Synthesis and Decomposition of Zinc Iodide
Partners: Sonya Pasia and Kristen Kobayashi
CH 085-01
20 September 2011
Zinc Iodide (ZnI2) was an interesting binary compound to experiment with. In this experiment, weakly acidified water (25mL distilled water with 18 drops 5M acetic acid solution) was used as an aid to bring molecules of the zinc and iodide atoms together, by dissolving iodine molecules, so that bonding would transpire to produce a reaction. Deprived of water, the Zn and I2 molecules would not be capable of moving close enough to each other, and a reaction would not occur. Deprived of acid, the reaction of Zn + I2 would have resulted in 2HI(aq) rather than ZnI2 (s), and it wouldn’t have appeared to follow the Law of
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Half of the solution was poured into an evaporating dish. Two copper wires were attached to a 9-volt battery on the electrodes. Wires were carefully bent and placed on opposite sides of the solution in the dish, where observations could be made regarding what has seen.
Different physical properties were noted during addition of substances throughout the course of this experiment. From the beginning, it was noted that 2.01 grams of Zn metal was silver and had shiny flakes throughout. The 1.985 grams of I2 was dark gray in color and also appeared to be shiny. Once 5mL of acidified water was added, the solution began to change colors, becoming a dark red orange, which was discovered to be closer to Zn(I3) 2. This reaction was endothermic in nature due to heat that was felt in the duration of the reaction. The heat melted some of the I2 to create the I3 in this solution once the color disappeared. After this point it was found that only 0.541 grams of Zinc was consumed in the reaction, making Zn the excess reagent and I2 the limiting reagent. After the ZnI2 was dried, a light yellow powder was noted, and recorded to be the color of ZnI2 (s). When performing the decomposition, a dark
By adding the 0.541 g Zn to the 1.985 g of I2 from the beginning of the experiment, it was predicted that there should be 2.526 g of ZnI2 in
10. As described in problem 7, a procedure was developed to determine the percent zinc in post 1982 pennies. In that procedure 50 ml of an HCl was used to react (dissolve) all of the zinc in the penny. To ensure complete reaction, the solution contains twice as many moles of HCl that is actually needed. To determine the percent zinc in the penny, the excess (unreacted) HCl was titrated with NaOH. Determine the concentration of NaOH needed if you want to use approximately 25 mL of NaOH to titrate the excess HCl.
When the zinc was dropped in the hydrochloric acid, the substance began bubbling vigorously, forming a precipitate. Eventually, the zinc dissolved completely. After the lit wooden splint broke the surface of the test tube, there was a loud popping noise. The gas that was released was hydrogen from the acid and the popping noise was a result of the Hydrogen being burned up by the fire creating a small explosion. Two chemical changes occurred in this test: one with the formation of a precipitate (a textbook sign of a chemical change), and the other when the explosion
No vigorous reaction occurred; rather, the zinc sample disintegrated slowly and turned red in color.
In this experiment, the scientists are testing which amount of zinc mossy would make a few normal cleaned pennies change their colors into silver when they are placed in the boiling solution of zinc sulfate and water and then turn their colors into the brightest yellow (gold) when they are exposed directly to heat,. The independent variables in this experiment are the amount of zinc and zinc sulfate measured in gram and the dependent variables is the shades of yellow (compared in decimal R,G, B code) of the penny after all steps of the experiment. The constant of this experiment will be the temperature of the tap water used to boil the solution, which is 100°C and the temperature of the hot plate used to heat up the silver penny, which is about 300°C. And the controlled of this experiment is the pennies with the brownish color.
There was an assortment of different changes indicating that chemical changes were taking place such as change in color or chemicals bubbling when combined with another chemical.
10. As described in problem 7, a procedure was developed to determine the percent zinc in post 1982 pennies. In that procedure 50 ml of an HCl was used to react (dissolve) all of the zinc in the penny. To ensure complete reaction, the solution contains twice as many moles of HCl that is actually needed. To determine the percent zinc in the penny, the excess (unreacted) HCl was titrated with NaOH. Determine the concentration of NaOH needed if you want to use approximately 25 mL of NaOH to titrate the excess HCl.
Purpose: The purpose of this experiment is to investigate the physical and chemical properties of pure chemical substances by subjecting them to various environmental extremes.
John. A. Avers conducted a purification experiment of Zirconium via ion exchange to determine the percentage of recovery of the element. John A. Avers used various resins to determine which would obtain the highest purification of zirconium compared to other ion exchange resins and the extraction of the element will be conducted by using Zirconium from ores. Some of the resins used consisted of rare earths, beryllium, titanium and iron. John A. Ayers took zirconium ore’s and broke them down by using concentrated sulfuric acid. 8 Once the zirconium ore was broken down, the soluble zirconium was then converted to its nitrate form, Zr(NO3)4. The zirconium nitrate is then passed through a column containing a resin. The hydrogen form of Amberlite I.R.-100 was used as the
The main purpose of this experiment was to show that single displacement reactions between metals according to their reactivity, with more reactive elements having the power to displace less reactive elements and take their place in a chemical compound (Beran, 2014). This was supported by the results of the experiment, where solid metals were combined with aqueous solutions that contained another element, and reactions only took place when the solid metal was more reactive than the other element in the compound. Only three attempted trials resulted in a failure to produce a reaction, namely the combinations of copper with hydrochloric acid, and copper with nickel sulfate. The outcomes of these trials are justifiably reasonable because copper is ranked lower in the
Using a circuit, if the substance was electrically conductive the lights would light up. The independent variables were the different samples, and the dependent variables were the conductivity testers. We used distilled water to act as a polar solvent for potential electrolytes. We hypothesized that the distilled water would not create an electric current and light up the conductivity testers, or produce bubbles. Distilled water is not electrically conductive based on Table 15.3.
Thomas and Fray (1981) studied leaching of oxide zinc materials with chlorine and chlorine hydrate. They found that the rate of leaching of the Adrar Turkish ore could be described using a shrinking core diffusion model, and that the rate of leaching was controlled by surface reaction. In all cases studied, lead was also leached with zinc. However, iron oxides remained virtually undissolved. Frenay (1985) investigated leaching of oxidized zinc ores in various solution media and obtained the best leaching results with sulfuric acid and caustic soda. Mineralogical studies showed that smithsonite can be completely leached but hemimorphite is relatively refractory to leaching. Abdel-Aal (2000) studied the kinetics of sulfuric acid leaching of low-grade zinc silicate ore and reported that about 94% of zinc can be leached after 180 min of reaction with 10% sulfuric acid at 70°C. Espiari et al., 2006 studied hydrometallurgical treatment of tailings with high zinc content using sulfuric acid and reported that the
In this experiment, a sample of K2S2O8 was prepared by the electrolysis of an aqueous solution of H2SO4 and K2SO4. The peroxodisulfate anion, S2O82-, was also observed for its ability to serve as a counterion for precipitation by preparing a copper (II) complex by reacting hydrated copper (II) sulfate with ammonium peroxodisulfate in the presence of pyridine. This same ability, coupled with its strong oxidizing ability allowed for stabilization of the unusual oxidation state of 2+ for silver which was observed by preparing an analogous silver (II) complex by reacting silver (I) nitrate with ammonium peroxodisulfate in the presence of pyridine. IR spectra for the three products were
The purpose of this lab was to evaluate our skills of decanting a supernatant liquid without losing the solid and successful completion of a series of reactions. This was done through five chemical reactions involving copper. In this lab, elemental copper was put through five different chemical reactions in order to convert it into different compounds. By the end of the fifth reaction, the copper was back to its elemental state.
12. The crocodile clips are attached to the copper electrodes of the experimental apparatus and the power supply is turned on. Simultaneously, the stopclock is started. The thermometer is checked every 30s. 13. After 300s the stopclock is stopped and the power supply is turned off. The negative cathode is carefully removed and is dried using a hair dryer. 14. When dry the negative cathode is placed on the electronic milligram balance and its final mass is recorded. 15. The positive anode and negative anode of the experimental apparatus are disposed and the electrolyte is poured out to ensure that the anode slime (impurities) does not contaminate the solution. 16. The electrodes of the experimental apparatus are replaced with new copper strips. 17. Steps 7 to 16 are repeated. However, this time, the rheostat is adjusted using the calibration apparatus until the multimeter shows approximate readings of 0.40 A, 0.60 A, 0.80 A and 1.00 A respectively. 18. Time permitting, the entire experiment is repeated. Safety Copper sulphate may cause irritation and burns if it comes into contact with the eyes. As standard lab procedure, safety goggles and lab coats must be worn at all times. Control of Variables Volume of Electrolyte Used
At the end of the experiment when the lid was removed, it was found out that the blue colour of the copper (II) sulphate solution has faded away. It was turned to pale grey and there were some precipitates present. It was the zinc powder that was in excess to ensure that the copper (II) sulphate solution could react fully with the zinc powder.