[Date Prev][Date Next][Thread Prev][Thread Next][Date Index][Thread Index]

**To**:**<aquatic-plants at actwin_com>****Subject**:**Counting by Molecules vs Counting by Mass****From**:**"James Purchase" <jppurchase at rogers_com>**- Date: Sat, 23 Nov 2002 08:45:44 -0500
- Importance: Normal

Steff asked: "Can someone straighten out for us all, when to count by molecule and when to reckon by weight?" You're confused. You're not alone. Sometimes basic concepts get glossed over or forgotten in the flow. Think of it this way - you want to make an omelet, but you're out of eggs. You go off to the store to buy what? Eggs are (usually) sold by the "dozen", a standard "number" of eggs (or actually, a standard number of "anything"). So, you pick up a "dozen" eggs and go home to make your omelet. You have 12 eggs in that "dozen" and depending upon your appetite, you start to break some shells. But enough of the cooking analogy, I'm getting hungry..... Molecules and atoms are extremely tiny things. As a commedian once observed, they are eenie, menny, minny (except I don't think the comic was referring to atoms). It would be impractical to actually "count" out atoms or molecules in everyday chemistry. Chemistry labs the world over would have to be awfully quiet places - imagine if you got distracted and lost count after counting out 1,456,328,673,967 atoms? You'd have start counting all over again from 1. A 19th century Italian scientist by the name of Amadeo Avogadro formulated a hypothesis which held that at a given temperature, an equal volume of any gas contains the SAME number of individual molecules as any other gas. The "number" of individual molecules is 6.02 x 10^23 (that's a 10 followed by 23 zeros). That's a BIG number. By inference, that means that individual molecules are REALLY small. In the year's since Avogadro first postulated his hypothesis, it have been recognized and accorded the status of a "law" - Avogadro's Law. It can be applied to anything, not just gases. The "number" he came up with, 6.02 x 10^23, is referred to as "Avogadro's Number", and has proved to be very handy as a way of "counting" anything that is really tiny. An "Avogadro's Number" of anything contains 6.02 x 10^23 individual particles (atoms or ions or molecules). Like the eggs you bought to make your omelet, this "number" has a easy to remember "name". It is referred to as a "Mole" (whoever came up with the names of these things probably didn't get out much). A "Mole" of any substance has the SAME number of individual particles as a "Mole" of any other substance. If you take two identical ballons, and fill one of them with Hydrogen gas and fill the other with Oxygen gas, they will both contain the same number of molecules of their respective gases. But if you tie them off and let them go, one will float up to the ceiling and the other will fall gently to the floor. The balloon filled with Hydrogen will float in air because the balloon doesn't weigh as much as an equivalent volume of air. Similarly, the balloon filled with Oxygen will sink to the floor because Oxygen weighs more that the air displaced by the balloon. This "weight" is actually the force of the Earth's gravity acting on the mass of gas within each balloon. Since we know that a Mole of any substance contains exactly the same number of individual particles as a Mole of any other substance, any difference in their "weight" MUST be attributed to the MASS of the individual particles in each Mole. A "Mole" of Hydrogen gas (6.02 x 10^23 atoms of H2) "weighs" 2.01588 (1.00794 x 2) grams, while a "Mole" of Oxygen gas "weighs" 31.9988 (15.9994 x 2) grams. In both cases, this is merely the Atomic Weight (mean relative mass) of the element multiplied by the number of atoms in the molecule. It is much easier and much more practical to determine the mass of a substance than it is to be counting quantities contining many orders of magnitude. When dealing with chemical compounds, you can easily determine the "weight" (actually the mass) of a Mole of the compound by summing the Atomic Weights of the individual elements in the compound. This is called the "Gram Formula Weight" of the compound. If you want a Mole of Potassium Nitrate (KNO3), you would measure out (on an accurate scale) 101.10 (39.0983 + 14.00674 + 15.9994 + 15.9994 + 15.9994) grams of the crystals. This "amount" of KNO3 contains 1 Mole of KNO3. You can also see that it contains 1 Mole of Potassium, 1 Mole of Nitrogen, 3 Moles of Oxygen and 1 Mole of Nitrate (NO3). If you dissolve this in 1 liter of water, guess how many parts per million K are in each cubic centimeter of the resulting solution? The element Nitrogen has an Atomic Weight of 14.00674. (that is a Nitrogen atom, not a molecule of Nitrogen gas). When combined with 3 atoms of Oxygen (each of which have an Atomic Weight of 15.9994), you get the compound Nitrate (NO3). A Mole of Nitrate would weigh 62.00494 (14.00674 + 15.9994 + 15.9994 + 15.9994) grams. When Roger referred to the factor of 4.4 when comparing Nitrogen-N with Nitrate, this is what he was talking about. So to cut this off (finally!) and to answer your initial question - "when to count by molecule and when to count by weight?" - the answer is that counting by weight is the same thing as counting by molecule, IF you stick to counting Moles of substances. And that is basically what we are doing here. James Purchase Toronto

- Prev by Date:
**Potassium source...** - Next by Date:
**What next?** - Previous by thread:
**Potassium source...** - Next by thread:
**What next?** - Index(es):