Teach Me Chemistry - Lesson 3
Right about now in any intro chemistry course, you could you could be forgiven if you thought you dozed off and sleepwalked into the history class next door. Most books choose to introduce the fundamental subatomic particles by putting them in the context of their discoveries.
I. Early Chemistry
For example, In 1774, Antoine Lavoisier proposed the following:
It is always true that the total mass of substances formed by a chemical reaction is the same as the total mass as the reactants before the reaction happened. In other words, matter can neither be created nor destroyed in a chemical reaction. This is the Law of Conservation of Matter. It means that in any chemical reaction, if you can find a way to measure exactly how much of every reactant you start with, and exactly how much of every product you end up with, you would find the masses of each group added all together would be the same.
If AB + DC –> AC + BD, then Mass (AB + DC) = Mass (AC + BD)
In approximately 1805 Joseph Proust described an initial version of what we now know as the Law of Definite Proportions: All samples of a compound have the same composition - that is, they have the same proportion of elements by mass.
For example, let’s think of one everybody knows, H2O. Every water molecule has a ratio of 2 hydrogens to one oxygen. If we were to, say, add an extra oxygen to the molecule to make it H2O2, we would now have hydrogen peroxide - which you can use to bleach your hair or clean out your earwax, but you wouldn’t want to drink or swim in. Hereby we see that if we change the ratio of atoms in the compound, we change the identity of the compound. Sometimes dramatically.
Then there’s good old John Dalton and his Atomic Theory, put forward in the early 1800’s:
- Chemical elements are composed of atoms, which are not created or destroyed in chemical reactions.
- All atoms of the same element have the same mass and other properties, but are different from atoms of every other element.
- Atoms combine in simple, whole number ratios to form compounds.
p.s. John Dalton was an English teacher!
II. Subatomic Particles
The ancient Greeks thought atoms couldn’t be divided. Now we know they’re made of protons, neutrons, and electrons. (Which are themselves made of still smaller particles, but that’s a way more advanced lesson.)
The particles we know as electrons were first discovered by Michael Faraday about a century and a half ago, when he called them “cathode rays” because they emanated from the negative end of a cathode ray tube. J.J. Thomson took over from there and established their charge/mass ratio and declared they were fundamental particles of all atoms. Robert Millikan established the charge of the electron, which is -1.6 x 10-19 coulombs.
Ernest Rutherford, beginning in 1909, set out to learn something about the position of electrons in atoms and ended up discovering protons. In his famous “gold foil” experiment, Rutherford fired positively charged particles at a thin sheet of, well, gold foil. The results are as follows:
- Most of the particles passed straight through the foil undeflected.
- Some were slightly deflected; a few were severely deflected.
- About 1 in 20,000 bounced right back like a rubber ball hitting a wall.
One of the most famous quotes in chemistry belongs to Rutherford. After the gold foil experiment he said of the bounced-back particles, “It’s about as credible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.”
The explanation for this bounce-back phenomenon was that the positively charged particles were encountering another positive charge which was repelling them back the direction they came from. In other words, 1 in 20,000 of the particles Rutherford fired were hitting the nuclei of the gold atoms head-on.
This experiment led Rutherford to propose the nuclear atom, where the mass and positive charge are concentrated at the center, or nucleus, of the atom while the negatively charged electrons are outside of it. The magnitude of the positive charge is different for the atoms of each element, and is about one half the atomic weight. Atoms are mostly empty space and are electrically neutral.
III. Summary
Protons
Electric charge = 1.602 x 10-19
Atomic charge = +1
Mass = 1 atomic mass unit
Location = nucleus
Neutrons
Electric charge = 0
Atomic charge = 0
Mass = 1 atomic mass unit
Location = nucleus
Electrons
Electric charge = -1.602 x 10-19
Atomic charge = -1
Mass = 0.0005 atomic mass units
Location = outside the nucleus