Models for the atom

Before any experimental evidence, the existence of atoms was posited by ancient philosophers as a consequence of the idea that there must be some limit to the divisibility of matter into smaller and smaller portions. Democritus (466-370 BCE) is usually credited with developing the concept during a period in which classical Greek civilization gave rise to an interest in natural philosophy. Democritus' atoms (the term was derived from the Greek atomos, meaning "uncuttable") were solid, indivisible, and indestructable. Furthermore, the atoms could stick together upon collision, and such agglomeration of atoms gave rise to the variety of substances observed in the world^{1, 2}. Remarkably, these speculative propositions bore a striking resemblance to Dalton's atomic theory, put forward 1200 years later. As chemistry developed in the 18th century, indirect evidence began to accumulate for the existence of atoms and for the proposition that the variety of pure substances is a result of the tendency of atoms to form different stable combinations in definite proportions. In other words, atoms combine to form molecules, which are thus the smallest unit of the pure substances known as compounds. By the first decade of the 19th century, the English scientist John Dalton had proposed the first recognizably modern version of atomic theory. Let's outline the development of atomic theory, and in so doing we will arrive at a model for the atom that will prove to be of great utility in understanding much of chemistry.

Dalton's atomic theory of matter. Dalton's postulates⁴ were as follows:

1. All matter consists of solid and indivisible atoms
2. All atoms of a given element are identical
3. Different elements have different kinds of atoms, differing in mass
4. Atoms are indestructible
5. Compounds are formed by combination of different atoms in small proportions

As we'll see, postulates 1, 2, and 4 would need to be revised in the light of later discoveries, but Dalton's theory was a major advance.

The Rutherford model. The indivisibility of atoms was the first of Dalton's postulates to require revision. Experiments throughout the 19th century suggested atoms were made up of smaller particles. The electron, a negatively charged particle very much smaller than the atom, was the first to be described and measured. Since atoms themselves are electrically neutral, if electrons are found in atoms, then there must also be some positively-charged component to balance the charge within the atom. In Rutherford's model, most of the mass and positive charge of an atom is centrally concentrated in a nucleus. In contrast, electrons were somehow diffusely distributed around the nucleus to define the size of the atom, but overall, most of the atomic volume was simply empty space. The nucleus is on the order of 10⁻¹⁵ m and an atom is on the order of 10⁻¹⁰ m in size. In visualizing the relative scales of nuclear and atomic sizes, images such as the "fly in the cathedral" or the "pea in the stadium" are commonly invoked. Subsequently, Rutherford showed that the nucleus of the hydrogen atom (the smallest and simplest nucleus) consisted of a particle with a unit of positive charge (exactly equal but opposite to the electron), and which makes up the positive charges of all other nuclei. Rutherford named this particle the proton. Later, the neutron was discovered as a second nuclear component, with a mass approximately equal to the proton (electrons are more than 1800 times lighter than protons and neutrons). All nuclei were determined to be made up of protons and neutrons.

At this point, let's summarize the properties of the subatomic particles that make up the atom in the following table.

Table 1: Subatomic particles
	Particle	Symbol	Mass (g)	Mass (amu)*	Charge†
	Proton	p+	1.672623 × 10−24 .	1.00728	1 +
	Neutron	n	1.674929 × 10−24 .	1.00866	0
	Electron	e−	9.10939 × 10−28 .	5.48680 × 10−4	1 −
	* Atomic mass unit, amu is defined as exactly 1/12 the mass of a carbon-12 atom; 1 amu = 1.660539 × 10−24 g. † Charges are given in elementary charge units, e  =  1.602176 × 10−19 C  (coulomb, C), which is the charge of an electron. The electron shell model This model can be thought of as building on Rutherford's by describing the distribution of electrons about the nucleus. Due to electrostatic potential energy, the electrons in an atom all experience an attractive force due to the positive charge of the nucleus, and simultaneously there is a repulsive force between electrons. In the electron shell model (or as we'll on occasion refer to more simply as the shell model), the electrons are organized into shells - analogous to the layers of an onion - with the electrons in shells close to the nucleus referred to as core electrons, and those in the outermost shell are called valence electrons. The electron shell model is a chemical model for the atom inasmuch as it is the valence electrons of various elements that greatly influences their reactivity.

The Diagram of the shell model, showing the coarse internal structure of the atom. At the center is the dense nucleus, containing the protons and neutrons, and accounting for nearly all of the mass of an atom. Surrounding the nucleus are the diffusely distributed electrons, which are pictured as organized into shells. The core electrons are those in the innermost shells and are found for the most part close to the nucleus. The electrons in the outermost shell are the valence electrons. Note the size of the nucleus relative to the atom is vastly exaggerated in this figure.

The valence electrons are located much further on average from the nucleus, so they would be less strongly attracted to it by its attractive electrostatic potential energy. Thus, they more freely participate in electron donation or sharing in chemical bonding between the atoms in molecules and molecular ions. The electron shell model is quite successful at explaining the chemical bonding behavior of different atoms by focusing attention on the valence electrons.

The shell model and chemical periodicity. A shell model for how the electrons are organized within atoms would seem to suggest several features.

First, the electrons within any given shell would all be located at about the same distance from nucleus - at least on average. Second, this average distance is significantly different for each shell, increasing as the number of shells increases. This follows from what we already noted for valence electrons, the outermost shell being furthest from the nucleus. Third, as shells are added in order to accommodate the electrons in atoms of larger atomic number, the number of electrons that can fit in a shell should increase.

These features can account in general for the structure of the periodic table and the periodic variation in atomic sizes (see chart showing elements 1-20). The structure of the periodic table is consistent with a first, innermost shell that can hold only two electrons, a second and third shell that hold eight electrons each, a fourth and fifth that hold 18 electrons each, and so on. If we look at the sizes of atoms (these can be experimentally determined by a technique called X-ray diffraction) we see that the atoms of elements in period 2 are much larger than those in period 1, and significantly smaller than those of period 3. Interestingly, atomic size actually decreases within a period, from left to right. Does the shell model offer a possible explanation? (Hint: Consider the effects of electrostatic potential energy.)

Another important feature of the electron shell model that is not immediately apparent is that filled shells are a stabilizing factor for atoms and ions. This is a necessary postulate to address the remarkable lack of chemical reactivity observed for the Group VIII elements, the noble gases. While this is a rather ad hoc postulate, it is a useful one for rationalizing chemical properties, even of elements other than noble gases. In this view, elements exhibit a strong tendency to combine chemically in ways such that each atom within the molecules or ions formed attains a filled-shell electron configuration, like that found in isolated atoms of a noble gas. Atoms achieve this by gaining, losing, or sharing valence electrons. For instance, the chemistry of Group I and VII elements is largely accounted for by the loss and gain, respectively, of one electron to become ions. In either case, the ion formed has filled shell electron configuration.

The quantum mechanical atom. 

The most modern model, the quantum mechanical view of the atom accounts for the wave-like properties of subatomic particles. In particular, this model explains why the electrons in atoms behave as if they were organized into shells.

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CHEM 101 - Atoms: An introduction. (2018). Guweb2.gonzaga.edu. Retrieved 11 October 2018, from http://guweb2.gonzaga.edu/faculty/cronk/CHEM101pub/atoms01.html