D.M. TAN1 AND P.B. ALEGRO2
1DEPARTMENT OF MINING, METALLURGICAL, AND MATERIALS ENGINEERING, COLLEGE OF ENGINEERING 2 INSTITUTE OF CHEMISTRY, COLLEGE OF SCIENCE
UNIVERSITY OF THE PHILIPPINES, DILIMAN QUEZON CITY, PHILIPPINES RECEIVED JANUARY 15, 2013
RESULTS AND DISCUSSION
A. Iron- Silver Equilibrium
The first part of the experimentation focuses in the iron-silver system. Silver nitrate (AgNO3) was added to ferrous sulfate (FeSO4) shown in this equation: 2AgNO3(l) + FeSO4(l) Ag2SO4-(aq)+ Fe(NO3)2(aq)
The additon of AgNO3to FeSO4 equilibrium had an effect on the equlibria. Though no precipitate was formed after just mixing the solutions, a precipitate formed after centrifugation. This was done to achieve separation of silver precipitate to the supernate.
The silver precipitate is formed due to the unsolubility of AgNO3 which displaces the chemical equilibrium of FeSO4 towards the silver compound. This is an application of Le-Chatelier fundament which states that an equilibrium will attempt to shift in a direction that will counteract a stress that is placed on it.
The supernate was then tested for the presence of three ions namely, Fe2+, Fe3+, and Ag+. In the test for Fe2+, he supernate was added with K3Fe(CN)6 that resulted to net ionic equation: Fe2+(aq) + K3Fe(CN)6(aq) KFe[Fe(CN)6](s)
The precipitate which is KFe[Fe(CN)6](s) is a prussian blue precipitae which indicates the presence of Fe2+ ions. Next, for Fe3+, KCNS was added in the supernate shown in this net ionic equation: Fe3+(aq) + SCN-(aq) Fe(SCN)2+(aq)
A blood red solution was achieved which indicates the presence of Fe3+ ions. Lastly, for the Ag+ test, the HCl was added to the supernate shown in the ionic equation below: Ag+(aq) + Cl-(aq) AgCl(s)
A white precipitate formed which indicates the presence ogf Ag+ ions.
Since the reactants and the products are both present at the same time, the system has reached its equilibrium. This means though products are produced, the reverse reaction still proceeds. For that it could be concluded that the range of Keq is between 10-2 and 102, significant amounts of both product and reactant will remain on the reaction and these substances, depending on the particular chemical reaction, can be easily be detected via qualitative test.
B. Copper-Ammonia Equilibrium
The second part of the experimentation focuses on the copper(II)-ammonia system. Ammonia (NH3) was added to cupric sulfate as shown in this chemical equation: (with limited NH3) CuSO4(aq) + NH3(aq) Cu(OH)+(s)
Cu(OH), a light blue precipitate formed upon adding smaller amount of ammonia. The precipitate formed because the concentration of the reactants were not equal thus no equilibrium attained. But with excess amount of ammonia added, as shown in this equation: (with excess NH3) CuSO4(aq) + 4NH3(aq) Cu(NH3)42+(aq)
The precipitate would dissolve due to the formation of soluble complex compound. The reaction was forced to go on completion which resulted to a deep blue complex.
When HCl was added, the production of NH4+ is triggered. The decolorization of the solution is due to the system at equilibrium given by the reaction below: NH3(aq)+H+(aq)⇆NH4+(aq)
The addition of HCl removes the available NH3 that will lead to the formation of the precipitate, copper (II) hydroxide and the complex ion, tetraamminecopper(II). Therefore, by the addition of HCl, the general chemical equilibrium is favored in the direction of the reverse reaction.
The amount of HCl added to shift the equilibrium to the reverse direction is 1/3 less than the amount of NH3 needed to complete the equilibrium (HCl- 10 drops, NH3-15 drops). This indicates that it is easier for HCl to shift the direction of the equilibrium in reverse than for NH3 which needed more amount to complete the reaction and attain...