CHEMICAL EQUILIBRIUM

Katherine Borlongan Marilag Sadicon HAB 1

Objectives

to evaluate how the change in concentration and temperature affect the equilibrium system

to explain the effect of these changes in the equilibrium system

to interpret the results based on Le Chatelier’s Principle

Chemical Equilibrium Chemical reactions

The state in which the forward reaction (formation of products) is happening at the same rate as the backward reaction (formation of reactants).

A + B C The reaction continues but the concentrations

of the reactants and products stay the same.

H2 + I2 2HI 2HI H2 + I2

Source: http://www.wellesley.edu/Chemistry/equil1.gif

Le Chatelier’s Principle If a system at equilibrium is subjected to a

change in pressure, temperature, or the number of moles of a substance, there will be a shift in the direction such that the effect is reduced and equilibrium is attained once again.

Three (3) Factors that can disturb chemical equilibrium:

1.Change in reactant or product concentrations 2.Change in the pressure 3.Change in temperature

Methodology20 drops of 1 M Fe(NO3)3 solution

20 drops of 1 M KCNS solution

10 mL

7 mL of water

Color of mixture?

110 drops of

distilled water

210 drops 1 M

Fe(NO3)3

310 drops 1 M

KCNS

410 drops 1 M

KCL

510 drops 1 M

AgNO3

6 1 pinch of NaF

7 Ice bath

8 Heat

Data and ResultsReagent/ Treatment

Observation Direction of shift

0.1 M Fe(NO3)3 Darker Forward

0.1 M KCNS Darker Forward

0.1 M KCL Lighter Backward

0.1 M AgNO3 Milky white Backward

Pinch of NaF Milky white Backward

Increase in temp.

Lighter Backward

Decrease in temp.

Darker Forward

Discussion Fe 3+ + CNS- FeCNS 2+

Orange colorless blood-red

Effect of concentration changes Test tube 2 and 3 An increase in concentration of the reactant To reduce amount of reactant (and attain

equilibrium), shift towards formation of more product (forward direction)

A forward reaction resulted in the creation of more FeSCN2+ making solution darker

Test tube 4 KCl dissociates into K+ and Cl- ions K+ ions also form a complex with CNS

Fe(CNS)3 + 3KCl FeCl 3+ 3KCNS

Decrease in amount of reactant CNS-

To increase amount of reactant (to attain equilibrium), the product FeSCN2+ is broken down (reverse shift)

Breakdown of product turns solution lighter  Light brown color can also be attributed to FeCl3

molecules Used for industrial applications

Test tube 5 Addition of AgNO3 caused a precipitation

of insoluble compound  AgCNS

Fe(CNS)3+ AgNO3  AgSCN + Fe(NO3)3

Decrease in the amount of CNS- ions as AgCNS was being formed

To increase amount of reactant (to attain equilibrium),the product FeSCN2+ was broken down (reverse shift)

Breakdown of product turns solution lighter

Although Fe(NO3)3, a reactant, was also produced, the effect of the shift towards the breakdown of the FeCNS2+ was favored because the AgCNS formed is solid and insoluble.

Test tube 6

F– reacted with Fe, decreasing iron’s concentration

There was a backward shift- formation of more reactants

Iron (III) fluoride commercially used for the production of ceramics,

it comes in a pale, milky color

Effect of temperature changes Equilibrium constant changes with changes in

temperature. Treat heat, H, as if it were a chemical reagent Endothermic:  Reactants + H product Exothermic: Reactants product + H

When the temperature is increased, it is as if we have added a reactant, or a product, to the system of equilibrium. The equilibrium shifts in the direction that consumes the excess reactant(or product), namely heat.

Endothermic: like adding more reactants Reduce amount of reactant and attain equilibrium

by shifting towards formation of more products Increase in temperature causes the shift to the

right

Exothermic: like adding more products Reduce amount of the product and attain

equilibrium by shifting towards formation of more reactant

Heating test tube 7 resulted in a solution of a lighter color, indicating a shift towards the production of reactants.

Placing test tube 8 in an ice bath produced a darker color of solution, indicating a shift towards production of products. Hence, it is an exothermic reaction

Guide Questions Explain your observations on the basis of Le Chatelier’s

Principle.  According to Le Chatelier’s Principle, if a system is at

equilibrium and we add a substance, the reaction will shift so as to reestablish equilibrium by consuming part of the added substance.

If reactants are added or products are removed, the system will reestablish equilibrium by moving in the forward direction and converting the reactants into products. Such is the case in test tubes 2, 3, and 8 (if you treat heat as a reactant). On the other hand, if the concentration of the reactants is decreased, shifting the reaction backwards and decomposing the products into reactants reestablishes equilibrium. Such as demonstrated by test tubes four, five and six and seven (if you treat heat as a reactant).

Which species (ions) in the added reagents are effective in altering the state of the system?

Fe3+ and CNS are effective.

Is the reaction endothermic or exothermic? Why? The reaction is exothermic, because the reaction

shifts backward (a lighter solution was obtained) with an increase in temperature, indicating that the reaction is heat releasing.

Conclusions and Recommendations Le Chatelier’s principle has

been demonstrated in this experiment. An increase in the concentration of the

reactant or a decrease in the concentration of the product shifts the direction of the reaction towards the production of more products to return to the equilibrium position.

The opposite, a decrease in the concentration of the reactants or an increase in the concentration of the product shifts the equilibrium position towards the production of more reactants.

If temperature is treated as a reactant, in an exothermic reaction, an increase in temperature shifts the direction towards the production of more reactants.

A decrease in temperature in an exothermic reaction causes a shift towards the production of products, in order to reestablish equilibrium.

References

Lower, S. (n.d.) Chemical equilibrium: A Chem1 reference text Retrieved from http://www.chem1.com/acad/pdf/chemeq.pdf