Atoms - they make up everything we see & feel but what are they?

Like pretty much everything in science, the answer is "we don't exactly know". The best science can usually do is say "Well here is a model of how it probably works - it predicts lots of the things we see, but who knows, there may be details we don't understand yet.". Some of the models that have been proposed are considered here. There have been, and no-doubt will be, others but the gist should be pretty much right by now.

The word atom comes from the Greek word meaning "indivisible".
In fact as long ago as 500 BC, Democritus proposed that the elementary substances (thought to be earth, air, fire and water) were composed of tiny particles and gave them this name.

From a practical point of view everything is made of atoms. There are atoms in the air and in anything you can see or feel. Essentially anything large enough to count as a "thing" of any sort is made up of atoms. Atoms are the smallest parts that things can be broken up into by chemical reactions and the use of heat, light, electricity etc. Early chemists thus thought them indivisible and called them atoms.

If you really go mad and start sticking atoms in nuclear reactors or smacking them into each other at pretty much the speed of light then you can break them down further, but from a normal, everyday, point of view atoms are pretty much all individual, indestructable and impossible to interconvert.

What are they like?

The key thing about atoms is that they are very very small. So small in fact that if you lined up a row of hydrogen atoms one inch long, you would need around 320 million of them.* Unlike Democritus, we now believe there are several hundred types *, forming the hundred or so known chemical elements. Because they are so small, you need a massive number of them to make anything worthwhile. In fact 360 ml of water (that's about a glass-full) contains around 12,000,000,000,000,000,000,000,000 oxygen atoms and twice as many hydrogen atoms!

In the early 1800's John Dalton showed that you could determine the relative masses of different atoms and by 1816 William Prout had proposed that the masses of all atoms were whole-numbers and whole-multiples of the mass of hydrogen. This model was rubbished when chlorine was discovered to have a mass of around 35.5 hydrogens but, as we will see, it's a better approximation than some people thought then.

What are they made of?

Daft though it may sound, the thing that is most important about these apparently indivisible atoms is what they are made of!

There are three basic building blocks of an atom (called sub-atomic particles); protons, neutrons and electrons. Their properties are approximately as follows:

As you can see, the mass of an atom resides (almost) entirely in the protons and neutrons, while the charge exists as positive protons and negative electrons.

If a normal, stable atom had an overall positive or negative charge then all hell would break loose! Electrostatic forces would make them repel each other massively and the whole place would pretty much fly apart. Thankfully, atoms usually have more or less the same number of protons and electrons, so the net charge is zero. Where they have a few electrons more or less, there is usually another atom near-by with the opposite charge to balance this out.

The thing that makes one atom different to the next is the number of protons and neutrons that it has. The total of these alters the mass of the atom and the number of protons controls the number of electrons and these in turn control an atom's properties. The number of protons in an atom is called its "atomic number" and the total of the protons and neutrons together is called its "atomic mass".

A type of atom with a particular atomic number (i.e. number of protons and therefore number of electrons) is known as an "element". The name derives from the fact that these were thought to be the elementary building blocks of matter.

Clearly you could have a certain number of protons but change the number of neutrons. This would make the same element but the atoms would be heavier or lighter (i.e. same atomic number, different atomic mass). Two or more atoms with the same atomic numbers but different atomic masses are called "isotopes". These tend to be pretty much identicle in chemical properties but vary slightly because they have different masses.

For reasons that are hugely complicated, only a certain proportion of neutrons to protons tends to be stable and so you can only get a few stable isotopes of each element. The less stable ones are radioactive and break up into smaller bits, the really unstable ones just don't stick around long enough to be measured. As atoms get heavier, they tend to take in a slightly higher proportion of neutrons.

So William Prout wasn't that far out when he said that all atoms were whole-multiples of hydrogen. It turns out that hydrogen * has one proton and no neutrons so its mass is 1 (from the table above). What was happening in chlorine was that natural chlorine is a mixture of isotopes, one with 18 and one with 20 neutrons (both have an equal number of protons obviously - 17 in this case). There is about 3 times as much chlorine-35 as there is chlorine-37 so the average mass was 35.5.

In 1869, Dmitri Mendeleev published a short article called "The Correlation Between Properties of Elements and their Atomic Weights*. This documented how the properties of the known elements seemed to vary in a regular cycle with their atomic mass *. This arrangement was eventually developed into the modern "Periodic Table of Elements" and, as we will see, was the first indication of an emerging pattern on the sub-atomic scale. How the sub-atomic particules are arranged within the atoms is the key to this.

Atomic Structure

There have been quite a number of models of atomic structre over the years. Every time someone comes up with an experiment showing results that can't be explained by the current model, a new one has to be devised.

Before we knew anything about the prtons and neutrons that make up atoms, Sir Joseph Thomson devised and experiment to show that electrons existed. This lead him to a model of the atom as a sort of drop of fluid positive charge with negative electrons dotted around. It was known as the "plumb pudding" model because the positive chage was like a pudding containing electron "plumbs"!

Some years later, Rutherford fired very high speed radiation at a very thin metal film and observed that a small amount bounced straight back. He is quoted as saying "It was 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!". Rutherford's explanation for this was that rather than being a uniform drop, the atom had all the positive charge in the middle (called the nucleus) with the electrons around the outside. This was the "planetary" model in which the positive charge was the "sun" and the electrons orbited like "planets" round the outside. There was a problem with this model as well though because, according to electromagnetism, a negatively charged electron orbiting an atom would give out raditation and quickly spiral inwards, crashing into the nucleus. Another theory was needed.

It was Neils Bohr who suggested the almost unthinkable - if Rutherford's atomic model and classical elecromagnetism did not agree, maybe it was elecromagnetism that was wrong! He combined Rutherford's model with the newly emerging theory of "quantum mechanics" to suggest a model in which electrons occupied stable orbits and emitted radiation only when they changed form one orbit to another.

Bohr's explanation was that the electron orbiting the atomic nucleus had angular moomentum* and that this angular momentum could only take certain values. By allowing only certain orbits to contain electrons he could side-step the classical idea that they would lose energy and spiral inwards. Instead, radiation with a specific energy would be lost or gained when the orbit changed. Bohr's orbits were circular and each was given a number. This number "n" is an integar known as the principal quantum number.

Later, further quantum numbers were proposed to explain the details observed in the spectra of the known elements. These were "l", which could take any value between 0 and n-1, "m" and which could take any value from +l to -l. These correspond to the various possible shapes of non-circular orbits around the atom. Finally, each eletron had a "spin", which represented the two directions in which an electron could be "spinning" around it's own axis*.

The various combinations of these quantum numbers can be used to explain a great deal about the periodic table and the properties of the elements. For example, when n=1, l and m can only be 0 and the spin can be up or down, so there are only two possibilities. These correspond to the two elements in the first period of the periodic table - Hydrogen with 1 electron and Helium with two.

The question remained however - Why do these electrons stay in their discrete orbits rather than spiraling inwards? The answer to this came in 1923 when DeBroglie suggested that electrons could show wave particle duality, like light. Within a few years wave mechanics had been developed which described the various "orbits" of the electrons as solutions to a wave-equation. Each orbital as they are known was a solution giving a stationary "wave" of electrons around the nucleus. These orbitals confined the electrons into various regions without requiring them to move and so removed the conflict. One solution to this wave equation was provided by each combination of quantum numbers giving various shells of orbitals. How the electrons arrange themselves in these orbitals is described in
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