Lab #5 – Formation of a Salt Introduction: It is amazing that food is seasoned with an ionic compound that is composed of two deadly elements – sodium and chlorine. The gain or loss of electrons can make a big difference in properties. Reacting sodium bicarbonate (baking soda) with hydrochloric acid (HCl), the acid found in your stomach, produces carbon dioxide, water, and salt according to the following equation: NaHCO3 + HCl NaCl + CO2 + H2O If we evaporate the water, then all that should remain is the salt, NaCl. Materials: 0.5M HCl Distilled water (dH2O) NaHCO3 Bunsen burner 100-mL beaker Ring stand 10-mL graduated cylinder Ring clamp Plastic pipet Wire gauze Buret with buret clamp Balance Methyl red indicator Procedure: 1. Mass a clean, dry 100-mL beaker 2. Mass 0.50g sodium bicarbonate (NaHCO3) on a piece of weighing paper, and add it to the beaker 3. Add about 15mL dH2O (about 15mL, but record the exact amount) to the beaker and swirl the solution gently to dissolve the sodium bicarbonate. Add more water if necessary to dissolve the powder completely. 4. With the pipet already in the Erlenmeyer flask, add 2-3 drops methyl red indicator solution to the beaker. The solution should be a pale yellow in color. Place a piece of white paper under the beaker to view the color of the solution better. 5. Assemble the buret apparatus as shown by your instructor. 6. Add approximately 25mL 0.5M HCl to the buret. Remember – to measure on a buret, you must subtract the number you read from 50mL to get the true volume (that is, if you read 35mL; 50mL – 35mL = 15mL solution added). 7. While gently swirling the beaker, put your lab goggles back on and add the hydrochloric acid drop by drop until the color of the solution changes to a definite red. 7a. As you add each drop, what color does the solution change immediately next to where the drop is added? 7b. As the critical amount of HCl approaches, does it take the solution longer to return to yellow? Maybe you should wait a bit longer between adding drops to make sure it will return… 8. RECORD the amount of HCl added to your solution. Again, remember how to read a buret correctly. 9. Place the beaker atop a ringstand and wire gauze and begin to heat the contents to evaporate the water. CAUTION: Do not heat the solution too much or it will spatter out of the beaker. 10. When only about 5mL of water is left in the beaker, shut off the flame, put your goggles back on, and allow the heat of the beaker to evaporate the rest. 11. Allow the beaker to cool for at least five minutes. Remember, the beaker will appear cool before it is ready to be handled. 12. Mass the cooled beaker with the white powder. You did remember to mass it initially in step 1, right? 13. Examine the contents of the beaker. Use a hand lens to see if the powder has the characteristic cubic shape of sodium chloride. Cleanup and Disposal: 1. Place unused chemicals in the solid or liquid waste containers. 2. Rinse out the contents of your cooled beakers in the sink. 3. Make sure the balance and fume hood areas are left in the same condition as how you found them. 4. Wipe down desk surfaces with a wet paper towel. Questions to Consider: As you added the hydrochloric acid, what did you observe? What gas was released during the chemical reaction? The sodium bicarbonate underwent a chemical change. What evidence do you have of this change? Describe the white powder that formed. How can we get a solid from two liquids? How can you identify the product as being different from the reactant? To ensure that the white powder was all sodium chloride and not mixed with sodium bicarbonate, would you need to add a little less or a little more hydrochloric acid to the reaction? Explain your choice. Mass is neither created nor destroyed. Yet the mass of the product (NaCl) was less than the mass of the reactant (NaHCO3). How can you explain this? Where did the mass go? What might have affected the accuracy of this experiment? Sodium bicarbonate is a common ingredient in antacid remedies. Knowing this, and using information from the equation for the reaction, explain how this chemical could relieve a stomach that contains excess acid. What gas would be formed? Studies have proven conclusively that fluoride is an effective tooth decay preventative. As a result, in the late 1960’s and 1970’s, many communities in the United States began adding trace quantities of fluoride to their drinking water supplies. However, strong opposition arose against this “tampering” with the water supply. One of the common arguments was that fluorine was known to be a deadly gas. What would be your response to this argument?