Hi, I am not well at all, and need to submit this lab as soon as possible and am busy with alot of other stuff too. Please help with it. Will be really thankful and would rate good.
Introduction
Decomposition of Sodium Chlorate Mass, Moles, and the Chemical Equation
Sodium chlorate is used as a source of oxygen in emergency oxygen generators. So- called oxygen canisters or oxygen candles are found on airplanes, submarines, even the space station â anywhere oxygen might be in short supply in case of an emergency. Sodium chlorate decomposes upon heating or in the presence of metals to give oxygen gas. What is the chemical equation for the decomposition of sodium chlorate?
Concepts
ï· Moles ï· Balanced chemical equation
ï· Molecular formula ï· Stoichiometry
Background
Sodium chlorate, NaClO3, is a colorless, odorless, white solid that melts at 248 °C. When heated above 300 °C, it begins to lose oxygen. The ultimate products of the thermal decomposition of sodium chlorate are oxygen gas and a white solid. Based on the molecular formula of sodium chlorate, three possible reactions will account for the loss of oxygen gas upon heating (Equations 1-3). Note that equations 1-3 are not balanced.
Equation 1 2NaClO3(s) â 2NaClO2(s) + O2(g)
Equation 2 NaClO3(s) â NaClO(s) + O2(g)
Equation 3 2NaClO3(s) â 2NaCl(s) + 3O2(g)
What is the actual chemical equation for the thermal decomposition of sodium chlorate? All of the possible sodium-containing products n Equations 1-2 are real compounds: sodium chlorite, NaClO2; sodium hypochlorite, NaClO; and sodium chloride, NaCl. All are white solids at room temperature. It is possible to determine the chemical equation for the decomposition of sodium chlorate by applying the principles of stoichiometry to the masses of the reactants and products.
Safety Precautions:
Sodium chlorate is a strong oxidizing agent and a dangerous fire risk; it is slightly toxic by ingestion. Contact with metal powders or combustible organic compounds may cause fires. Keep away from contact with organic materials, including rubber stoppers, rubber tubing, etc. Avoid contact with eyes and skin. Do NOT dispose of exess sodium chlorate in the trash.
When heating sodium chlorate, use only a small amount of solid (0.25-0.40 g) in a Pyrex test tube. Inspect the test tube for chips and cracks before use and handle the test tube using a test tube clamp. Distribute the solid evenly along the bottom of the test tube and heat the test tube clamp. If any smoke is produced, remove the test tube from the flame until the reaction subsides. Do not inhale the smoke. Allow the test tube to cool completely on a wire gauze with a ceramic, heat-resistant center. Wear chemical splash goggles and chemical-resistant gloves and apron.
Procedure:
Measure all masses to the nearest 0.01 g and record all data in the data table.
Obtain two medium test tubes. Label them 1 and 2.
Measure and record the mass of each test tube.
Using a spatula, add approximately 0.2-0.4 g of sodium chlorate to the bottom of each test tube.
Measure and record the combined mass of each test tube and sodium chlorate.
Place test tube #1 in a test tube clamp. Holding the test tube in an almost horizontal position, gently tap the test tube to distribute the solid sodium chlorate evenly along the bottom one-third of the test tube.
Light a Bunsen burner.
Holding test tube #1 with the test tube clamp, slowly move the test tube back and forth through the burner flame to gently heat the solid. The solid should begin to melt and bubble slightly. Caution: Do not aim the opening of the test tube at anymore.
Quickly remove the test tube from the flame if a white smoke is given off from the open end of the test tube. The smoke represents a loss of starting material due to evaporation of molten sodium chlorate and will reduce the yield of product. A small amount of smoke may be inevitable and should not interfere with the accuracy of the results.
Continue to heat the test tube in the burner flame until no more bubbling is observed and the material in the test tube has solidified.
Allow the test tube to cool completely on a wire gauze with a ceramic center.
While test tube #1 is cooling, repeat steps 5-11 sing test tube #2.
When test tube #1 has cooled to room temperature, measure the combined mass of the test tube and its solid contents.
Reheat test tube #1 gently for an additional three minutes.
Cool the test tube on the wire gauze until it is room temperature. Measure the mass a second time. If the second mass differs from the first mass by more than 0.02 g, repeat the heating cycle a third time.
Repeat steps 13-15 using test tube #2.
Data Table:
Trial 1
Weight of beaker --- 68.1887g
Weight of test tube -- 19.3538g
Weight of NACLO3 -- 0.3341g
Everything together -- 87.9155g
1st time burning:-- 87.8851g
2nd time burning:-- 87.7630g
3rd time burning:-- 87.7626g
Trial 2
Weight of beaker --- 68.1886g
Weight of test tube -- 22.8518g
Weight of NACLO3 -- 0.3154g
Everything together -- 91.3508g
1st time burning:-- 91.2078g
2nd time burning:-- 91.2103g
Post-Lab Calculations and Analysis:
1) For each trial, find the initial mass of sodium chlorate and calculate the number of moles of sodium chlorate.
2) Based on the molar masses of the three possible solid products in Equations 1-3, and the number of moles of reactant in each case, calculate the expected masses of the three possible products for Trials 1 and 2.
3) Compare the actual mass of product obtained in Trials 1 and 2 with the calculated or expected masses. What is the balanced chemical equation for the decomposition of sodium chlorate? Explain.
4) Use the following equation to determine the percent error in the mass of solid obtained.
â (actual mass â expected mass) / expected mass
5) Consider the following potential sources of error in this experiment. Explain whether they would have caused the actual mass of solid product to be lower or higher than the expected value.
a)The sodium chlorate did not decompose completely.
b)The sodium chlorate was heated too fast, allowing considerable white smoke to escape from the test tube.
c)The sodium chlorate absorbed some moisture from the air before the experiment began and was not completely dry.
