🏷️ LIMITED TIME OFFER: GET 20% OFF GRADE+ YEARLY SUBSCRIPTION →
Log in
HomeStudy Guides420,000US230,000

01:160:311 Study Guide - Quiz Guide: Eutectic System, Boiling Chip, Seed Crystal

451 views19 pages
School
Rutgers University
Department
Chemistry
Course
01:160:311
Professor
All

For unlimited access to Study Guides, a Grade+ subscription is required.

Document Summary

This will hasten absorption of chemical through the skin. The greatest harm comes from breathing the hot vapors that rise past the mouth. Aim it at the base of the fire & squeeze the handle and don"t hold on to the horn b/c it will become cold. A very small fire in a beaker or a flask could possibly be put out by simply covering the vessel with a flat object. A wire gauze or watch glass could put out a small fire b/c it will prevent oxygen from fueling the fire. Why is it dangerous to wear contact lenses in an organic chemistry laboratory? (be specific) The lenses could be fused to one"s eyes. Too much will cause differential in temperature throughout the sample. Coarse particles will cause air pockets which will cause an inaccurate temperature. Define eutectic point in terms of the two components of a mixture and their liquid solutions.

Get access

Grade+20% off
$8 USD/m$10 USD/m
Billed $96 USD annually
Grade+
Homework Help
Study Guides
Textbook Solutions
Class Notes
Textbook Notes
Booster Class
40 Verified Answers

Related Documents

01:160:311
Final Exam
Study Guide
01:160:311 Study Guide - Final Guide: Alpha And Beta Carbon, Le Chatelier'S Principle, Epoxide
yellowox812
01:160:311
Final Exam
Study Guide
01:160:311 Study Guide - Eutectic System, Boiling Point, Paper Towel
cerisehound184
01:160:311 Lecture Notes - Lecture 10: Borate, Lithium Aluminium Hydride, Sodium Borohydride
cyanserval624

Related Questions

Ethyl ether (CAS# 60-29-7; Formula C2H5OC2H5) is a very useful chemical that is found in many laboratories. It is a Class IA Flammable Liquid, with an OSHA PEL of 400 ppm, a flammable range of 1.9% to 36%, and an auto-ignition temperature of 320 F (160 C). Because it is a good solvent, evaporates easily, and is only slightly soluble in cold water, you decide to use it as an extraction solvent in a separatory funnel. After the extraction, you plan to rapidly evaporate the ether inside a fume hood using an electric hot-plate that is set at 318 degrees F. Unfortunately, while preparing your extraction process you dropped the bottle of ether and it broke on the floor. NO ONE is injured, but you now have a large mess to clean up. Your lab is located in a small town and you are the only member of the local volunteer Haz-Mat response team. This means you have to take care of the mess yourself - you can't just leave the lab and call 911 because you ARE 911. Before you begin the cleanup, you determine that in the area of the spill the concentration of ether vapors is 2000 ppm and the oxygen concentration is 19.3%. The most appropriate PPE to wear during the cleanup is:

Level D (gloves, lab coat, eye protection, long pants, and closed shoes). No respiratory protection is necessary because the spill will be cleaned up long-before your 8-hour Time Weighted Average exposure reaches 400 ppm, and OSHA does not have a STEL for ether. Fire isn't an immediate concern because ether vapors are heavier than air and the spill is on the floor while the hot plate is several feet above in a fume hood.

Level C. Use an air-purifying respirator with "organic vapor" cartridges installed that are capable of reducing the ether concentration that you breathe to less than 100 ppm. Also wear solvent-resistant gloves and a chemical suit. Before you start to clean up the spill, you will turn off the hot plate.

Level B. Use a Self-Contained Breathing Apparatus (SCBA) because of the relatively high vapor concentration and the relatively low oxygen concentration. Although you will wear appropriate gloves and a chemical suit, there is no need for an encapsulated (level A) suit because your skin does not need to be completely protected from the vapors like your lungs need to be protected.

Level A. Top-of-the-line, first-class protection. SCBA with fully-encapsulated, chemical-resistant suit and gloves. As soon as you dropped the bottle, you shut off the hot plate and left the area to go put on your enhanced PPE (because you are smart, you were already wearing level "D" with all-cotton clothes, leather shoes, and a fire-resistant lab coat). You will use tools while cleaning the broken glass so that you do not accidentally cut a hole in your very expensive Level-A suit.None of the above. In the box below, briefly explain why options (A) through (D) are not entirely correct and what you believe is a better way to manage this spill.

Explain your answer:

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 – 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

 Moles  Balanced chemical equation

 Molecular formula  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) → 2NaClO2(s) + O2(g)

Equation 2 NaClO3(s) → NaClO(s) + O2(g)

Equation 3 2NaClO3(s) → 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.

‌ (actual mass – 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.