Atom (science)

An atom (from the Greek atomos, indivisible) is the smallest physical entity that can retain its chemical properties. Moreover, each chemical element is a weighted average of the properties of all the atoms with the same number of protons. As such, atoms are the building blocks of all the matter we come in contact with on a daily basis. Individually atoms are so small, $10^{-10}$ that it takes literally trillions of them to form most objects we are familiar with. In fact, atoms were once thought to be the fundamental building blocks of the entire universe. We now however, understand that atoms are made up of smaller subatomic-particles.

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Introduction

The idea of an atom dates back to the ancient Greeks who postulated that all matter in the universe was made of tiny indivisible chunks. While this concept had no empirical support at the time, the search for the most fundamental "chunks" in the universe has remained one of most important the driving forces in science to this day. The particles that today, bare the name "atom" were first encountered by early chemists who discovered that they could not be broken down nor transformed through any chemical reaction.

The modern atom serves as the fundamental particle for chemistry and as the smallest stable structure for engineering. Nuclear energy is generated either by breaking an atom into two or more smaller atoms fission, or by combining two smaller atoms into a larger one fusion.

History of the Atom

The nineteenth century

In the beginning of the nineteenth century men as John Dalton, Joseph-Louis Gay-Lussac and Amedeo Avogadro hypothesized that matter consists of minute particles, atoms as Dalton liked to call them, or molecules (a word coined by Avogadro). However, the difference between the two concepts was not clear. One of the main topics of the first international chemistry conference, the historic 1860 Karlsruhe conference, was to clear up the confusion about the difference between atoms and molecules. Stanislao Cannizaro offered for the first time in the history of science, a very clear definition of atoms as distinguished from molecules. To him the atom was the smallest quantity of each element which enters as a whole into the molecules which contain it. His statements represented a most remarkable contribution to the clarification of the issues debated at the time concerning the relations between atoms and molecules in both organic and inorganic compounds.

J.J. Thomson's "Plum Pudding" Model

The first widely accepted model of the atom was J.J Thomson's Model of a positively charged cloud with negatively charged "corpuscles," electrons, interspersed throughout. Thomson proposed and searched for configurations for which these particles had normal modes of vibration and were stable. He built his model using only electrostatic forces, which required that the electrons be in constant motion, this led to further difficulties and he was never able to create a model that matched observed data.

Rutherford Model

Ernest Rutherford, one of Thomson's students, disproved his theory in his now famous scattering experiment by showing that there was a dense, positively charged core in each atom. Rutherford postulated that the structure of an atom more closely resembled a solar system with the nucleus at the center and electrons orbiting around it. Unfortunately, Classical Mechanics predicts that these orbits were unstable and would decay in less than a microsecond, while emitting a continuous spectrum of light. However, it was already known that each element emitted a unique spectrum of discrete lines. Rutherford could not explain why these orbits did not decay nor why atoms did not emit the continuous spectrum that these decays would have caused.

Bohr's model of the Hydrogen Atom

In 1913 Niels Bohr used the Schrödinger equation to produce the first quantum mechanical model of an atom. He resolved the difficulties of Rutherford's model by making the non-classical assumption that there were only certain specific energy states allowed for these electrons. Bohr conjectured that the angular momentum of each electron was constrained to be an integer times Planck's constant divided by 2π, $\hbar$. Since, electrons could now only change energies between these accepted states, the discrete spectral lines could be explained by photons of equal value to the energy change between the acceptable states.

Bohr was only able to explicitly solve the Schrödinger equation for electrons orbits the hydrogen atom, though he conjectured that larger atoms would have a similar structure.

Today's view

For most practical purposes Bohr's model has proven an acceptable approximation. However, the quantum mechanical interpretation is that the electrons are spread out in various probability distributions around the nucleus rather than actually orbiting it.

Structure of the Atom

Atoms are made of a dense nucleus formed by combinations of the two nucleons (positively charged protons and zero charge neutrons) and surrounded by a much larger "cloud" of electrons. The number of protons contained in the nucleus determines the atomic number and in turn which element it is classified as. The number of neutrons further specifies the isotope number of that element. The number of electrons surrounding the nucleus is typically assumed to be equal to the number of protons in order to keep the entire atom electrically neutral. Atoms that are not neutral are called ions, they are designated by their charge in units of elementary charge, which is equal to the negative of the number of surplus electrons present around the atom. Thus an atom with one extra electron is charged -1, and one missing is charged +1.

Forces in the atom

The nucleus of the atom contains a high concentration of positively charged particles with no counter balancing negatively charged particles to keep it stable. In order to explain the existence of the nucleus scientists introduced the two nuclear forces, the strong force and the weak force. We now understand that the nucleus is held together by the residual strong force despite its significant positive charge. The electrons which surround the atom are electro-statically attracted to the nucleus due to their negative charge.

Electron Quantum States

Electrons surround the nucleus of an atom, where each is in a unique quantum mechanical state in the atom's structure. These states of well defined energy are called atomic orbitals and are organized from the lowest energy to highest. Each orbit can be uniquely defined by four quantum numbers, 'n' the shell, 'l' the angular momentum, 'm' the magnetic angular momentum, and 's' the spin. Shell numbers can only be natural numbers but no shell number over 7 has ever been observed. Each shell has n² spatial orbitals which are grouped by angular momentum 'l' into subshells. These subshells are often given the names 's', 'p', 'f', 'g', and 'h' instead of numerals. Each of these subshells contains ((2*l)+1) suborbitals 'm' which are numbered (−l...,0,...l). Finally, each suborbital can contain two electrons of different spin $+{\tfrac {1}{2}}$ and $-{\tfrac {1}{2}}$. For a total of 2n² electrons per shell.

The valence, or the highest energy shell which contains electrons, is responsible for most of an elements chemical properties.

An atom with electrons in higher energy states without first filling lower states is considered "excited" and is likely to emit this excess energy in the form of a photon of energy equal to the difference between the current state and a lower energy state until it is no longer excited. These states are assembled into groupings called "shells" which have great importance in determining the chemical properties of an element.

Nuclei Quantum States

The nucleus of an atom itself also has a shell like behavior not dissimilar to electron shells. This leads to some extra stable and semi-stable states when the number of protons or neutrons in the nucleus is equal to the "magic numbers" 20, 28, 50, 82, and 126 which is believed to correspond to complete shells.

Decay

Most combinations of nucleons are inherently unstable and undergo a number of radioactive decays in order to form more stable nuclei. In all of the common decays a particle is emitted from the nucleus in order to compensate for some instability in the atom.

• Alpha decay most frequently occurs in atoms which are simply too big, atoms are limited in size because the residual strong force which holds them together only acts over very small distances so that the rate of electrostatic repulsion grows faster than the rate of strong attraction as the nucleus grows. Alpha decay emits an Alpha particle which is denoted with the Greek letter α.
• Beta+ decay occurs in atoms which are proton heavy. In this decay a neutron decays into a proton and emits both a positron and a neutrino. The only natural Beta+ emitter is ⁴⁰K.
• Beta- decay occurs in atoms which are neutron heavy. In this decay a proton decays into a neutron emits both an electron and a neutrino. A lone proton can beta- decay.
• Gamma decay occurs in atoms where the nucleus is in an excited state. The nucleus will turn this excess energy into a high energy photon, typically called a gamma ray. Gamma decays typically follow either an alpha or beta decay which leaves the nucleus with excess energy.
• Electron capture occurs when a proton inside the nucleus captures an electron and becomes a neutron.

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