Chemical Equilibrium A. Arteta Department of Chemical Engineering, College of Engineering University of the Philippines, Diliman, Quezon City Date Performed: April 18, 2013 Instructor’s Name: Maro Peña
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RESULTS AND DISCUSSION
The experiment was aimed to study the effect of temperature and concentration of involved species on the equilibrium. equilibrium. The first part of the experiment involved the equilibrium in iron (III) – silver ions system. The chemical equation for the system is given below. FeSO4(aq) + AgNO3(aq) ⇌ Fe(NO3)3(aq) + 2Ag(s) + SO42-(aq) Fe2+(aq)
+
Ag+
(aq)
⇌
Fe3+(aq)
+ Ag(s)
The reaction, as shown in the equation, yielded a white precipitate (Ag). To separate the precipitate from the supernate, the solution was first centrifuged to gather the precipitate at the bottom of the container. Afterwards, the supernate was separated from the precipitate by means of decantation. The supernate was then tested for the presence of Fe 2+, Fe3+, and Ag+ ions by reacting it with test reagents namely K3Fe(CN)6, KCNS, and HCl. Ions Fe2+
Fe3+
Reagent Used K3Fe(CN)6
KCNS
solution HCl White Positive precipitate Table 1. Data for iron (II) – silver ions system Ag+
Table 1 implies that the supernate contains all of the mentioned ions. The presence of Fe2+, Fe3+, and Ag+ were confirmed by the formation of Prussian blue precipitate, blood red solution and white precipitate, respectively. The equations for the reactions were listed below. For Fe2+: 3 Fe2+(aq) + 2 Fe(CN)63-(aq) → Fe3[(CN)6]2(s) For Fe3+: Fe3+(aq) + SCN-(aq) → FeSCN2+(aq) For Ag+: Ag+(aq)+ Cl-(aq) → AgCl(s) Meanwhile, the range of K eq for experiment would be from 10 -2 to 10 2. Keq =
this
Observation
Result
Prussian blue precipitate Blood red
positive
Since Keq = 0.1 and is between 10-2 and 102, it means that significant amounts of both products and reactants will remain in the reaction and these substances can be easily detected by using qualitative tests.
positive
The second part involved the equilibrium of the copper (II) – ammonia system. The
chemical equation for the reaction was shown below.
2 Cr2O42-(aq) + 2 H+(aq) ⇌ Cr2O72-(aq) + H20(l) Cr2O72-(aq) + 2 OH-(aq) ⇌ 2 Cr2O42-(aq) + H20(l)
Cu2+(aq) + 2 NH3(aq) + 2 H2O(l) ⇌ 2 NH4+(aq) + Cu(OH)2(s) The reaction yielded blue-green precipitate, Cu(OH)2(s). The precipitation was formed from the reaction of Cu 2+ to OH-. Addition of excess NH3 to the previously attained system caused the formation of deep blue complex, copper tetramine. Cu(OH)2(s) + 4 NH3(aq) ⇌ [Cu(NH3)4]2- + 2 OHAfterwards, the solution was added with HCl, causing the formation of the initial reactants. This happened because the H + ions from the acid, reacts with NH 3 to form NH4+ as shown in the equation below. NH3(aq) + H+(aq) → NH4+(aq) Meanwhile, it was observed that the system requires an equal of amount of the added acid and base. Addition of 7 drops of NH 3 changed the color of the solution from light blue to deep blue. Conversely, addition of 7 drops of HCl changed the color of the solution from deep blue to light blue. Table 2 summarizes the data gathered in this part of the experiment. Number of Drops Added N/A 7
Color
Initial Light blue Addition of Deep NH3 blue Addition of 7 Light blue HCl Table 2. Data for copper-ammonia system The third part of the experiment involved the equilibrium in chromate –dichromate system and examined how pH affects equilibrium. The chromate solution (yellow) was placed at wells 1-2 and the dichromate solution and the dichromate solution (orange) was placed at wells 3-4.The equations for the reactions are given below.
Well No. 1
Reagent H2SO4
2 3 4 Table 3. system.
Data
Visible Result Yellow to orange NaOH No visible change H2SO4 No visible change NaOH Orange to yellow for chromate-dichromate
Table 3 shows that the addition of acid changed the color of the chromate solution from yellow to orange and the addition of base changed the color of the dichromate solution from orange to yellow. H2SO4 was used to acidify because it is a strong acid and therefore, it completely dissociates into H + and SO42- easily. The added H + would then favor either forward or reverse reaction, depending on how the equilibrium will be attained. Based on the observations in the experiment, it can be said that chromate is stable under acidic conditions and dichromate is stable under basic conditions. The fourth part of the experiment showed how the varying amounts of reactants and products would affect the equilibrium of the iron (III) chloride-thiocyanate system. The equation for the reaction is shown below. FeCl3(aq) + KSCN(aq) ⇌ Fe(SCN)3(aq) + 3 KCl(aq) Fe3+(aq) + 3 SCN-(aq) ⇌ Fe(SCN)3(aq) The table below shows the data gathered in this part of the experiment. Color Fe3+
Deep orange
Shift in Equilibrium To the right
SCNDeep orange To the right NaCl Orange To the left Table 4. Data for irom (III) chloridethiocyanate system.
reaction is endothermic since the heat can be considered as a reactant.
Table 4 tackles the color of the solution after a specific reagent was added to the solution. Fe3+ and SCN- are included in the original reactants’ side, therefore, addition of these ions would result to a change of color from orange to deep orange.
CONCLUSION
On the other hand, the addition of Cl - (NaCl) would react with Fe 3+, forming the complex, FeCl-4. Since the amount of Fe3+ in the reactants’ side was reduced, the reaction would proceed to the reactants’ side. Thus, the color of the solution changed to a lighter shade of orange. The fifth and final part of the experiment involved the equilibrium in cobalt-cobalt chloride system with the ionic equation: Co2+(aq) + 4 Cl-(aq) ⇌ CoCl42The observations gathered in the experiment are shown below. Color Pink
Before addition of HCl After addition of HCl Blue At room temperature Pink At boiling-water Blue temperature Table 5. Data for cobalt (II) ions system. The color of the reactant, Co 2+, is pink and the color of the product, CoCl 42-, is blue. Also, as shown in Table 5, the color of the solution under room temperature is pink (similar to the original color) and turned blue (color of CoCl42-) at boiling-water temperature. This signifies that a higher temperature favored the forward reaction of the system. Hence, the
The experiment was aimed to study the equilibrium states in five different systems. Each part of the experiment shows that in all systems, Le Chatelier’s Principle of Equilibrium can be applied. Applying a stress to the system at equilibrium would either result to a forward reaction or a backward reaction until equilibrium is finally restored. Also, it was shown that certain amounts of reactants and products will remain throughout the reaction. Moreover, it can be concluded that acidity, concentration and temperature are among the factors that affect the equilibrium of a system.
REFERENCES
Petrucci, R. H., Herring, F. G., Madura, J.D. and Bissonnette, C (2011). General Chemistry – Principles and Modern Applications. 10th edition. Pearson Education, Canada. General Chemistry II: Laboratory Manual . June 2011. University of the Philippines, Diliman, Quezon City.
