EXPERIMENT 2 Synthesis of Potassium Tris (1).docx

**Experiment 1: Degree of dissociation of double salts and complex compounds** A double salt is a combination of two salt compounds. A complex salt is a molecular structure that is composed of one or more complex ions. Double salts can give simple ions when added to water. Complex salts do not give...

**Experiment 1: Degree of dissociation of double salts and complex compounds** A double salt is a combination of two salt compounds. A complex salt is a molecular structure that is composed of one or more complex ions. Double salts can give simple ions when added to water. Complex salts do not give simple ions when added to water. A double salt dissociates fully into simple ions in water, whereas a coordination complex dissociates with at least one complex ion in water. Mohr\'s salt, FeSO~4~(NH~4~)~2~SO~4~, is one example. K~4~\[Fe(CN)~6~\] is a complex, while 6H2O is a double salt.      **EXPERIMENT 2 Synthesis of Potassium Tris(oxalato)ferrate(III) Trihydrate (K3\[Fe(C2O4)3\].3H2O)** **Objectives** - synthesis a coordination compound, potassium tris(oxalato)ferrate(III) trihydrate under carefully controlled conditions. **Introduction** Synthesis is a useful technique in all areas of chemistry. This technique is important because it is the basis for developing new compounds, which may be useful for animals and vegetation. Most developments in the pharmaceutical industry as well as the introduction of new and less harmful pesticides are made possible because the chemicals involved are synthesized and tested in the laboratory. In addition, many chemical compounds used in everyday life are synthesized from simpler materials. This experiment involves preparing a substance by reacting to known quantities of chemicals. The expected product is potassium tris(oxalato)ferrate(III) trihydrate, K~3~\[Fe(C~2~O~4~)~3~\]·3H~2~O. The preparation of this compound involves several steps. Firstly, ferrous ammonium sulphate, Fe(NH~3~)~2~(SO~4~)~2~.6H~2~O, is dissolved in distilled water to which excess oxalic acid, H~2~C~2~O~4~, is added to make it slightly acidic and the following reaction occurs. Fe(NH~4~)~2~(SO4)~2~.6H~2~O (aq) + H~2~C~2~O~4~ (aq) → FeC~2~O~4~ (s) + H~2~SO~4~ (aq) + (NH~4~)~2~SO~4~ (aq) + 6H~2~O (l) \...\...\...\.....(Eq 1) Ferrous oxalate, FeC~2~O~4~, is a finely divided precipitate and tends to be colloidal. However, heating the solution causes it to coagulate and facilitates separating the precipitate from the solution. Next, potassium oxalate solution, K~2~C~2~O~4~ (aq), is added to the FeC~2~O~4~ precipitate which produces a slightly basic solution to facilitate the oxidation of ferrous ion to ferric ion by hydrogen peroxide, H~2~O~2~. The following reaction takes place: \[Oxidation / 2 electrons lost\] 2Fe^2+^ → 2Fe^3+^ + 2e− \[Reduction / 2 electrons gained\] H~2~O + HO^2−^ + 2e^−^ → 3OH^−^ \[Overall net reaction\] 2Fe^2+^ + H~2~O + HO^2−^ → 2Fe^3+^ + 3OH^−^\...\...\.... (Eq 2) Note that FeC~2~O~4~ is the source of Fe^2+^ ions in Eq. 2. The OH^−^ ion concentration of the solution is high enough so that some of the Fe3+ ions react with OH^−^ ions to form ferric hydroxide (brown precipitate) as follows: Fe~3+~ + 3OH^−^ → Fe(OH)~3~ (s)\...\...\....(Eq 3) With the addition of more H2C2O4, Fe(OH)3 dissolves and the soluble complex K~3~\[Fe(C~2~O~4~)~3~\].3H~2~O is formed according to: 3K~2~C~2~O~4~ (aq) + 2Fe(OH)~3~ (s) + 3H~2~C~2~O~4~ (aq) →2K~3~\[Fe(C~2~O~4~)~3~\].3H~2~O (aq) + 3H~2~O (l)..........(Eq.4)  Alcohol is added to the solution to cause the complex iron salt to precipitate since it is less soluble in alcohol than in water. The complexity of the series of reactions described in equations 1 -- 4 may be greatly simplified by following the Fe^2+^/Fe^3+^ ion throughout. It can be seen that for every mole of Fe(NH~3~)~2~(SO~4~)~2~.6H~2~O used as the starting material, one mole of K~3~\[Fe(C~2~O~4~)~3~\].3H~2~O will be obtained as the final product. Fe(NH~4~)~2~(SO~4~)~2~.6H~2~O (aq) + H~2~C~2~O~4~ (aq) → FeC~2~O~4~ (s) + \...etc (see Eq.1) FeC~2~O~4~ + K~2~C~2~O~4~ + H~2~O~2~ → Fe(OH)~3~(s) + \...etc (see Eq.2 & 3) (Fe^2+^ = Fe^3+^) Fe(OH)~3~ + H~2~O~2~ + K~2~C~2~O~4~ → K~3~\[Fe(C~2~O~4~)~3~\].3H~2~O (aq) + \...etc (Eq.4) **Safety precautions** All chemicals are toxic. The organic solvents are highly flammable and can be irritating to body tissues and respiratory tract. Acids are corrosive. Work with care and wear safety goggles. Wash hands thoroughly with soap and water before leaving the laboratory. **Apparatus** Balance, 125 mL, Erlenmeyer flask, Measuring cylinder, 50 mL (2 pieces), Beakers (100 mL, 3 pieces), Beaker, 400 mL, Glass rod, Ice bath, Bunsen burner, Tripod, stand, Thermometer, Buchner funnel, Vacuum filtration apparatus, Filter paper, **Chemicals** Ferrous ammonium sulphate, Fe(NH~4~)~2~(SO~4~)~2~.6H~2~O, 10% oxalic acid, H~2~C~2~O~4~ (aq), 1 M potassium oxalate monohydrate, K~2~C~2~O~4~·H~2~O (aq), 10% hydrogen peroxide, H~2~O~2~ Ethanol 1:1 ethanol/water solution, Deionised water, **Procedure** Students are supposed to write the procedure on their own. Make sure to use the correct tenses (past tense, passive voice) Note: The complex is photosensitive and should not be exposed to light. Store in a sample bottle/vial wrapped in foil. Keep the crystals for the next experiment Useful Molar Masses K~2~C~2~O~4~ = 166.22 g/mol Fe(NH~4~)~2~(SO~4~)~2~.6H~2~O = 392.13 g/mol K~3~\[Fe(C~2~O~4~)~3~\].3H~2~O = 491.24 g/mol **Example lab report of Synthesis of potassium tris (oxalato) ferrate (III) trihydrate** **Purpose** - to synthesis potassium tris (oxalato) ferrate (III) trihydrate , K3 \[Fe (C2O4)3\]. 3H2O. **Introduction** Ferrous ammonium sulfate, Fe(NH4)2(SO4)2. 6H2O is dissolved in a slightly acid solution, excess oxalic acid, H2C2O4, is added and the following reaction takes place: Fe(NH4)2(SO4)2. 6H2O + H2C3O4 FeC2O4(s) + H2SO4 + (NH4)2SO4 + 6H2O FeC2O4 is finely divided precipitate and tends to be colloidal. However, heating the solution causes it to coagulate and facilitates separating the precipitate from the solution. Potassium oxalate is added to the FeC2O4 precipitate, which produces a slightly basic solution for the oxidation of the ferrous ion to the ferric ion, by hydroxide, H2O2. The following reaction takes place: H2O + HO2- +2Fe2+ 2Fe3+ + 3OH- The OH- ion concentration of the solution is high enough so that some of the Fe3+ reacts with OH- to form ferric hydroxide(brown precipitate) as follows: Fe3+ + 3OH- Fe(OH)3 With the addition of more H2C2O4, the Fe(OH)3 dissolves and the soluble complex K3\[fe(c2o4)3\]. h20 is formed according to : 3k2C2O4 + 2Fe(OH)3 + 3H2C2O4 2K3\[Fe(c2o4)3\]. 3H20 + 3h2o Ethanol is added to the solution to cause the complex iron salt to precipitate. **Data Analysis and Discussion** In this experiment, I have studied how to synthesize coordination compounds. Coordination compounds are formed when a neutral metal atom: Fe acting as a Lewis acid, reacts with some neutral molecules, acting as Lewis bases; or when a metallic cation, acting as a Lewis acid, reacts with any of a variety of organic or inorganic molecules, cations, or anions, acting as Lewis bases. These Lewis bases: C2O4 and H2O are called ligands. (Lewis acids are electron pair acceptors and Lewis bases are electron pair donors. Ferrous ammonium solution is added with oxalic acid dihydrate solution will form yellow solution with yellow precipitate. Fe(NH4)2(SO4)2. 6H2O + H2C3O4 FeC2O4(s) + H2SO4 + (NH4)2SO4 + 6H2O Then it is heated boiling and the supernatant is decanted. As it is added with solid potassium oxalate, it is allowed to heat at 40 0 C and drop wise added with H2O2 and the solution turns to brown with precipitate for the oxidation of the ferrous ion to the ferric ion. H2O + HO2- +2Fe2+ 2Fe3+ + 3OH- Fe3+ + 3OH- Fe (OH) 3 Next, more oxalic acid dihydrate is added until the solution turns to colourless. 3k2C2O4 + 2Fe (OH) 3 + 3H2C2O4 2K3 \[Fe (c2o4)3\]. 3H20 + 3h2O The colourless solution is boiled then it turns to pale green solution. The solution is filtered then leaves for crystallization. After that, the green crystal is filtered and washed with 1: 1 ethanol/ water and cooled acetone. The mass of bright (luminescent) green crystals is obtained which is 3. 2822 g. So, the percent yield of K3\[Fe(C2O4)3\]. H2O that I have obtained is 47. 72 %. The precautions that we must take are while heat the solution of ferrous ammonium sulfate and solution of oxalic acid dihydrate as it will bump. Next, beware of temperature (at least 40 0 C) of solution when add H2O2 into the solution. The next experiment is determination of the percentage of ligands in coordination compounds. **Conclusion** I have studied how to synthesis coordination compound which is potassium tris (oxalato) ferrate (III) trihydrate , K3 \[Fe (C2O4)3\]. H2O. The mass of bright (luminescent) green crystals is obtained which is 3. 2822 g. So, the percent yield of K3\[Fe(C2O4)3\]. 3H2O that I have obtained is 47. 72 %. **Reference** 1\. Hadariah Bahron, Kamariah Muda, S. Rohaiza S. Omar, Karimah Kassim (2011). 2\. Inorganic Chemistry. Experiments for Undergraduates, UPENA UiTM 2008. 3\. http://chem. science. oregonstate.edu/courses/ch221 **Experiment 3: Synthesis of Hexaammine Cobalt(III) Chloride** **Introduction** The purpose of this experiment was to synthesize a 6-coordinate cobalt(III) compound from CoCl~2~ 6H~2~O. This is made difficult by the fact that Co^2+^ ion is more stable than Co3+ for simple salts. There are only a few salts of cobalt(III), such as CoF~3~ , that are known. However, cobalt(III) can be made stable when in octahedrally coordinated compounds.. Originally, it was questioned whether the chlorine atoms in hexaammine cobalt(III) chloride were coordinated or ionic. It has since been determined that the chlorine atoms are indeed ionic. There are many ways to ascertain this, the most successful of which is to complex the cobalt iodometrically and titrate the liberated iodine with sodium thiosulfate solution. A difficulty in this experiment is the oxidation of cobalt(II) to cobalt(III). This could be accomplished through the addition of hydrogen peroxide, but this method is not suitable for this experiment. A more suitable method is the air oxidation of cobalt with carbon as a catalyst. An additional benefit of carbon as a catalyst is its ability to shift the equilibrium in favor of the desired product. **Theory** As was indicated in the introduction, cobalt(II) is oxidized to cobalt(III) via the following reaction: CoCl~2~ 6H~2~O + 5NH~3~ + NH~4~Cl ‡ \[Co(NH~3~)~6~\]Cl~3~ + 6H~2~O + H^+^ (1) in the presence of air and carbon. The UV-visible spectrum may then be used to judge the quality of the product. **Experimental** A vacuum filtration apparatus was assembled and connected. 6.0065 g of CoCl2 6H2O was massed and placed into a 150 mL beaker. Then, 4.0164 g of NH4Cl was massed and added to the 150 mL beaker. To this beaker was added 5 mL nanopure water, and the solution was stirred until most of the crystals dissolved. This solution was a dark redpurple color. Next, 13.0 mL of concentrated ammonia was added to the beaker. This caused the color to change from dark purple to a brighter red. After time, it appeared that there were two layers: one dark purple and another the brighter red color. The contents of the 150 mL beaker were then transferred to a 125 mL sidearm flask and 0.13 g activated carbon was added. A water aspirator was hooked to the 125 mL sidearm and air was passed through the solution in order to oxidize the cobalt. The air was allowed to pass through the solution for 2 1/2 hours, but the cobalt did not appear to be oxidized at that point. Thus, the contents of the sidearm flask were stored for 7 days. After storage, the solution appeared purple in color. This solution was filtered on a Buchner funnel. This filtrate was then added to 40 mL water to which 0.5 mL concentrated HCl had been added. To insure that the solution was acidic, 4 more drops of con. HCl were added. The solution was then heated to 60°C and filtered to remove the decolorizing carbon. To the warm, filtered solution was added 10 mL con. HCl. This was allowed to cool to room temperature. After the HCl was added, orange crystals began to form in the bottom of the beaker to which the solution was transferred. Once this solution had cooled to room temperature, the beaker was placed into an ice bath and allowed to cool to 0°C. Upon cooling, more orange crystals formed. These crystals were filtered on a Buchner funnel and washed with 60% and 95% solutions of ethanol. The orange crystals were then dried until completely free of water. The contents were saved and the UV-visible spectrum was taken of a dilute solution of the product. **Results and Discussion** The UV-visible spectrum of hexamine cobalt(III) chloride is shown below: The absorption peak from approximately 400 nm to 600 nm means that the compound absorbs light in the violet-blue-green range. The sum of the non-absorbed, or reflected, wavelengths gives the product is orange color. The two peaks on the spectrum are d-d transitions for the d6 cobalt complex. Judging by the Tanabe-Sugano diagram for d6 complexes, the ground state 5D would be split into a 5T~2g~ and a 5E~g~. The stronger of the two peaks is most likely the transition between these two states. The weaker peak may be a spin-forbidden transition, which cannot be accurately predicted. The hexaammine complex is most likely a weak field, with four unpaired electrons. Since the ligand NH3 is of intermediate field strength, this assumption is based on a calculation of Dq/B of 1.8 for the complex, which is in the weak field region of the Tanabe-Sugano diagram. A table of products and reactants is shown below:  79.1% yield of hexaammine cobalt(III) chloride was obtained. This was very reasonable considering the literature stated to expect approximately an 85% yield. If this experiment was to be performed again, a different method of oxidation may be chosen, due to the large amount of time required for air oxidation to be effective. **Reference** G. Pass, H. Sutcliffe (1974), Practical Inorganic Chemistry, Chapman and Hall, Ltd: New York. 80-81. J. Tanaka, S.L. Suib (1999), Experimental Methods in Inorganic Chemistry, Prentice Hall: Upper Saddle River, NJ. 272-74 **Experiment 4: Preparation of sodium thiosulphate** **Objective** To prepare sodium thiosulphate (Na~2~S~2~O~3~) **Theory** Sodium Thiosulphate is an inorganic salt with the chemical formula Na2S2O3. It is a white colorless crystal or even powder. When degraded into sulfide and sulfate in the air, the material is known to be alkaline. Sodium thiosulphate is soluble in water, producing thiosulfate ions, excellent reducing agents. When it is decomposed by heat, it generates very poisonous vapors of sulfur oxides. It is frequently discovered in toxic waste from the dye industry. It is an ionic compound composed of two sodium cations and one thiosulfate anion, with one sulfur atom connected to three oxygen atoms and one more sulfur atom. Double bonds connect the oxygen atoms to the sulfur atom, and a single bond connects the other sulfur atom. It has tetrathionate stability, which means it can form stable tetrathionate complexes in solution. Because of its capacity to effectively neutralize chlorine, sodium thiosulphate is a common component in water treatment dechlorination operations. These properties, combined with their low toxicity, contribute to sodium thiosulphate's wide range of applications in sectors such as photography, medicine, and analytical chemistry. Historically, sodium thiosulfate has been utilized as a nephroprotectant during the administration of cisplatin and as an antidote for cyanide poisoning. It is believed to have cation-chelating as well as antioxidant effects. Preparation of sodium sulphate can be done through three steps: **i). From sodium bisulfite** Sodium thiosulphate and water are formed when sodium hydroxide (NaOH) interacts with sulfur (S) and sodium bisulfite (NaHSO3). **2NaOH + S + 2NaHSO3 → Na2S2O3 + 2H2O** **ii. From sodium hyposulphate** It can be prepared by the decomposition of sodium hyposulphate Na2S2O4. Na2S2O4 → Na2S2O3 + Na2SO3 **iii. From SO2** It can be prepared by passing SO2 gas into waste liquor obtained in the manufacture of Na2S. Waste liquor contains Na2S, Na2CO3 and Na2SO3. **2 Na2S + Na2CO3 + 4 SO2 → 3 Na2S2O3 + CO2** **2 Na2S + Na2SO3 + 3 SO2 → 3 Na2S2O3** Sodium thiosulphate is a chemical with a unique structure made up of sodium cations Na+ and thiosulphate anions S2O32- The sodium ion Na+, which has a positive charge, is at the heart of sodium thiosulfate. Sodium is an alkali metal that is well-known for its reactivity and tendency to create cations by losing its outermost electron. Thiosulphate anions, which contain two sulfur atoms (S) and three oxygen atoms (O), surround the sodium ion. A double bond connects the sulfur atoms, forming a core S-S link. A single connection connects one of the sulfur atoms to an oxygen atom, forming an S-O bond. The other sulfur atom forms double bonds with two oxygen atoms, resulting in S=O bonds. The thiosulfate anion has a distinct structure that resembles a bent or "U" shape, with the sulfur atoms at the bends and the oxygen atoms projecting outward. Ionic bonding attracts the sodium cations to the negatively charged thiosulfate anions, which keeps the molecule stable. **Physical Properties of Sodium thiosulphate** Sodium thiosulphate is an inorganic chemical that appears as a white transparent, colorless crystal. It is a water-soluble chemical that is also soluble in turpentine oil, but not in alcohol. The melting point of the material is between 48 and 52 degrees Celsius. This chemical compound is extremely stable in nature and is claimed to be incompatible with certain strong oxidizing agents and acids. Thiosulfate anion easily reacts with dilute acids, releasing sulfur, sulfur dioxide, and water. The compound has a density of approximately 1.667 g/mL. **Chemical properties of Sodium thiosulphate** **Reducing property** In some chemical processes, sodium thiosulphate acts as a reducing agent. When coupled with an oxidizing agent such as iodine, sodium thiosulfate converts iodine to iodide ions. The following is the reaction: **2 Na2S2O3 + I2→2 NaI+Na2S4O6** **Reaction with salt** A dilute solution of sodium thiosulphate reacts with AgNO3, and a white precipitate of silver thiosulphate (Ag2S2O3) is formed which changes from yellow, orange, brown, and finally to a black precipitate of Ag2S. **2 AgNO3 + Na2S2O3 → Ag2S2O3 + 2 Na2NO3** **Ag2S2O3 + H2O → Ag2S+ H2SO4** **Complex formation reaction** Sodium thiosulfate can form complexes with a variety of metal ions. This feature is useful in analytical chemistry, where sodium thiosulphate is widely employed as a titrant to detect the concentration of particular metal ions in solution, such as silver or copper. **6CuSO~4~ + 11Na~2~S~2~O~3~ → 6Na~2~SO~4~ + 3Na~2~S~4~O~6~ + 2Na\[Cu~6~(S~2~O~3~)~5~\]** **Reaction with acid** When sodium thiosulphate Reacts with acids (HCl, H2S04, HNO3), sulfur dioxide gas is released and Colloidal sulphur (White or yellow turbidity) is formed. The reaction can be expressed as follows: **Na~2~S~2~O~3~ + 2HCl → 2 NaCl + H~2~O + SO~2~ + S** **Na~2~S~2~O~3~ + H~2~SO~4~ → Na~2~SO~4~ + H~2~O + SO~2~ + S** **Na~2~S~2~O~3~ + HNO~3~ → NaNO~3~ + H~2~O + SO~2~ + S** **Dechlorination reaction** Sodium thiosulphate can be used to dechlorinate water. It interacts with chlorine to produce sodium sulphate and harmless chloride ions, making it very useful in water treatment operations to remove excess chlorine from water supplies. **Na~2~S~2~O~3~ + Cl~2~ + H~2~O → Na~2~SO~4~ + HCl + S** **Decomposition reaction** Under normal conditions, thiosulphate is a stable chemical. However, when heated, it decomposes to form sodium polysulfide and sodium sulfate. **Uses** - It is essential in the development and fixing of photographic films. - In the medical field, sodium thiosulphate is extremely important. It is used as an antidote for cyanide poisoning, successfully neutralizing cyanide's poisonous effects by interacting with it to generate a less hazardous molecule. Because of its feature, sodium thiosulfate is a vital component in emergency medical kits and a lifeline in dire conditions. - It is used for the determination of the strength of iodine or the determination of the strength of oxidizing agents like potassium permanganate and potassium dichromate. - Another application for sodium thiosulphate is in the treatment of water. It works as a dechlorinating agent, eliminating chlorine from water. This is especially crucial in aquariums because chlorine can be toxic to aquatic life. - When dissolved in a large amount of warm water, the chemical can be used as a cleaning agent. - It is used in the textile industry as a bleaching chemical to help remove undesired dyes and stains from materials. Furthermore, in certain chemical reactions, this molecule works as a chlorine scavenger, preventing the degradation of other compounds due to the presence of chlorine. In industrial environments, where chemical stability is critical, such protective characteristics are highly valued. **Health hazards and precaution** Direct contact with sodium thiosulphate may cause skin and eye irritation. To avoid any potential contact, it is recommended to wear protective gloves and safety eyewear when handling the compound. Inhaling sodium thiosulphate dust or fumes might cause respiratory discomfort. When handling significant amounts of material or generating dust, it is best to operate in a well-ventilated location or wear suitable respiratory protection. Some people may experience sensitization or allergic reactions to sodium thiosulphate such as skin rashes, itching, or respiratory symptoms. **References** https://www.sciencedirect.com/topics/biochemistry-genetics-and-molecular-biology/sodium-thiosulfate. https://byjus.com/jee/sodium-thiosulphate/ https://www.vedantu.com/jee-main/chemistry-sodium-thiosulphate https://www.ncbi.nlm.nih.gov/pmc/articles/PMC4750847/