Particles and forces
Before discussing what happened in the immediate aftermath of the bang (because the bang itself falls outside the scope of present day physics), it is appropriate to say something about the particles of nature and the forces which hold them together and govern their behaviour and their provenance.
The word “atom” derives from the Greek atomos: ‘that which cannot be cut”, and the Greek Democritus was among the first who (in the fifth century BC) supposed matter to consist of vast numbers of indivisible atoms ceaselessly moving in a vacuum. This ancient heritage notwithstanding, at the beginning of the twentieth century the very existence of atoms was still a subject of dispute in the scientific community until in 1905 Einstein and the Australian William Sutherland both independently demonstrated that their existence could be verified. After introducing the concept of an electrical force between atoms, Sutherland went on to use this idea to work out how to determine the size of atoms, and how different atoms might mix together [1].
By the 1930s, 1932 to be quite exact, understanding had progressed to the point where it was known that, far from being the most elementary constituent of matter, atoms consisted of a nucleus, containing protons and neutrons, surrounded by a swarm of orbiting electrons. The path by which this came about is interesting.
The electron was first conceived in theoretical terms in 1892 by the Dutch physicist Hendrik Antoon Lorentz, and in 1897, Joseph John Thomson demonstrated experimentally that electrons really did exist, so one might say that the electron was conceived in 1892 and delivered in 1897 [2]. Thomson was using a vacuum tube to investigate the nature of cathode rays, measuring to what extent magnetic and electric fields deflected them. As the rays passed through these fields, they bent towards a positively charged electric plate, so Thomson realised that they must be negatively charged particles (only later were they called electrons) coming from inside the atoms of the cathode, thereby providing the proof that the atoms themselves were made up of smaller parts hidden within them. However, since it seemed impossible to extract electric charge from an atom, and yet atoms as a whole were known to be electrically neutral, Thomson suggested that the negative electrons were moving around inside a uniform sphere of positive charge that accounted for most of the atom’s mass. This caused him to devise his so-called ‘plum pudding’ model of the atom with randomly dotted negatively charged electrons in a positively charged cloud-like sphere that contained most of the atom’s mass. In this model, there was no nucleus [3].
The electron was therefore the first atomic particle to be discovered. Then in 1909, Hans Geiger and Ernest Marsden, working under Ernest Rutherford, fired alpha particles at a thin sheet of gold leaf, leading to the finding that atoms actually concentrated their mass and positive charge in a small area at the centre – the nucleus, and in 1911 Rutherford proposed that the electrons orbited the nucleus of an atom like moons around a planet. In 1918, Rutherford identified particles in the atom with a positive electric charge that was equal and opposite to the electron’s negative charge. He had discovered the proton, but they did not account for an atom’s entire mass – the remainder seemed to lack any electric charge at all. Rutherford named these chargeless particles neutrons.
In July 1913 Niels Bohr, one of the leading lights in the emerging discipline of quantum mechanics, published a seminal paper [4] postulating that:
- electrons orbiting the nucleus are in no way similar to or comparable with planets orbiting a sun - which travel in nice regular orbits - because they are only capable of occupying well defined orbits with certain allowed energies, and when in such orbits, they do not radiate energy;
- within that framework they can flit from one orbit to another by absorbing or emitting a single photon without visiting the space in between – a quantum jump or leap - the photon's energy being the difference between the two orbital energies.
These well-defined orbits prevented the electrons from falling into the nucleus. Instead they appeared like a cloud, “a zone of statistical probability marking the area beyond which they only very seldom stray”, or like the whirring blades of a spinning fan, “seeming to be everywhere at once while filling every bit of space in their orbits simultaneously” - but the blades only seem to be doing this, while electrons - comprising a collection of properties like mass, electric charge and with no apparent physical size - actually do [5]. So the physics world came to realise that an atom was not like some monolithic billiard ball, but instead an elegant mechanism with several moving parts: a nucleus consisting of protons and neutrons orbited by a swarm of electrons – “not unlike a hive with its restless swarm of bees” [6].
It took another 14 years after Rutherford’s discovery of the proton and his theoretical construct of the neutron before the neutron’s existence was finally proven by James Chadwick in 1932 [7]. In the same year, using the first particle accelerator, the British physicists John Cockroft and Ernest Walton discovered the source of the energy that bound protons and neutrons together in the strong nuclear force, replete with an energy loss duly accounted for, in due course, by E=mc². [8]
1932 was a seminal year in nuclear physics, generally regarded as its birth year. The unsatisfactory proton and electron model of the nucleus, under which:
nuclei consisted of how ever many protons were required to account for their mass, plus how ever many electrons were needed to reduce the net positive charge to be equal to the number of electrons that were outside the nucleus [9],
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was now replaced by a proton-neutron model in which the number of protons in the nucleus equals the atomic number (and also the number of extra-nuclear electrons) and neutrons make up the difference between the element’s mass number (the total number of protons and neutrons in its nucleus) and its atomic number (its position in the Periodic Table [10]).
Protons and neutrons are almost the same size as each other, and much larger than electrons, and in this context Dawkins and Wong elaborate upon the process of ionization, what it means and the significance of the numbers of protons and neutrons in the nuclei of the different elements. Thus:
Protons and neutrons are almost the same size as each other, and much larger than electrons, and in this context Dawkins and Wong elaborate upon the process of ionization, what it means and the significance of the numbers of protons and neutrons in the nuclei of the different elements. Thus:
"Unlike a neutron, which is electrically neutral, each proton has one unit of electric charge (arbitrarily designated positive), which exactly balances the negative charge of one electron “in orbit” around the nucleus. A proton can be transformed into a neutron if it absorbs an electron, whose negative charge neutralises the proton’s positive one. Conversely, a neutron can transform itself into a proton by expelling a unit of negative charge – one electron. Such transformations are examples of nuclear reactions. Chemical reactions leave the nucleus intact. Nuclear reactions change it …..
"Each element has a characteristic number of protons in its atomic nucleus, and the same number of electrons in “orbit” around the nucleus: one for hydrogen, 2 for helium, 6 for carbon, 11 for sodium, 26 for iron, 82 for lead, and 92 for uranium. This is the element's so-called atomic number which (acting via the electrons) largely determines an electron’s chemical properties. Whilst the neutrons have little effect on an element’s chemical properties, they do affect its mass and its nuclear reactions”. [10.1] |
The mass number appears at the top and the atomic number at the bottom of a chemical element’s symbol in the Periodic Table: thus, using chlorine as an example:
This is the full symbol for a chlorine atom which has 17 protons and also 17 electrons, because the number of protons and electrons in an atom is the same. It also tells us that the total number of protons and neutrons in the chlorine atom is 35, and the number of neutrons can also be deduced from the mass number and atomic number: thus 35 – 17 = 18 neutrons.
Again: protons and electrons both have the same size of electrical charge, but the proton is positive and the electron negative, and neutrons are neutral, and since, the number of electrons in an atom is equal to the number of protons in its nucleus, this means atoms have no overall electrical charge. In an uncharged atom, the atomic number equals the number of electrons, and all the atoms of a particular element have the same atomic number, that is, the same number of protons [11].
This is the full symbol for a chlorine atom which has 17 protons and also 17 electrons, because the number of protons and electrons in an atom is the same. It also tells us that the total number of protons and neutrons in the chlorine atom is 35, and the number of neutrons can also be deduced from the mass number and atomic number: thus 35 – 17 = 18 neutrons.
Again: protons and electrons both have the same size of electrical charge, but the proton is positive and the electron negative, and neutrons are neutral, and since, the number of electrons in an atom is equal to the number of protons in its nucleus, this means atoms have no overall electrical charge. In an uncharged atom, the atomic number equals the number of electrons, and all the atoms of a particular element have the same atomic number, that is, the same number of protons [11].
On the following pages, we will have a look at the Particle Landscape describing the panorama lurking within the atom and its proton and neutron constituents [12].
[1] “Honouring our own Einstein – relatively”: The Age, July 31, 2005.
[2] Frank Wilczek, “Happy Birthday, Electron”, Scientific American, June 2012, 11.
[3] Source: Adam Hart-Davis (ed), Science – The Definitive Visual Guide, London, 2009, 286-7, 292-3.
[4] "On the Constitution of Atoms and Molecules", Philosophical Magazine and Journal of Science, Series 6, Volume 26, Issue 151, July 1913. Bohr won the Noble Prize in Physics for his insight.
[5] Bill Bryson, A short history of nearly everything, Broadway Books, 2013, pp 141-145. See also Gavin Hesketh, The Particle Zoo - The search for the Fundamental Nature of Reality, Quercus, Hhachett, london, 2016, 7-8.
[6] Robert P. Crease, The Great Equations – Breakthroughs in Science from Pythagoras to Heisenberg, Norton, New York, 2008, 172.
[7] Ibid.
[8] Robert P. Crease, The Great Equations, Ch 7, Celebrity Equation E=mc², 172. This energy loss is considered in more detail below under “E=mc²”.
[9] Associate Professor Michael Box's WEA course, What are atoms made of? Session 1, Discovering the Divisible Atom: the Birth of Modern physics, 24 October 2016.
[10] The Periodic Table may be found on the page Supernovas and the Periodic Table of Elements. An excellent interactive model may be found at http://excelhero.com/periodic-table/
[10.1] Richard Dawkins and Yan Wong, The Ancestor's tale - A Pilgrimage to the Dawn of Time, Weidenfeld and Nicolson, London, 2005, 2nd edition, 2016, 593-594.
[11] http://www.bbc.co.uk/schools/gcsebitesize/science/aqa/fundamentals/atomsrev3.shtml
[12] Source for graphic: http://www.scholarpedia.org/article/Nuclear_Forces