Complex Formation and Precipitation free essay sample

Complex Formation and Precipitation. INTRODUCTION Precipitation is the formation of a solid in a solution as the result of either a chemical reaction, or supersaturating a solution with a salt resulting in solid material collecting on the bottom of the beaker (Housecroft amp; Constable, 2006). When the solution has been supersaturated by a compound and no more material can be supported by the solution, it considered as precipitate. Commonly, the solid will fall out of the solute phase and sinks to the bottom part of the solution. However, the solid may float on top of the solution or suspended within the solution if the precipitate is less than the aqueous solution. On the other hand, complex formation is the soluble compound that can increase the solubility of an insoluble salt form in aqueos solution (Housecroft amp; Constable, 2006). They are formed from a cation such as water or ammonia that has pair of electrons that become attracted to the positively charged cation. In terms of the solubility, the formation of a complex lowers the concentration of the cation in the solution. In this experiment, there are 3 parts need to be done which is complex formation, weak acids and bases and precipitation reaction. During complex formation experiment, student needs to observe the different reaction occur when adding ammonia or sodium hydroxide in small quantity and in excess. At the second part, student needs to measure the pH of different solution by using universal indicator paper and determine whether each solution is acidic, basic or neutral. Lastly, the precipitation reaction is observed by using phosphates, chlorides and silver halides groups. AIMS The main objective of this experiment was to identify certain inorganic reaction types and the differences in solubility of silver halides. Thus, a series of inorganic reactions was carried out as an introduction to complex formation and precipitation. All the reaction in this experiment involved alteration of equilibrium by changing concentration to make precipitation to form or dissolve. Other than that, weak acid and bases of the solution can be identified and students able to write the balanced equations for inorganic reactions. PRE-LAB QUESTIONS 1. Precipitation reaction is a type of reaction that forms a solid in a solution as a result of chemical reaction or supersaturating a solution. This reaction involves changing the soluble ions into insoluble compound and as a result precipitate is formed when the solution has been saturated by a compound (Heitz, 2005). 2. Inorganic complex consists of a central ion or atom (usually metallic) that is surrounded by anions or molecules bound to the central ion/metal via chemical bonds that range in strength from weak to strong (Heitz, 2005). 3. Trigonal planar, tetrahedral and octahedral. METHODS * PART A: COMPLEX FORMATION (AMMONIA COMPLEXES) . 0. 5 cm3 of dilute zinc nitrate, cupric nitrate and lead nitrate solutions was poured into 3 different test tubes. 2. 4 M of ammonia solution was added dropwise in each test tube until each solution turn alkaline (test with pH paper). 3. Ammonia solution was added more and observes if any precipitate dissolves. 4. The experiment was repeated by re placing the ammonia solution with 2 M of sodium hydroxide. * PART A: COMPLEX FORMATION (IODIDE COMPLEXES) 1. Saturated solution of potassium hydroxide was added dropwise to 0. 5 cm3 of 0. 1 M mercuric nitrate solution until the precipitate of HgI2 just dissolves by forming [HgI4]2-. . Saturated solution of potassium hydroxide was added dropwise to 5 drops of 0. 1 M lead nitrate solution until the precipitate of PbI2 just dissolves by forming [PbI4]2-. 3. The solution of [PbI4]2- was added to about 10 times its volume of distilled water in a medium test tube. 4. The amounts of I- required to dissolve the precipitates formed was compared between Hg (II) and Pb (II). * PART B: WEAK ACIDS AND BASES 1. The pH of distilled water, sodium chloride, aluminium sulphate, ammonium chloride, trisodium phosphate and sodium acetate was measured by using universal indicator paper. 2. Each solution was determined whether it is acidic, basic or neutral. 3. The main ions present in solution and acid-base equation for each solution was recorded. * PART C: PRECIPITATION REACTIONS (PHOSPHATES) 1. 1 cm3 of each of potassium chloride, manganous chloride, magnesium chloride, calcium chloride and cupric sulphate was placed into separate test tubes. 2. 1 cm3 of trisodium phosphate solution was added to each test tube and comment on the possible reasons for the formation of a precipitate. 3. A few drops of 1 M HCl were added to any precipitate formed, shake carefully and any observations were recorded. PART C: PRECIPITATION REACTIONS (CHLORIDES) 1. 1 cm3 of concentrated hydrochloric acid was added to 2 cm3 of saturated sodium chloride solution. 2. The precipitate was allowed to settle before the liquid was pouring off into the mixture and add water. * PART C: PRECIPITATION REACTIONS (SILVER HALIDES) 1. Place 10 drops of sodium chloride, sodium bromide and sodium iodide solut ion into three separate test-tubes. 2. 3 drops of silver nitrate solution was added to each test tube. Then, the precipitate was allowed to settle and pour off the liquid. 3. The solubility of any precipitates was tested in 4 M ammonia solution. . Steps 1 and 2 of this experiment were repeated. 5. The solubility of any precipitated was tested in concentrated ammonia solution (density 0. 880 g cm-3) RESULTS AND DISCUSSION * PART A: COMPLEX FORMATION (AMMONIA COMPLEXES) Table 1: Result of adding ammonia solution to form Ammonia Complex Reactants| Observation| Equation| Zinc nitrate and ammonia solution| * Turn red litmus paper into blue. * White precipitate was formed with cloudy solution. | Zn2+aq+2OH-aq-ZnOH2S| Zinc nitrate and ammonia solution (excess)| * Turn red litmus paper into blue. Clear solution is obtained with all precipitate dissolved in excess ammonia. | Zn(OH)2(S)+4HN3 Zn(NH3)4(aq)2++2OH-aq| Cuprix nitrate and ammonia solution| * Turn red litmus paper into blue. * Blue precipitate was formed. | Cu2+aq+2OH-aq? Cu(OH)2(S)| Cuprix nitrate and ammonia solution (excess)| * Turn red litmus paper into blue. * Clear blue solution is obtained with all precipitate dissolved in excess ammonia. | Cu(OH)2(S)+4HN3Cu(NH3)4(aq)2++2OH-aq| Lead nitrate and ammonia solution| * Turn red litmus paper into blue. * White precipitate was formed with clear solution| Pb2+aq+2OH-aq? Pb(OH)2(S)| Lead nitrate and ammonia solution (excess)| * Turn red litmus paper into blue. * Clear solution is obtained with all precipitate dissolved in excess ammonia. | Pb(OH)2(s) + 4NH3(aq) ? [Pb(NH3)4](aq) + 2OH-(aq)| DISCUSSION| Zinc nitrate, cuprix nitrate and lead nitrate produce white precipitate when reacted with 4 M ammonia and dissolve in excess ammonia solution because they were partially hydrolyzed in water. The solubility of all three precipitate was enhanced by excess ammonia ions although they were insoluble in water. Ammonia acts as ligands where OH group has been replaced with NH3 to form a complex ion. Table 2: Result of adding sodium hydroxide solution to form Ammonia Complex Reactants| Observation| Equation| Zinc nitrate and sodium hydroxide solution| * Turn red litmus paper into blue. * White precipitate was formed with cloudy solution. | Zn(NO3)2aq+2NaOHaqZn(OH)2(S)+2NaNO3aq| Zinc nitrate and sodium hydroxide solution (excess)| * Turn red litmus paper into blu e. * Clear solution is obtained with all precipitate dissolved in excess sodium hydroxide. | Zn(OH)2(aq)+2OH-aqZn(OH)4(aq)2-| Cuprix nitrate and sodium hydroxide solution| * Turn red litmus paper into blue. Blue precipitate was formed with cloudy solution. | CuNO32aq+2NaOHaqCu(OH)2(S)+2NaNO3aq| Cuprix nitrate and sodium hydroxide solution (excess)| * Turn red litmus paper into blue. * Precipitate did not dissolve in excess sodium hydroxide result in dark blue solution. | CuNO32aq+2NaOHaqCu(OH)2(S)+2NaNO3aq| Lead nitrate and sodium hydroxide solution| * Turn red litmus paper into blue. * White precipitate was formed with cloudy solution. | Pb(NO3)2aq+2NaOHaqPb(OH)2(S)+2NaNO3aq| Lead nitrate and sodium hydroxide solution (excess)| * Turn red litmus paper into blue. Clear solution is obtained with all precipitate dissolved in excess sodium hydroxide. | Pb(OH)2(aq)+2OH-aqPb(OH)4(aq)2-| DISCUSSION| The precipitate was formed when sodium hydroxide were reacted with all the three solution and were dissolved when the amount of sodium hydroxide was excess. This shows that the solubility was enhanced by an excess of sodium hydroxide ions as more sodium hydroxide ions can bind with each of the solutions. | * PART A: COMPLEX FORMATION (IODIDE COMPLEXES) Table 3: Results of adding Potassium Iodide (KI) to form Iodide Complex Reactants| Observation| Equation| Mercuric nitrate solution and saturated potassium iodide| * The solution turns to orange colour at the first drop. * The clear solution is obtained after 6 drops of saturated potassium iodide. | * First dropHg2+aq+2I-aq? HgI2(S) * After six dropsHgI2(s) + 2I-(aq) HgI42-(aq)| Lead nitrate solution and saturated potassium iodide| * The solution turns to bright yellow colour at the first drop. * The clear solution is obtained after 15 drops of saturated potassium iodide. * After 10 times of its volume of distilled water (80 ml) was added, the precipitate appeared again| * First dropPb2+aq+2I-aq? PbI2(S) * After fifteen dropsPbI2(S)+2I-aq? PbI42-(aq) * Adding waterPbI42-(aq) + H2O(l) ? PbI2(s) + 2I-(aq)| DISCUSSION| Mercuric nitrate solution shows lower stability constant as compared to lead nitrate solution because only six drops of saturated potassium iodide is needed to dissolve the precipitate formed (Lead nitrate need fifteen drops to dissolve the precipitate). Thus, lead nitrate have higher relative stability constant as compared to the mercuric nitrate solution. | * PART B: WEAK ACID AND BASES Table 4: Reaction, Ions and their Relationship in Weak Acid/Bases Reactants| Observation| Equation| Comment| Distilled water| * pH = 6 * Acidic solution * Main ions present:H+,OH- | * H2Ol ? H3O+aq +OH-aq| * Distilled water supposed to have neutral pH condition which is 7. * The different in the result might due to the presence of the other ions in the distilled water. | Sodium chloride| * pH = 5 * Acidic solution * Main ions present:Na+, Cl-,H+,OH-| * NaClaq ? Na+aq+ Cl-aq| * Sodium chloride supposed to have neutral pH condition because it is neutral salt. * The different in the result might due to the presence of contaminant in the solution used during experiment. Aluminium sulphate| * pH = 2 * Acidic solution * Main ions present:AlH2O62+,SO42-| * AlH2O62+(aq)+H2Oaq? AlH2O3OH(aq)+H+aq| * Due to the presence of H+ ion after dissociation, the solution is acidic and one of the strongest acid as the pH value is lower. | Ammonium chloride| * pH = 5 * Acidic solution * Main ions present:NH4+,Cl-| * NH4+(aq)+H2Ol? H3O+aq+NH3aq| * The pH value shows almost to neutral value as the present H3O+aq r esult to acid properties and NH3aq gave result to basic properties. Trisodium phosphate| * pH = 12 * Basic solution * Main ions present:Na+,PO43-| * HPO42-aq+H2Ol? H2PO4+aq+OH-aq| * Due to the presence of OH- ion after dissociation, the solution is basic and one of the strongest basic as the pH value is higher. | Sodium acetate| * pH = 9 * Basic solution * Main ions present:Na+,CH3COO-| * CH3COO-aq+H2Ol? CH3COOHaq+OH-aq| * Due to the presence of OH- ion after dissociation, the solution has basic properties solution. | * PART C: PRECIPITATION REACTIONS (PHOSPHATES) Table 5: Results of adding trisodium phosphate (Na3PO4) and HCl to different solution. Reactants| Observation| Equation| Potassium chloride and trisodium phosphate solution| * White precipitate were formed | Na3PO4(aq) ? 3Na+(aq) + PO43-(aq)3K+(aq) + PO43-(aq) ? K3PO4(aq)| Potassium chloride, trisodium phosphate and hydrochloric acid| * All precipitate was dissolved and solution is soluble in HCl. | K3PO4S+HCl(aq)? 3KCl(aq)+H3PO4aq| Manganous chloride and trisodium phosphate solution| * White precipitate was formed with cloudy solution. | Na3PO4(aq) ? 3Na+(aq) + PO43-(aq)3Mn2+aq+2PO43-aq? Mn3PO42S| Manganous chloride, trisodium phosphate and hydrochloric acid| * All precipitate was dissolved and solution is soluble in HCl. | Mn3PO42S+6HCl(aq)? 3MnCl2(aq)+2H3PO4aq| Magnesium chloride and trisodium phosphate solution| * White precipitate was formed with cloudy solution. | Na3PO4(aq) ? 3Na+(aq) + PO43-(aq)3Mg2+aq+2PO43-aq? Mg3PO42S| Magnesium chloride, trisodium phosphate and hydrochloric acid| * All precipitate was dissolved and solution is soluble in HCl. | Mg3PO42S+6HCl(aq)? 3MgCl2(aq)+2H3PO4aq| Calcium chloride and trisodium phosphate solution| * White precipitate was formed with cloudy solution. Na3PO4(aq) ? 3Na+(aq) + PO43-(aq)3Ca2+aq+2PO43-aq? Ca3PO42S| Calcium chloride, trisodium phosphate and hydrochloric acid| * All precipitate was dissolved and solution is soluble in HCl. | Ca3PO42S+6HCl(aq)? 3CaCl2(aq)+2H3PO4aq| Cupric sulphate and trisodium phosphate solution| * Blue precipitate was formed with cloudy solution. | 3Cu2+(aq) + 2PO43-(aq) ? Cu3(PO4)2(s)| Cupric sulfate, trisodium phosphate and hydrochloric acid| * All precipitate was dissolved and the solution become soluble (blue in colour)| Cu3(PO4)2(s) + 6HCl(aq) ? 3CuCl2(aq)+ 2 H3PO4(aq)| DISCUSSION| Potassium chloride, manganous chloride, magnesium chloride, calcium chloride and cupric sulphate formed precipitate when reacted with trisodium phosphate solution. All the precipitate then dissolve when hydrochloric acid was added into the mixture to produce metal phosphates and sodium chloride. | * PART C: PRECIPITATION REACTIONS (CHLORIDES) Table 6: Results of adding Concentrated HCl and H2O to saturated NaCl Solution Reactants| Observation| Equation| Potassium chloride and trisodium phosphate solution| * White precipitate were formed | Na+aq+ Cl-aq? NaClsHCl(aq) ? H+(aq) + Cl-(aq)| Potassium chloride, trisodium phosphate solution and water| * White precipitate formed was completely dissolved. | NaCl(s) + H2O(l) ? NaOH(aq) + HCl(aq)| DISCUSSION| The white precipitated was formed because the product, NaCl which is a neutral common salt was white in colour. When water was added, it elimininates the excess Cl- and H+ ions, favoured the backward reaction and thus NaCl is soluble in water. As a result white precipitate dissolved and colourless solution formed. | * PART C: PRECIPITATION REACTIONS (SILVER HALIDES) Table 7: Results of adding AgNO3 and NH3 solution to Sodium Halides Reactants| Observation| Equation| Comment| Sodium chloride and silver nitrate solution| White precipitate was formed with cloudy solution. | AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++Cl-? AgCl (s)| The white precipitate is AgCl because it does not dissolve in water. | Sodium chloride, silver nitrate and ammonia solution| Precipitate dissolved completely to give a colourless solution| AgCl (s)+2NH3(aq)? AgNH32+ (aq)+Cl-(aq)| When AgCl was reacted with ammonia solution, AgNH32+was formed and it was soluble in water as the NH3 ions was able to bind with water molecules. Sodium bromide and silver nitrate solution| Creamy precipitate was formed with cloudy solution. | AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++Br-? Br| AgNH32+is soluble in water as it can bind with water molecules. | Sodium bromide, silver nitrate and ammonia solution| Precipitate was almost unchanged because it did not dissolve completely (partially) in ammonia solutio n| AgBr (s)+2NH3(aq)? AgNH32+ (aq)+Br-(aq)| This precipitate partially dissolves in 4 M ammonia solution. | Sodium iodide and silver nitrate solution| Pale yellow precipitate was formed with cloudy solution. | AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++IAgI (s)| The reaction of NaI with silver nitrate forming a precipitate of AgI| Sodium iodide, silver nitrate and ammonia solution| Precipitate did not dissolve in the mixture of ammonia solution| AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++IAgI (s)| This compound was very insoluble in water, so it will not bind with water molecules, even if reacted with ammonia solution| Table 8: Results of adding AgNO3 and concentrated NH3 solution to Sodium Halides Reactants| Observation| Equation| Comment| Sodium chloride and silver nitrate solution| White precipitate was formed with cloudy solution. | AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++Cl-? AgCl (s)| The white precipitate is AgCl because it does not dissolve in water. | Sodium chloride, silver nitrate and concentratedammonia solution| Precipitate dissolved completely to give a colourless solution| AgCl (s)+2NH3(aq)? AgNH32+ (aq)+Cl-(aq)| When AgCl was reacted with concentrated ammonia solution, AgNH32+was formed and it was soluble in water as the NH3 ions was able to bind with water molecules. | Sodium bromide and silver nitrate solution| Creamy precipitate was formed with cloudy solution. | AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++Br-? Br| AgNH32+is soluble in water as it can bind with water molecules. Sodium bromide, silver nitrate and ammonia solution| Precipitate dissolved completely to give a colourless solution| AgBr (s)+2NH3(aq)? AgNH32+ (aq)+Br-(aq)| By adding concentrated ammonia solution, the precipitate (AgBr) dissolved completely. | Sodium iodide and silver nitrate solution| Pale yellow precipitate was formed with cloudy solution. | AgN O3(aq) ? Ag+(aq) + NO3-(aq) Ag++IAgI (s)| The reaction of NaI with silver nitrate forming a precipitate of AgI| Sodium iodide, silver nitrate and ammonia solution| Precipitate did not dissolve in the mixture of ammonia solution| AgNO3(aq) ? Ag+(aq) + NO3-(aq) Ag++IAgI (s)| Adding either diluted or concentrated NH3 solution, did not make the precipitate (AgI) became soluble because it will not dissolve. Thus, it can be said that solubility of silver halides decreases down the group. | QUESTIONS 1. The relative stability constants of iodide complex of Hg (II) is less than the iodide complex of Pb(II). This is because from the result obtained, only 6 drops of iodide were required to dissolve the precipitates formed by the mercuric nitrate, Hg (II) as compared to lead nitrate, Pb(II) that required 15 drops to dissolve the precipitate. Thus, iodide complex of Pb(II) is more stable than iodide complex of Hg(II) because more drops of iodide were required to dissolve lead nitrate, Pb(II) than mercuric nitrate, Hg(II). 2. Commonly, using a pH meter and an indicator paper during the experiment would result in different pH value obtained. This is because pH meter is more precise in giving pH values up to 2 decimal places while indicator paper only gives an approximation of pH based on colour with reference to chart. Thus, due to difficulty to differentiate the colours observed as referred to chart, indicator paper is not the best method to determine the accurate pH value of the mixture. 3. Equation of the reaction: NaCl Na+ + Cl- The precipitate formed was sodium salt crystal (NaCl). According to Le Chatelier’s principle, this change in reaction acts to counteract the increase of Cl- concentration (Silberberg, 2006). When concentrated HCl was added, concentration of chloride ion (Cl) increases. Thus, equilibrium will shift to the left. Therefore, reverse reaction was favoured result in more solid salt precipitate (NaCl) was formed. 4. In the case of silver halides, the chloride is the most soluble in ammonia and the iodide is the least while for the case of calcium halides the reverse trend is seen. This is due to the solubility of the group 2 sulphates which decreased down the trend. Thus, silver and metal halides were soluble in ammonia as they formed complex ions which were soluble. CONCLUSION In conclusion, the complex formation and precipitation reactions were able to establish thus met the aims of this experiment. Thus, it can be concluded that ammonia and hydroxide ion can act as a Lewis base since a complex ion was formed when a metal cation is covalently bonded to ligands. Other than that, weak acid and base will dissociates partially in water while strong acid and base dissociate completely in water. Finally, the formation of the precipitate (product) can be controlled by controlling the pH of the reaction. REFERENCES Heitz. , M(2005) General Chemistry: An Integrated Approach, Chapter 16: More Equilibria in Aqueous Solutions: Slightly Soluble Salts and Complex Ions, 4th Edition, Prentice Hall, Inc. , State University of New York. pp. 715-718 Housecroft, C. E. amp; Constable, E. C. (2006), Introduction to spectroscopy, in Chemistry, 3rd Edition, Pearson Education, England. pp. 951-970 Silberberg, M. S. (2006), Chemistry: The Molecular Nature of Matter and Change, 4th Edition, Mc Graw Hill International Edition, New York, pp. 1013-1019.