Tuesday, September 14, 2010

law of definite proportions


As I was preparing yet another lecture for my introductory chemistry class I realized something: this class is all about grasping the law of definite proportions. Once the fundamental concept of relationships in chemical quantities is realized and accepted, students can propel themselves into much bigger and sophisticated equations and chemical concepts.

What is the law of definite proportions? Let's go to my friend Wikipedia for a brief description. (I realize some readers may have an aversion to Wikipedia but I go there when my textbook is otherwise preoccupied with another instructor as it is currently.) Wikipedia is generally correct although I will admit there are some QA issues with the site.

"In chemistry, the law of definite proportions and also the elements, sometimes called Proust's Law, states that a chemical compound always contains exactly the same proportion of elements by mass. An equivalent statement is the law of constant composition, which states that all samples of a given chemical compound have the same elemental composition. For example, oxygen makes up 8/9 of the mass of any sample of pure water, while hydrogen makes up the remaining 1/9 of the mass. Along with the law of multiple proportions, the law of definite proportions forms the basis of stoichiometry."




This little paragraph nicely describes my overall definition of the law of definite proportions.

This scientific law was incorporated into Dalton's Atomic Theory in the early 1800s. Now it probably just blends in as a fundamental tenant of atomic theory, I doubt most people know it originated with Proust.

While Atomic Theory (especially modern atomic theory) is important for introductory chemistry, the class can be defined at a more fundamental level as applicable to the law of definite proportions.

Why? Here are the reasons: The law of definite proportions explains why the following are true.

1. The relationships between amu mass/molecule is the same as the number of grams in one mole of something. A strange and difficult relationship to swallow especially when it would seem experimentally impossible to prove this. I'm not really sure how they did it. (If anybody reads this and knows please fill me in)
2. The manipulation of numbers between grams, moles and particles depends on the relationships in #1 (previous bullet). Without these relationships, the theoretical yield of a chemical reaction cannot be calculated and compared with the actual experimental value. Hence, without this concept there is no percent yield calculation.
4. While the rest of atomic theory is critical to chemistry it does not fundamentally affect calculations (for which we grade the students). 
5. None of the other tenants of the theory allow for such simple yet profound manipulation of the experimental data to give predictions, results and information!



Here is the Wikipedia definition for law of multiple proportions:
If two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small whole numbers.[4]
For example carbon oxide: CO and CO2, 100 grams of carbon may react with 133 grams of oxygen to produce carbon monoxide, or with 266 grams of oxygen to produce carbon dioxide. The ratio of the masses of oxygen that can react with 100 grams of carbon is 266:133 ≈ 2:1, a ratio of small whole numbers. The Law of Multiple Proportions, is just what the name suggests, the law of multiple proportions of one constant element within differing compounds sharing the same type of chemical bonding.

Last semester our learning objectives were to:
1. Calculate and predict theoretical yield in a chemical equation. Use this with experimental data to calculate percent yield.
2. Use Lewis theory to build molecules and predict shapes. Use shapes to predict molecular behavior.
3. Predict behavior of atoms based on periodic table trends

#1 is by far the most important components of this class.