bus stop and you observed that the rate at which people ride the bus is equal to

the rate at which people leave the bus, the number of people riding the bus and leaving the bus remains

constant and can be said to be in dynamic equilibrium, much like in chemicals reactions. They can occur

in both forward and reverse directions, and when the rates of reactions became equal, the concentrations

of reactants and products remain constant; at that point, the chemical system is at equilibrium

(McMurry& Fay, 2008).

Chemical equilibrium as defined by Lower is a chemical reaction in which there is no tendency for

the quantities of reactants and products to change.

Chemical change occurs when the atoms that make up one or more substance rearrange in such

a way that new substances are formed. These substances are the components of the chemical reaction

system; those components which decrease in quantity are called reactants while those which increase are

called products. A given chemical reaction system is defined by a balanced net chemical reaction which is

conventionally written as :

Reactants -> products

There are some factors which instigate these changes to disturb a system at equilibrium namely,

the concentration, pressure and temperature.

The effect of any change in the reaction condition of a system in equilibrium can be described

using the Le Chatelier’s Principle. The principle states that if a system at equilibrium is subjected to a

change of pressure, temperature, or number of moles of a substance, there will be a tendency for a net

reaction in the direction that tends to reduce the effect of this change.

B. Objectives

At the end of the experimentation, the student was able to:

1. to determine how different factors affect a system in chemical equilibrium;

2. to explain the effect of these factors in terms of the Le Chatelier’s principle.

II. MATERIALS

A. Reagents 0.02 M KSCNH2O0.02 M Fe(NO3)3

NaN02

0.02 M NaH2PO4

6 M HCLCoCl2

ice

B. ApparatusTest tubeStirring rodCentrifuge tubesSyringeRubber stopperHot plateBeaker

III. PROCEDURE

A. Effect of Concentration

5 ml of 0.02 M KSCN was mixed with 5 ml water in a test tube. Afterwards, 2-4 drops of

0.02 M Fe(NO3)3 was added to the solution in the test tube and was stirred. Observations were

noted and were recorded on the table provided.

Subsequently, the formed solution was divided into four labelled test tubes (1-4). The

first test tube served as the control in the experiment. The second was added with a small crystal

of KSCN. Test tube 3 was added with a drop of 0.02 Fe(NO3)3 and test tube 4 with 2 drops of 0.02

M NaH2PO4.

Observations on physical properties of the solution on each testtube were recorded on

the table.

B. Effect of Pressure

Under the fumehood, a pinch of NaNO2 and 3-4 drops of 6M HCl were mixed in a test

tube to generate nitrogen dioxide (NO2).

After sometime, 10 cc of brown gas was drawn into a syringe and was pressed into a

rubber stopper to prevent the gas from escaping. It was placed in a white background to note for

the observations on its color.

Next, as the syringe was constrained on a rubber stopper, the plunger was pressed from

the 10 cc until the 4 cc mark and observations on the initial change of gas color were

distinguished.

After few seconds of maintaining the same pressure, the color variations of the gas

inside the syringe was also noted.

Observations were tabulated on the given table.

C. Effect of Temperature

The prepared solution of CoCl2 in 12 M HCl was placed into three test tubes containing 0.5

ml of this solution.

One test tube was placed in a hot-water bath, the second in a cold-water bath, and the third,

in room temperature which served as the control. After 5 mins, observations on the color variations

of each solution were noted.

Afterwards, the test tube from the hot-water bath was transferred to the cold-water bath

and vice versa. It was allowed to stand for another 5 minutes then the observations were again noted

on a table.

IV. DATA AND RESULTS

Table 4.1 Observations on the addition of 0.02 M Fe(NO3)3to KSCN solution.

Step Observations

Preparation of 5 ml 0.02 M KSCN + 5 ml H2O Clear, colorless solution

Addition of 2-4 drops of 0.02 M Fe(NO3)3 Deep red solution

Table 4.2 Preparation of different solutions of KSCN.

TEST TUBE CONTENTS OBSERVATIONS

1 Prepared solution Bloody red in solution

2 Prepared solution + KSCN crystal Dark red solution

3 Prepared solution+ 1 drop 0.02 M

Fe(NO3)3

Darker red solution (in reference to test tube

no.2)

4 Prepared solution + 2 drops 0.02 0.02 M NaH2PO4

Yellowish solution

Table 4.3 Effect of pressure on color changes of gas.

Table 4.4 Effect of temperature to the CoCl2 solution.

Syringe plunger position Observations

B Brown gas

C. Light brown gas

A.Lighter brown gas

V. DISCUSSION

Chemical equilibrium is the state reached when the concentrations of reactants and products

remain constant over time.

There are some factors which instigate these changes to disturb a system at equilibrium namely,

the concentration, pressure and temperature.

Table 4.1 shows the observations as KSCN was added with drops of Fe(NO3)3. It can be observed

that when KSCN was first added with water, the solution formed was colorless and when the additional

drops of Fe(NO3)3 was dispersed in the test tube, the colorless solution turned to a bloody red solution

which is due to the formation of FeSCN2+ with a net ionic equation of:

Fe3+ + SCN -> FeSCN2+

Table 4.2 shows the variation in the color of the solution as different reagents were added on it.

Shifts in the position of its equilibrium can be detected by observing how the color of the solution

changes when various reagents were added.

For example, as an additional crystal of KSCN was added to the test tube containing FeSCN2+

solution, the color of the solution turned to a darker shade of red, in contrast to the bloody red color

of the test tube containing FeSCN2+ alone. Furthermore, as Fe(NO3)3 was added to another test tube

containing the prepared solution of FeSCN2+, the color also changes, presenting a a darker shade of

InitialTemperature

Observations ShiftingTemperature

Observations

Hot-temp Violet solution Col temp Lighter pink solution (refer to test tube in room temp.)

Cold temp Lighter pink solution (refer to test tube in room temp.)

Hot temp Violet solution

Room temp Pink solution Room temp Pink solution

red. However, when NaH2PO4 was added to the solution, the color of the solution changes to a

yellowish one which can be explained by using Le Chatelier’s Principle.

Le Chatelier’s Principle states that the concentration stress of an added reactant or product is

relieved by net reaction in the direction that consumes the added substance furthermore, the

concentration stress of a removed reactant or product is relieved by net reaction in the direction that

replenishes the removed substance.

As FeSCN2+ reacts with KSCN, the concentration of stress added of SCN-shifts the equilibrium

from left to right and the red color gets darker. Similarly, as Fe(NO3)3 was dropped in the solution,

the concentration stress of added FE3+ is relieved by net reaction from left to right which consumes

the Fe3+ and increase the concentration of FeSCN2+ making the solution change to a darker shade of

red (McMurry& Fay, 2008).

However, McMurry and Fay (2008) added that as Le Chatelier’s Principle is a handy rule for

predicting chanhes in the composition of an equilibrium mixture, it doesn’t explain why those

changes occur. So in order to thoroughly picture out how the principle works, we can use the

mathematical relationshipaA + bB ->cC + dD specifically the Law of Mass Action which is represented

by the following expression:

Keq = [C]c [D]d / [A]a [B]b

Wherein the quotient (Qc) is equal to Keq presenting a system at equilibrium.

When the equilibrium is disturbed by increasing the concentration of reactant, the denominator

of the equilibrium constant expression increases yielding a quotient (Qc) less than the Keq, so for the

system to move to a new state of equilibrium, the Qc must increase by increasing the numerator

(product) and decreasing the denominator.

In line with the experiment, as the FeSCN2+ was added with SCN, the concentration of the

reactant increases, making the quotient less than the supposedly Keq. So the tendency of the system

is to increase the quotient by raising the amount product and reducing the reactants so the reaction

will have a forward shift, thereby increasing the quotient to adjust the system to equilibrium.

The constants Kc (Keq) and Kp for the general gas-phase reaction aA + bB ->cC + dD are related

because the pressure of each component in a mixture of ideal gases is directly proportional to its

molecular concentration.

The equilibrium equation for Kp is therefore given by:

Kp= (PC) c (P D) d = [C] c [D] d * RT (c+d) - –a+b)

(PA)a(PB)b [A] a [B] b

Therefore, Kp= Kc (RT)(An)

Because gas pressures are easily measured, equilibrium equations for gas-phase reactions are

often written using partial pressures. The partial pressure of a gas in a mixture is the pressure the gas

would exert if it were the only one present and is independent of the partial pressures of the other

gases in the mixture.

In relation to this, Table 4.3 shows the variations of color as the pressure in the syringe was

increased showing the effect of pressure on chemical equilibrium.

It was observed that as the plunger was pushed until the 4cc mark, the brown gas turned to a

lighter shade of brown.

As Le Chatelier’s Principle state, An increase in pressure by reducing the volume will bring about

net reaction in the direction that decreases the number of moles of gas while a decrease in pressure

will bring about net reaction in the direction that increases the number of moles of gas.

2 no2 -> N2O4

Increasing the pressure on this equilibrium system will result in the equilibrium position shifting

to reduce the pressure, that is, to the side that has the least number of gas particles.There are 2 gas

particles on the left hand side of the reaction and 1 gas particle on the right hand side of the

reaction.Increasing the pressure on this system results in the equilibrium position moving to the right,

consuming NO2(g) and producing more N2O4(g). The system will become a lighter red-brown

colour.Thus, increase in pressure favors the backward reaction and vice versa.

Meanwhile, Table 4.4 shows the effect of temperature on the chemical equilibrium. As the

prepared solution was put on the cold water bath, the intial pink solution turned to a lighter shade of pink

and the test tube that was placed on the hot water bath turned to violet solution. Afterwards, as the temperature

was shifted from hot to cold temperature, and cold to hot temperature, the color of the solution also shifted from

light pink to violet and vice versa.

For an endothermic reaction such as the reaction of the prepared solution of CoCl2, heat is absorbed by

the reaction in the forward direction. The equilibrium therefore shifts to the product side at the higher

temperature which means that equilibrium increases with increasing temperature making the pink solution turned

to a violet one. And the exothermic reaction of CoCl2 wherein heat is released made possible the color alteration

of the pink solution to a lighter shade.

The equilibrium constant for an exothermic reaction decreases as the temperature increases wherein the

energy can be considered as a product of the reaction. Meanwhile, in an endothermic reaction, the equilibrium

constant is directly proportional to the temperature which means that as Kc increases, the temperature also

increases and the energy can be considered as a reactant of the reaction.

You can predict the way in which Kc depends on temperature by using Le Chatelier’s Principle. As it says, if

heat is added to an equilibrium mixture thus increasing its temperature, net recation occurs in the direction that

relieves the stress of the added heat.

To generalize the idea, in an Endothermic Equilibrium Systems, Increasing the temperature of the

equilibrium system will shift the equilibrium position to the side that does not include the energy term in order to

reduce the temperature, that is to the rightmakinf the color of the solution darker.Whereas in an Exothermic

Equilibrium Systems,Increasing the temperature of this equilibrium system shifts the equilibrium position to the

left, consuming some of the energy and products to produce more reactants making the color of the solution

lighter.

VI. CONCLUSION

Chemical equilibrium is the state reached when the concentrations of reactants and products

remain constant over time.

There are some factors which instigate these changes to disturb a system at equilibrium namely,

the concentration, pressure and temperature.

The effect of any change in the reaction condition of a system in equilibrium can be described

using the Le Chatelier’s Principle.

On the effect of concentration on chemical equilibrium, Le Chatelier’s Principle states that the

concentration stress of an added reactant or product is relieved by net reaction in the direction that

consumes the added substance furthermore, the concentration stress of a removed reactant or product is

relieved by net reaction in the direction that replenishes the removed substance which is supported by

the variations on the color of the FeSCN2+ solution as they were added with different substances.

Meanwhile, LeChatelier’s Principle states that An increase in pressure by reducing the volume

will bring about net reaction in the direction that decreases the number of moles of gas while a decrease

in pressure will bring about net reaction in the direction that increases the number of moles of gas which

explains why the color of the gas changes from brown to a lighter shade .

Furthermore, he described the effect of temperature on the equilibrium constant, in two ways for an

Endothermic reaction, Increasing the temperature of the equilibrium system will shift the equilibrium

position to the side that does not include the energy term in order to reduce the temperature, that is to

the right making the color of the solution darker which explains why the test tube that was put on the hot

water bath became violet. Whereas in an Exothermic reaction, Increasing the temperature of this

equilibrium system shifts the equilibrium position to the left, consuming some of the energy and products

to produce more reactants making the color of the solution lighter and which explains why the solution

placed on the cold water bath turns light pink.

VII. LITERATURE CITED

Lower, S.K. 2001.Chemical Equilibrium: A chemical 1 reference text. California: NP. McMurry, J.E & Fay, R.C. 2008.Chemistry. (5thed.).Pof pub: Prentice Hall.Atkin, P. & de Paula, J. 2010. Physical Cehmistry for the Life Sciences.(2nded.). Place of pub:

Oxford University Press. http://www.ausetute.com.au/lechatsp.html