Spontaneity-2 - July 2002 Equilibrium and Spontaneity...

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July 2002Equilibrium and SpontaneityIntroductionIn lecture, you’ve discussed thermodynamics, solubility, acids and bases and more. Although Chemistry seems to be broken up into several distinct topics, the goal is to get the ‘big picture’— to see how all the topics overlap and depend on each other. Understanding the big picture is a process that starts now, and continues until you stop asking “I wonder how that works?” The cool thing is that chemistry uncovers the secrets of the reallybig picture; how the universe works. Everything that happens in the universe is an interaction of mater and energy and that’s what chemistry is all about. In the words, “Understand chemistry and you understand the world!”This investigation is geared toward your putting several Chemistry ideas together to see a small slice of the big picture. You will use your understanding of acids, bases and titration to determine the solubility product (Ksp) of a compound commonly called Borax. As you have seen in lecture, the free energy (G) of a reaction and the equilibrium constant (like Ksp) depend on temperature. By varying the temperature of Borax solutions and measuring Ksp, you will be able to calculate the thermodynamic parameters of the Borax dissolution reaction. We will look at the mineral borax, which is also commonly found in lakebeds in the dessert southwest. An entire industry revolves around the collection of borax from dry lakebeds for use in soaps, solder, and as a preservative.. The tetraborate anion is a weak base. It reacts with water (hydrolyses) to form the hydroxide ion (OH-), a strong base (and boric acid, H3BO3, an extremely weak acid that we won’t be concerned with):The important thing is that amount of tetraborate in solution is precisely related to the amount of OH-. So if we can determine the number of OH-ions in solution, we know how many tetraborate ions were there. We will use titration to count the number of OH-ions in solution. Titration of OH-simply means we will add H+ions in the form of a strong acid (HCL), until we’ve added one H+for each OH-present. The H+and OH-react to form water. The point where all the OH-is used up (converted to water) is called the endpoint of the titration. It is usually not obvious to the naked eye, so to find the endpoint, we will use an acid/base indicator that will change color just when the OH- ions are gone. If we immediately stop adding acid (H+) at the endpoint, the number of OH-ions originally present is precisely the number of H+(acid) ions that we added. We can calculate the number of H+ions added from the volume and concentration of the acid used. Hence we can determine the number of OH-ions in the original solution and their concentration.

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