Introduction to Chemistry

Chm 1311 Lecture for 6 June 2000

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The course presumes that you've already had an introduction to Chemistry (however long ago); so it will skip the first chapter of Brady and Holum except for summary remarks on this the first class period. Nevertheless, the first exam and the final will consider the content of Chapter 1 fair game!


Chemistry is sometimes called "the central science" lying as it does between atomic physics and the macroscopic world of Life. From our perspective as living beings, Chemistry is the connection between the fundamental particles and fields and the workings of cells. But Chemistry isn't a bridge that leads only to Biology; it informs all of the other sciences and all of the world of energy and manufacture.
For Chemistry is the science of building blocks and their transformation: blocks which construct all of the matter of the universe. So if it's not the Science of All That Matters, it's at least the Science of All That's Matter! Moreover since some of these transformations involve Light, Chemistry directs its attention to the interaction of Light and Matter as well.
Chemistry, in common with all the Sciences, is based upon observation and inference. Those bulwarks of the Scientific Method prepare one to postulate hypotheses (theories about the workings of the material universe) which, however tentative, suggests new, possibly corroborative, observations to be made and inferences drawn therefrom.


Observations are recorded using a worldwide unit system called "Système International" whose basic units are the meter (m), the kilogram (kg), and the second (s) as well as the Kelvin (K), a unit of absolute temperature. There are also fundamental SI units for current, luminosity, and other measures to be discussed later, but one of critical importance to Chemistry is the measure of amount or the standard number of atoms or molecules; that unit is the mole (mol) and consists of 6.0221367×1023 molecules.
Given such massively large numbers, the system of scientific notation permits us to specify the scale with simplicity; so the mole is about 6 followed by 23 zeroes in decimal notation. Still more convenient are new names for each power of 1000 (103) such as kilogram meaning 103 grams. Or milligram (mg) as 10-3 grams. The Periodic Table given to you on the first class day has on its reverse a table of fundamental constants, conversion factors, and a list of those decimal multipliers. Correct usage is to use a multiplier which permits the magnitude of the number to lie in the range of 1 to 999; thus the speed of light in a vacuum, 2.99792458×108 m s-1 could be expressed also as 299.792458 Mm s-1 since mega (M) means 106. But since these decimal multipliers are merely shorthand, there's nothing wrong with the full scientific number expression.
Two multipliers not associated with multiples of 1000 are the common centi (10-2) as in centimeter (cm) and the less familiar deci (10-1) as in decimeter cubed (dm3). The latter means the cube of a decimeter, (10-1 m)3 = 10-3 m3 = 103 cm3 = 1 L, another favorite Chemistry measure, the liter.
Some SI
Giga, G Mega, M Kilo, k deci, d centi, c milli, m micro, m nano, n
10 9 10 6 10 3 10-1 10-2 10-3 10-6 10-9
Derived SI units are those obtained by mathematical combination of the fundamental SI units. Since (from Physics) kinetic energy is (½)mv², the SI unit of energy, the joule, is J = kg m2 s-2.
Another way energy changes, DE, is via changes in temperature, DT. This kind of energy is called "heat," and it is expressed as

DE = n CV DT

where n is the amount of the material involved (mol) and CV must have units of J mol-1 K-1.


Since the heat scales with the amount of material, n, this property of matter is called extensive, i.e., it increases or decreases with the amount of material present. Heat is extensive but temperature is intensive; if we add together two identical blocks heated to 400 K, the heat doubles, but the temperature remains 400 K. Such properties, invariant to change in the amount of material, are called intensive properties.


Finally, regardless of the properties being recorded, your calculator will be far more generous with its digits than you should be! It will presume that the digits you enter are known absolutely precisely...or at least to the precision (usually 10 places) retainable by the calculator. You know better. A measurement of 2 m is far less precise that one of 2.000 m; the former is known to only one significant figure whereas the latter is known to 4. All of those zeroes, if present, were meant to be taken as seriously as the 3s in 2.333 m. On the other hand, the zeroes in 0.0002 are only placeholders; so 0.0002 has only one significant figure.
An easy way to enlighten yourself about when zeroes are significant figures and when they aren't is to convert numbers to scientific notation. Then 0.0002 becomes 2×10-3, and it's obvious there's only one significant digit there. For larger numbers, 2.0×103 means 2000 where only the first two digits are significant.
In use, multiplication and division results should be quoted to no more than the least significant number of digits of their arguments. So 1.07×0.006 = 0.006 regardless of the first multiplier's superior precision.
Under addition and subtraction, the results are significant to the largest decimal position still significant. So 1.07+0.006 = 1.08 since we must round for greatest accuracy when fixing significant digits. Indeed, with multiple math operations, it is legitimate to retain as much precision as your calculator will permit all the way to the final answer, but then you must trucate it as appropriate to the correct number of significant figures!

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Last modified 4 June 2000. Chris Parr