Chlorine gas can be produced in the laboratory by adding concentrated hydrochloric acid to manganese(IV) oxide in the following reaction: MnO2(s) + 4HCl(aq) ( MnCl2(aq) + 2H2O(l) + Cl2(g) a. Calculate the mass of MnO2 needed to produce 25.0 g Of Cl2 ans: 30.7 g MnO2 b. What mass of MnCl2 is produced when 0.091 g of C12 is generated? ans: 0.16 g MnCl2 1. How many moles of ammonium sulfate can be made from the reaction of 30.0 mol of NH3 with H2SO4 according to the following equation: ans.
Once the zinc was added slowly to the solution, a gas was released and the solution began to change colors. The colors that occurred were light blue to light grey, to a clear blue, back to a dark blue, then grey color. Once the stirring process ended, the solution was slightly blue and not colorless. There were solids in the bottom. The reaction that occurred with this step was displacement and metathesis in the form of gas formation.
What volume of 2.0 M hydrochloric acid is needed to completely react with the amount of calcium carbonate in Part 2a above? c. Based on Parts 2a and 2b above, how many moles of water would be produced? 3. Ammonium chloride and calcium hydroxide react according to the following balanced equation: 2 NH4Cl(aq) + Ca(OH)2(aq) ⋄ CaCl2(aq) + 2 NH3(g) + 2 H2O(l) a. What mass of ammonium chloride is needed to make 3.0 liters of a 1.5 M ammonium chloride solution?
Calculate the number of moles of calcium chloride that would be necessary to prepare 85.0 g of calcium phosphate. 3CaCl2 + 2Na3PO4 —> Ca3(PO4)2 + 6NaCl 3. Sodium chloride (0.400 Mole) is allowed to react with an excess of sulfuric acid. How many moles of hydrogen chloride could be formed? 2NaCl + H2SO4 —> Na2SO4 + 2HCl 4.
Introduction: Thyme contains a surprising amount of iron compounds. This experiment enables us to determine the amount of iron(II) present in dried thyme by means of a redox reaction. Method: 1. Weigh accurately about 1g of dried thyme and put into a 250cm3 beaker. Record the mass of thyme used.
Percent Copper and Formula Weight of a Copper Compound Introduction: In many chemistry problems, you are asked to calculate the percent composition of each element in a compound. In this experiment a variation of this basic calculation will be employed. The amount of copper in a compound will be determined by dissolving the unknown compound in water. The copper ions in solution will be converted to metallic copper by reaction with magnesium metal. To ensure that all of the copper is removed from solution, an excess of magnesium will be used.
Single Replacement Reaction Laboratory Modified from Glencoe Chemistry - Matter and Change, Glencoe McGraw-Hill, 2002 Objectives Observe a single replacement reaction Measure the masses of iron and copper Determine the mole ratios and the limiting reactant Chemicals Iron filings (Fe) – 20 mesh Copper(II) sulfate pentahydrate, (CuSO4·5H2O) Distilled water Materials Stir rod 100-mL beaker 250-mL beaker 25-mL graduated cylinder Weigh paper Balance Hot plate Beaker tongs Wire mesh insulated pad screen Distilled water wash bottles |Lab Data - Reaction of Copper(II) Sulfate and Iron | | Mass of empty 100-mL beaker |(g) | | | Mass of 100-mL beaker
LAB 5: Taking the A train Focus Question Amphoteric material can exhibit properties of a base as well as properties of an acid. In our experiment we combine Zinc and Iodine to create Zinc Hydroxide, which suppose to be Amphoteric. So in order to check if our substance is indeed amphoteric we will check its behavior when tested for Acidity and for Base. If in both case it will exhibit a positive results, we will reaffirm the assumption it is Zinc Hydroxide and amphoteric. Materials used: 2.00 Gr granule Zinc – grayish hard material, light, no smell.
Standard Molar Volume The ultimate goal of this lab was to find the standard molar volume of hydrogen gas (H2). An unknown sample of metal to 3M H2SO4 and an eudiometer filled with water. The dense acid sank towards the bottom to react with the metal sample and form Hydrogen gas. The gas raising to the top of the container caused the pressure in the eudiometer to increase, which lead the water to be displaced. | Trail 1 | Trial 2 | Code | Skinny | Skinny | Mass of Metal | .041g | .027g | Temperature of Water | 296k | 296k | Vapor of Water Temperature | 21.1 mmHg | 21.2mmHg | Barometric Pressure | 76.632cm | 76.632cm | Volume of H2 collected | 29.15mL | 29.2mL | Height of Supported H2 column | 23.95cm | 23.15cm | After the O2 gas had fully reacted, measurements of mass, temperature, vapor, and H2 collected (as shown in the table above) .
h) A way to make hard water softer is to put an sodium nitrate and create a precipitate to mellow out the reaction. Another way of making it softer is by removing the calcium ions one way of doing that is by boiling the solution to take out some of the ions. Conclusion: Overall, we determined that sodium carbonate, Na2CO3, is the anion that can be used to precipitate the most metal cations. Also, we learned that the anion sodium chloride, NaCl, could be used to remove silver ions from solutions. The stuff that I found interesting was that how many colours you can get when you mix the cations and anions