1. A few snippets of the past and continuing history of the Periodic Table

Not all scientists are mentioned who perhaps should be, but I've tried to pick out a few 'highlights' and added some footnotes on what was happening in terms of the development of the detailed knowledge of the structure of atoms, so essential to the modern interpretation of the Periodic Table.

It is a good 'advanced' example of how science works i.e. the relationship between experimental data and theories to account for it, questions posed, questions answered, leading to more comprehensive and accurate theories developing.

the result is that the periodic table has extended from 92 naturally occurring elements to 118 due to the development of nuclear physics e.g. nuclear reactors and particle accelerators.

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1.1 The early classification of Antoine Lavoisier of 1789

• The understanding that an element as a unique atomic 'building block' which could not be split into simpler substances or compound is a chemical combination of two or more elements were not at all understood at the time of Lavoisier in the late 18th century.

• 'light' and 'caloric' (heat), were considered 'substances' and the last 'scientific' vestige of the elements of 'earth, fire, air and water' which had there conceptual origin in the Greek civilisation of 2300-2800 years ago.

• However, Lavoisier was correct on a few things e.g. recognition of the elements sulphur, phosphorus and carbon and correctly described their oxides as acidic e.g. they dissolved in water turned litmus turns red.

• Many metallic elements, were correctly identified though I doubt if they were very pure!

• What he described as the 'earthy elements' are of course compounds, a chemical combination of a metal plus oxygen or sulfur e.g. barium oxide (BaO) and barium sulfide (BaS).

• Lavoisier didn't have very high temperature smelting technology, or electrolysis to extract a reactive metal from its salts.

  • Note that electrolysis only developed in science from about 1806 onwards.

• So there barriers to 'separate' the elements in some way e.g. he couldn't extract a reactive metal!

• In other words, at this time, the wrong 'classification' was due to a lack of chemical technology as much as lack of knowledge.

  • A parallel atomic structure history note:

    • You can see from the 1789 'table' Lavoisier and his contemporaries did not have the experiment techniques, data or theoretical framework to clearly distinguish between 'elements' and 'compounds'.

    • It was only in 1808 Dalton proposed his atomic theory based on experimental data and produced the first list of 'atomic weights' (which we now call relative atomic masses, H = 1, O = 16 etc.).

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 1.2 The 1829 work of Johann Döbereiner

• Döbereiner was amongst the first scientists to recognise the 'group' idea of chemically very similar elements.

• Three groups he 'recognised' were (i) Group 1 Alkali Metals, (ii) Group 2 Alkaline Earth Metals, (iii) Group 7 Halogens.

  • A parallel atomic structure history note:

    • The physical and chemical likeness of the three members of these 'triad groups' should be evident and it was based purely on observation, however Döbereiner and contemporaries where unaware of the atomic and molecular nature of these elements e.g. the atomic nature of the metals (M atoms) and the molecular nature of the Halogens (X₂ diatomic molecules). In fact the concept of a 'molecule' was first realised by Avogadro in 1811 but it took 50 years before the genius of his experimental work and intuition was fully realised.

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 1.3 The work of John Newlands 1864

• Newlands recognised that every 7 elements, the 8th seemed to be very similar to the 1st of the previous 7 when laid out in a 'periodic' manner and he was one of the first scientists to derive a 'Periodic Table' from the available knowledge.

  • His system was based on the 8th element in the sequence because the Noble Gases were unknown, today for periods 2 and 3 it is the 9th element in a sequence where you observe repeated chemical and physical properties.

• e.g. his 'table' consists of almost completely genuine elements (Di was a mix of two elements), classified roughly into groups of similar elements and a real recognition of 'periodicity'

• He also recognised that the 'groups' had more than 3 elements (not just 'triads'), and was correct to mix up metals and non-metals in same group e.g. in 5th column there is carbon, silicon, tin (Sn) from what we know call Group 4. However, indium is in group 3 but Ti, Zr have a valency of 4, like Group 4 elements and do form part of vertical column in down what we know call the Transition Metal series

• Other correct 'patterns' if  not precise are recognisable in terms of the modern Periodic Table e.g. half of column 2 is Group 1, half of column 3 is Group 2, half of column 5 is Group 4, half of column 6 is Group 5, half of column 7 is Group 6. If we put his column 1 as column 7, it would seem a better match of today!

• Although none of his vertical column groups match completely, the basic pattern of the modern periodic table  was emerging.

  • However column's 1 and 7 do seem particularly mixed up compared to the modern periodic table.

  • The progressive work of John Newlands was initially rejected because not all the elements fitted the pattern and you may notice that many metals and non-metals are rather mixed up..

  • However, he was very much on the right track and deserves more credit than he is often given because he was a pioneer in the idea of setting out the elements to give vertical columns of 'like elements', which we now call 'groups', and you see this in the contents of most of the columns.

• A parallel atomic structure history note: A 'recap' and a good wedge of history at this point!

  • The Greeks Leucippus and Democritus ~500-400 BC wondered what was the result of continually dividing a substance i.e. what was the end product or smallest bit i.e. what was left that was indivisible - the word atom/atomic is from Greek adjective atomos meaning 'not divisible'.

    • They considered that matter is made of atoms that are too small to be see and cannot be divided into smaller particles. They speculated that there was empty space between solid atoms and that atoms were the same throughout a cross section and atoms could have different sizes, shapes and masses.

    • These were brilliant ideas for their time and such concepts were the result of excellent intuitive thinking BUT the famous and much more eminent and revered philosopher Aristotle, didn't think much of their theory, and so atomic theory never developed for nearly 2000 years!

      Its worth commenting further on the Greeks. Although brilliant in intellectual discourse on many subjects and legendary mathematicians, they were NOT very good at science. Most Greek intellectuals did not consider doing experiments to test out theories as very important, and therefore over 2000 years ago they actually rejected the principal methods by which we today practice science and it took another 2000 years to really get stuck into lots of good scientific experiments!

  • However, the Greeks idea of atoms was not completely forgotten and later revived by Boyle and Newton but with little progress.

    • Robert Boyle (1627-1691) in his book 'The Sceptical Chemist' talked about tiny identical particles that were indivisible but could be joined in various ways to make 'compounds'.

  • But, in 1808, Dalton (1766-1844) proposed his atomic theory that all matter was made up of substances of some kind of 'atomic nature' and the different types of atoms (elements) combined together to give all the different substances of the physical world.

    • His theory included the idea that atoms in an element are all the same and an element was not divisible into more fundamental substances.

    • In 1808 there was no actual proof that individual atom particles existed but Dalton envisaged an element as a fundamental type of substance that could not be split into simpler substances.

    • Dalton considered that a compound is made by joining at least two different elements together to form a new substance in specific proportions (we now write as a formula, and atoms do not change themselves in a reaction but from the original reactants they re-arrange to form the products.

      • This is real progress! Most of his ideas were correct except the 'indivisibility' of the atom! but it would take nearly another 60 years before the idea of 'atomic structure' would take shape from experimental results showing the nucleus was made of protons and neutrons and surrounded by shells of electrons.

    • He also produced the first list of 'atomic weights' (we now call relative atomic masses) on a scale based on hydrogen which was given the arbitrary value of 1 since it was lightest element known, and, as it happens, correctly so.

    • Dalton also devised symbols for the different elements, but his 'picturesque' symbols were not universally adopted and today's elements letter symbols were introduced and promoted by the Swedish chemist Jons Jacob Berzelius in 1811.

  • In 1876 Goldstein and Jean Perrin in 1895 passed a high-voltage electrical discharge through various gases and discovered beams of negatively charge particles where formed.

    • They where called cathode rays and, where in fact, what we now know as negative electrons (but they didn't know this!).

    • The electrons were emanating from the negative cathode electrode and being accelerated towards the positive anode.

    • They were unaware that positive ions were also produced and beamed in the opposite direction.

    • Up till then, it was just assumed that matter consisted of Daltons 'atoms' i.e. particles that could not be broken down into smaller particles, so did not have any meaningful structure but just combined in various ways to make different compounds.

    • This was the real start of research into 'atomic structure', especially as it was soon found later on that a stream of positive particles was travelling in the opposite direction to the 'negative electrons'!

    • Goldstein's and Perrin's experiments also provided the experimental basis for the development of the mass spectrograph by Aston - what we know now as a mass spectrometer.

Background developments in identifying metallic elements

Before proceeding further it is pertinent to consider the history of metal extraction, since most elements in the periodic table are metals and the more elements known, the more the structure of the periodic table can emerge.

The ease of extraction and ultimately being identified as an element is intimately connected to how easy it is to extract a metal. A short summary, based on the reactivity series of metals and methods of extracting metals from ores is outlined below.

• Order of decreasing reactivity related to the earliest know date of extraction and use:

(BP means 'before present time ~2000)

• francium (1939, very radioactive), caesium (1860), rubidium (1861) all from electrolysis

• potassium (1807, 1855 from electrolysis), sodium (1807, from electrolysis)

• lithium (1817, electrolysis?), calcium (1808, from electrolysis)

• magnesium (1755, 1808 from electrolysis), aluminium (1825, by electrolysis)

• zinc (before 1500, ), iron (extracted with charcoal before 3000 BP)

• tin (~4500 BP, used to make bronze)

• lead (over 9000 BP, archaeologist have found lead beads 9000 years old, later used by the Romans for plumbing well over 2000 years ago)

• copper (~11000 BP extracted via charcoal from ores >4000 years ago, found 'native' and was beaten out of rocks and into a useful shape!)

• silver (~7000 BP, used by ancient civilisations)

• gold (~8000 BP, used by ancient civilisations, e.g. Egyptian civilisations, found 'native' in streams and extracted by 'panning')

• platinum (~1735, rare metal but known to ancient South American civilisations before Europeans arrived in the 15th century, brought to Europe ~1750)

  • Detailed notes on the 'Reactivity Series of Metals'

• The date is quoted as the 'normal' year (BCE/AD) or BP meaning years before present year (I've not used BC/BCE).

• The understanding of electricity and the development of d.c. electrical supplies in the early 19th century (1807 onwards, well before Mendeleev) e.g. using simple voltaic batteries meant that the more reactive metals could then be extracted by electrolysis.

• Once more reactive metals could be produced in larger quantities by electrolysis, these metals themselves were then used to extract other metals e.g. chromium which were often difficult to extract by conventional smelting furnaces using carbon.

• So, by the time we reach Mendeleev's periodic table and the work of others like Lothar Meyer (from the 1866s to the 1890s) , quite a large number of elements were well known and characterised, so the time was ripe for further development.

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1.4 Dmitri Mendeleev's Periodic Table and Lothar Meyer's Graphs of 1869

• Mendeleev (Russian chemist) first published his 'Periodic Table' work simultaneously in 1869 with the work of Lothar Meyer (German chemist) who looked at the physical properties of all known elements.

• Lothar Meyer noted 'periodic' trend patterns e.g. peaks and troughs when melting point, boiling points, specific heat and atomic volume (volume of atomic mass in g, = modern volume of 1 mole ) values were plotted against 'atomic weight' - what we now call relative atomic mass.

• My modern versions of Lothar Meyer's graphs are shown on a separate pages, plus others and now the properties are plotted against atomic/proton number and I've managed to collect most data up to element 96.

  • Elements Z = 1 to 20 covering Periods 1-3 and start of Period 4 Elements Z = 1 to 38 covering Periods 1-4 and start of Period 5 Elements Z = 1 to 96 covering Periods 1-6 and start of Period 7

  • The atomic volume graph is shown below clearly showing the 'periodic' highest volumes for the alkali metals - the least dense of the elements in liquid or solid form.

My modern version of Lothar-Meyer's 'atomic volume' curve

and below one of Mendeleev's early versions of the Periodic Table

• Mendeleev laid out all the known elements in order of 'atomic weight' (what we know call relative atomic mass, A_r) except for several examples like tellurium (Te, A_r = 127.60) and iodine (I, A_r = 126.90) whose order he reversed because chemically they seemed to be in the wrong vertical column! Smart thinking!

  • Argon (Ar, A_r = 39.95) and potassium (K, A_r = 39.10) is the 2nd example, but that was not a problem for chemists at the time, because the Group 0 Noble Gases hadn't been discovered by then!

  • These 'anomalies' in the order of 'atomic weights' are explained by the existence of isotopes which were discovered ~1916 and the neutron finally characterised in 1932.

  • Isotopes of elements are atoms of the same proton number with different numbers of neutrons, hence atoms of the same element with different mass numbers.

  • The most abundant stable isotope of potassium is ³⁹K, and that of argon is ⁴⁰Ar, hence the anomaly.

  • Naturally occurring iodine is 100% ¹²⁷I, but tellurium has a range of isotopic masses from ¹²⁰Te to ¹³⁰Te but more the heavier isotopes are more abundant than the lighter isotopes.

• By 1869, Mendeleev and Lothar Meyer had an advantage over Newlands (1864) because by then there was an increased number of known elements and hence 'groups' of similar elements were becoming more clearly defined.

• Mendeleev used a double column approach which is NOT incorrect, i.e. a sort of group xA and xB classification.

  • This is due to the 'insert' of transition metals, some of whom show chemical similarities to the vertical 'groups'.

This is how Mendeleev's periodic table looked in an early Russian publication (in Russian). The left image doesn't look quite as familiar, BUT, if you rotate it round 90o it begins to look much more familiar! All 'familiar 7 vertical groups (1-7, also now numbered 1-2 and 13-17) show up, remember Noble Gases had not been discovered yet. I've added comments that partially explain why Mendeleev got some of the groupings wrong in terms of our modern groups 1 to 7 of the periodic table. Note that despite it being in Russian from the late 19th century, most of the chemical symbols should be familiar to you! that's the idea - a universal language! Group 1 is correct bar Tl and radioactive francium was unknown. Thallium is (Tl) is in group 3 but does have a valency of 1. Group 2 is partly correct, but two wrong, Zn and Cd, but the latter two d block elements have a valency of 2 just like Be and Mg Group 3 B and Al 'correct' but included an unknown and U & Y, the latter two (but have a valency of 3). Group 4 three correct and one unknown predicted (Ge) and lead (Pb) in the wrong place, but the principal valency of lead is 2 so it was included with the group 2 metals Ca, Sr and Ba (Mg is missing?). Group 5 is all correct, quite remarkable since you go down the group from non-metals to a metal. Group 6 is all correct, again quite remarkable grouping, only the unknown radioactive element polonium is missing. Group 7 is all correct, brilliant again,  but couldn't have known about radioactive astatine at that time. Although the complications due to the transition metal series and lanthanide and actinide series of metals due to the electronic sub-groups we now recognise as d blocks or f blocks, Mendeleev still recognised some as 'blocks of ' metals with some similarity. So, no wonder he is given great historical credit for his insight and foresight into the development of the Periodic Table.

• Mendeleev's 'presentation' was sufficiently accurate to predict missing elements and their properties * e.g. germanium (Ge) below silicon (Si) and above tin (Sn) in Group IV and Mendeleev is rightly called the 'father of the modern Periodic Table'.

Parallel atomic structure history notes:

• In 1897 Wien and J J Thompson measured the charge mass ratio of the 'particles' of the cathode rays (electrons) and also showed that the smallest positively charged particle was obtained from hydrogen gas. This 'smallest particle' we now know is the proton.

• J J Thompson ~1897 proposed his 'plum pudding' theory based on the growing evidence that atoms where themselves composed of even small more fundamental particles and the mass and charge of the proton and electron.

  • Thompson envisaged a plumb pudding atom consisting of a positively charged 'pudding' with just enough lighter negatively charged electrons embedded in it to produce a neutral atom.

  • The positive balancing the negative idea was correct but the relative size and nature of the nucleus were not.

• Between 1910-1914 Millikan established the value of the electric charge on an electron in his famous 'oil drop' experiments, hence the mass of the electron could be calculated.

• From 1902-1910 Rutherford, Geiger and Marsden and others used alpha particle scattering experiments - atomic structure notes) to establish the concept of the nucleus.

  • They were even able to make an estimate of the value of its positive charge (which we now know equals the atomic/proton number).

  • Even at that stage it was recognised that this positive nuclear charge bore some relationship to the order of the elements, as given by 'atomic weights', which Mendeleev and others had used to construct their periodic table.

• Experimentally the 'atomic number' of an element was established by Chadwick in 1920 from beta particle scattering experiments - an atom's electrons deflect the bombarding beta particle electrons - from the scattering pattern you can deduce the number of 'atomic' electrons causing the scattering.

• Earlier, the X-ray spectra results of Moseley in 1913 showed that when atoms were bombarded with cathode rays (electrons) X-rays where produced.

  • It was found that the square root of the highest energy emission line (called the K alpha line, K_α) gave a linear plot with the apparent atomic number - the number order in the periodic table - the proton number was not confirmed yet.

  • However the plot of √K_α against atomic weight (relative atomic mass) gave a zig-zag plot.

  • Therefore finally establishing that the really important 'chemical identity number' of an element was the charge on the nucleus, i.e. what we know as the atomic/proton number and this would be the crucial number for ordering the elements, ultimately into the modern periodic table.

• However, there was still the problem of why the atomic mass and atomic number where different i.e. in the case of the lighter elements, the atomic weight was often about twice the atomic number.

• In 1919 Aston developed a cathode ray tube i.e. like those used by Wien and Thompson etc. into a 'mass spectrograph', which we now know as a mass spectrometer notes.

  • This showed that atoms of the same element had different masses but there was no experimental evidence that they had different atomic numbers (which of course they didn't). In 1920 Rutherford suggested there might be a 'missing' neutral particle and in 1932 Chadwick discovered the neutron by bombarding beryllium atoms with alpha particles which produced a beam of neutrons

  • ⁹₄Be + ⁴₂He ==> ¹²₆C + ¹₀n

    • Incidentally, the neutrons are unaffected by electrical and magnetic fields and not directly 'observed', they were primarily detected because they produced a beam of protons on collision with molecules of a hydrocarbon wax by a sort of snooker ball collision effect.

    • The protons are readily detected and characterised (mass 1, charge +1 and their formation linked to the presence of a neutral particle of the same mass (neutron mass 1, charge 0).

• Once the nature of the neutron was finally deduced by Chadwick, it completely explained the nature of isotopes and backed up the ideas from Moseley's work that the fundamentally important number that characterises an element is its atomic number and NOT the atomic mass.

• So, we are now ready to construct the full modern periodic table based on the order of atomic number and wide range of data (formulae, spectroscopy, chemical reactions etc.) on most of the elements up to 92 and now beyond to element 118.

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1.5 A modern version Periodic Table based on the electronic structure of atoms

The electronic basis of the periodic table is explained in Part 2.

• With increasing knowledge of the elements of the Periodic Table it is now laid out in order of atomic (proton) number.

• Due to isotopic masses, the relative atomic mass does go 'up/down' occasionally (there is no obvious 'nuclear' rule that accounts for this, at least at GCSE/GCE level!). BUT chemically Te is like S and Se etc. and I is like Cl and Br etc. and so are placed in their correct 'chemically similar family' group and this is now backed up by modern knowledge of the electron structure of atoms.

• We now know the electronic structure of elements and can understand sub-levels and the 'rules' in electron structure e.g.

  • 2 in shell 1 (period 1, 2 elements H to He),

  • 8 in shell 2 (period 2, 8 elements Li to Ne),

  • there is a sub-level which allows an extra 10 elements (the transition metals) in period 4 (18 elements, K to Kr). this also explains the sorting out of Mendeleev's A and B double columns in a group.

  • The periods are complete now that we know about Noble Gases.

• The use and function of the Periodic Table will never cease! Newly 'man-made' elements are being synthesised. 

• In the 1940's Glenn Seaborg was part of a research team developing the materials required to produce the first atomic bombs dropped on Hiroshima and Nagasaki.

  • He specialised in separating all the substances made in the first nuclear reactors and helped discover the series of 'nuclear synthesised' elements beyond the naturally occurring limit of uranium (₉₂U).

  • From element 93 to 118 are now known, so the structure of the bottom part of the periodic table will continue to grow.

  • There is plenty of scope for present day, and future Mendeleev's!!!! (will you be one of them!?).

• Parallel atomic structure history note:

  • From 1913 onwards the electron structure of atoms was gradually being understood and paralleling the developing knowledge of the structure of the nucleus and its importance in determining which element an atom was i.e. the atomic/proton number.

  • The Bohr theory of the hydrogen spectrum (see section 2.6)  postulated that the electrons surrounding the positive nucleus could only exist in specific energy levels and that any electron level change must involve a specific input/output of energy - the quanta e.g. a photon of light or X-rays etc.

  • In the 1920's and 1930's scientist-mathematicians like Heisenberg and Schrödinger were developing the mathematical equations known as wave mechanics. These mathematical theories describe the detailed behaviour of electrons, and out of these equations come the four quantum numbers from which are derived the set of rules we use to assign electrons in their respective levels (see section 2.2), which ultimately determines the chemistry of an element.

The 'many' names used to indicate the various groups and series of elements in the periodic table

 (I have included all the alternative group numbers - historic and modern)

Alkali metals – The very reactive metals of group 1: Li, Na, K, Rb, Cs, Fr
Alkaline earth metals – The metals of group 2: Be, Mg, Ca, Sr, Ba, Ra
Pnictogens – The elements of group 5/15: N, P, As, Sb, Bi (non-metals ==> metals)
Chalcogens – The elements of group 6/16: O, S, Se, Te, Po, Lv (non-metals ==> metals)
Halogens – The elements of group 7/17: F, Cl, Br, I, At
Noble gases – The elements of group 0/8/18: He, Ne, Ar, Kr, Xe, Rn
Lanthanoids – Elements 57–71: La, Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu
Actinoids – Elements 89–103: Ac, Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr
Rare earth elements – Sc, Y, and the lanthanoids
Transition metals – Elements in groups 3 to 11 or 12. (eg 3d block Sc to Cu or Zn)

Other miscellaneous 'names' comments, not standard IUPAC descriptors
Lanthanoids and actinoids may be referred to as lanthanides and actinides respectively.
Post-transition metals – metals of groups 13–16: Al, Ga, In, Tl, Sn, Pb, Bi, Po.
Metalloids – elements with properties intermediate between metals and non-metals: B, Si, Ge, As, Sb, Te, At.
Diatomic nonmetals – nonmetals that exist as diatomic molecules in their standard states: H, N, O, F, Cl, Br, I.
Superactinides – hypothetical series of elements 121 to 155, which includes a predicted "g-block" of the periodic table.
Precious metal – non-radioactive metals of high economical value eg silver, gold, platinum
Coinage metals – various metals used to mint coins eg the coinage metals Ni, Cu, Ag, and Au.
Platinum group – Ru, Rh, Pd, Os, Ir, Pt
Noble metal – vague term for corrosion resistant metals like silver and gold and the platinum-group metals
Heavy metals – metals like lead, on the basis of their density, atomic number, or toxicity
Native metals – metals that can occur pure in nature eg gold and copper
Transuranium elements – elements with atomic numbers greater than 92 (U)
Transactinide elements – elements after the actinides with atomic numbers greater than 103 (Lr)
Transplutonium elements – elements with atomic number greater than 94 (Pu)

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1.6 Where did elements come from originally? Where do WE get the elements from?

(a) Where did elements come from originally? It all starts in the STARS!

• The ultimate origin of all elements is the nuclear reactions that go on when stars are formed from inter-stellar dust and gas forming a huge combined mass due to gravity, and then 'chunks' of a star cool down to form planets.

• The heaviest elements are formed in nuclear fusion reactions when stars self-destruct in super-nova explosions.

  • The nuclear synthesis of light elements up to Z = 26 (Fe, iron) occurs in stars formed from the condensation of hydrogen and helium atoms.

  • Eventually, as the mass increases, the force of gravity causes such compression that the temperatures rise considerably at high matter densities and nuclear reactions spontaneously begin.

  • Up to Z = 26 nuclei, they are usually formed energy releasing fusion processes or the decay of unstable nuclei

  • There are hundreds of possible nuclear transformations possible, so, below, I've chosen some examples of possible nuclear reactions, whose products fit in with the isotopes, mass numbers, relative atomic masses etc. which A level chemistry students are likely to come across ...

  • ... in the nuclear equations, for the nuclide symbol ^A_ZX, A = mass number, Z = atomic number, X = element symbol

    • and gamma photons, and neutrinos, often formed in nuclear changes, are NOT shown for simplicity.

  • ¹₁H + ¹₁H ==> ²₁H + ¹₀n

    • The formation of hydrogen-2 (deuterium) occurs slowly at a temperature of 1.5 x 10⁷ K e.g. in our Sun (~15 million ^oC).

  • ²₁H + ¹₁H ==> ³₂He + Ɣ

    • This is one way the next heaviest element, helium can be formed.

  • ³₂He + ³₂He ==> ⁴₂He + 2¹₁H

    • From helium-3, the formation of helium-4, the most common isotope of helium we find on earth.

    • From helium-4, by what is known as the alpha process, a succession of heavier elements can be synthesised in subsequent nuclear reactions ...

  • 2⁴₂He ==> ⁸₄Be

    • beryllium-8 is very unstable, but is readily converted on collision with another helium-4 nucleus to give stable carbon-12 (and ultimately life on earth!)

  • ⁸₄Be + ⁴₂He ==> ¹²₆C

    • These sorts of nuclear reactions need star temperatures of 1 x 10⁸ K (~100 million ^oC)

  • ¹²₆C + ⁴₂He ==> ¹⁶₈O

    • From carbon-12 you can get oxygen-16, the most common oxygen isotope on earth today.

  • ¹⁶₈O + ⁴₂He ==> ²⁰₁₀Ne

    • From oxygen-16 you can get neon-20, the most common neon isotope on earth today.

  • ²⁰₁₀Ne + ⁴₂He ==> ²⁴₁₂Mg

    • From neon-20 you can get magnesium-24, the most common magnesium isotope on earth today.

  • ²⁴₁₀Mg + ⁴₂He ==> ²⁸₁₄Si

    • From magnesium-24 you can get silicon-28, the most common silicon isotope on earth today.

  • ²⁸₁₄Si + ⁴₂He ==> ³²₁₆S

    • From silicon-28 you can get sulfur-32, the most common sulphur isotope on earth today.

  • ³²₁₆S + ⁴₂He ==> ³⁶₁₈Ar

    • From sulphur-32 you can get argon-36, the most common argon isotope on earth today

  • ³⁶₁₈Ar + ⁴₂He ==> ⁴⁰₂₀Ca

    • From argon-36 you can get calcium-40, the most common calcium isotope on earth today.

• You can see from the Periodic Table of relative atomic masses how the alpha-process ('helium burning' has produced the values for C, O, Ne, Mg, Si, S, Ar and Ca from the principal isotope of multiples of four mass units.

• There are lots of other possibilities involving H and He nuclei and particularly complicated nuclear fusion cycle involving carbon nuclei e.g. the six step cycle ...

  • ¹²₆C + ¹₁H ==> ¹³₇N

  • ¹³₇N ==> ¹³_cC + ⁰₊e

  • ¹³₆C + ¹₁H ==> ¹⁴₇N

  • ¹⁴₇N + ¹₁H ==> ¹⁵₈O

  • ¹⁵₈O ==> ¹⁵₇N + ⁰₊e⁺  (decay of oxygen-15 by positron emission)

  • ¹⁷₇N + ¹₁H ==> ¹²₆C + ⁴₂He

  • You can also see how other isotopes of an element can be formed and in the cycle carbon-12 is reformed to continue these particular nucleosynthesis pathways.

  • There is a good analogy here with auto-catalytic cycles in chemistry.

• The heavier elements beyond iron i.e. Z > 26 Co cobalt onwards must be formed by energy absorbing processes including neutron capture e.g. the formation of technetium from molybdenum.

• The conditions necessary are only found in supernova explosions because extremely high temperature are needed to create particles of sufficient kinetic energy to overcome the massive repulsion forces between nuclei of a multiple positive charge.

  • Apart from fusion reactions like those described for 'our sun', heavier elements can also be formed by neutron absorption and subsequent radioactive decay e.g.

  • ⁹⁸₄₂Mo + ¹₀n ==> ⁹⁹₄₂Mo

    • neutron absorbed by molybdenum-98 nucleus to give an unstable Mo nucleus

  • ⁹⁹₄₂Mo ==> ⁹⁹₄₃Tc + ⁰_-e

    • Mo-98 nucleus decays by beta particle emission to give an atom of technetium, with a higher atomic number.

  • Similarly, gallium can be formed from zinc, i.e. again forming an element of higher atomic number ...

  • ⁶⁸₃₀Zn + ¹₀n ==> ⁶⁹₃₀Zn    followed by   ⁶⁹₃₀Zn ==> ⁶⁹₃₁Ga + ⁰_-e

• So you can see that these nuclear fusion, neutron or proton capture, nuclear decay etc., can over time, gradually produce all the heavier elements up to element 92 uranium, the last of our naturally occurring elements.

• Even though small amounts of ²³⁸₉₂U are eventually formed, it requires the highest of temperature e.g. in a super-nova explosion of giant stars a lot bigger than our sun!

• Some more examples of nuclear fusion to form heavier elements are quoted on Where do heavier elements come from?

• All the elements from atomic numbers Z = 1 to 92 (H to U) naturally occur on Earth, though some are very unstable-radioactive and decay to form more nuclear stable elements.

  • Many isotopes of elements after lead, ₈₂Pb are unstable.

  • After uranium, ₉₂U, the vast majority of the isotopes of the elements of atomic number >92 are inherently unstable.

  • They will not have survived even if they were formed billions of years ago in the Sun, and retained or formed in the initial 'spin-off' material that formed the 'very early' Earth.

  • However, the advent of nuclear reactors has enabled up to kg quantities of e.g. plutonium, ₉₄Pu (used in nuclear reactors and weapons) and americium, ₉₅Am (used in smoke alarms) to be produced.

  • Cyclotrons, particle bombardment linear accelerators, have enabled 'super-heavy' elements up to Z = 118? to be 'synthesised', but only a few atoms at a time (The Russia-US space race seems to have been partly replaced by 'who can synthesize the biggest atom'!).

  • One things for certain, the Periodic Table still keeps growing with newly synthesised elements!

(b) Where, and how, do we get the elements from the earth?

• Everything around you is made up of the elements of the periodic table, BUT most are chemically combined with other elements in the form of many naturally occurring compounds e.g.

  • hydrogen and oxygen in water, sodium and chlorine in sodium chloride ('common salt'), iron, oxygen and carbon as iron carbonate, carbon and oxygen as carbon dioxide etc. etc.!

• Therefore, most elements can only be obtained by some kind of chemical process to separate or extract an element from a compound e.g.

  • Less reactive metals are obtained by reduction of their oxides with carbon and more reactive metals are extracted by electrolysis of their chlorides or oxides (see GCSE/IGCSE/AS notes on Metal Extraction)

  • Non-metals are obtained by a variety of means e.g. chlorine is obtained by electrolysis of sodium chloride solution (see GCSE/IGCSE notes on Group 7 The Halogens).

• However some elements never occur as compounds or they occur in their elemental form as well as in compounds e.g.

  • The Group 0/8/18 Noble Gases are so unreactive they are only present in the atmosphere as individual atoms. Since air is a mixture, these gases are separated from air by a physical method of separation by distillation of liquified air.

  • The elements oxygen and nitrogen are obtained from air at the same time, which is far more convenient than trying to get them from compounds like oxides and nitrates etc.

  • Gold/platinum is are the least reactive metals and are usually found 'native' as the yellow/silver elemental metal.

  • Relatively unreactive metals like copper and silver can also be found in their elemental form in mineral deposits as well as in metal ores containing compounds like copper carbonate, copper sulphide and silver sulphide.

  • The non-metal sulphur is found combined with oxygen and a metal in compounds known as sulfides, but it can occur as relatively pure sulphur in yellow mineral beds of the element.

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1.7 APPENDIX 1. ALL THE KNOWN ELEMENTS

Elements from Z = 1 to 118 in alphabetical order,

so, given the atomic number, you can find it on the full modern periodic table above (section 1.5)

Chemical Symbol Element name Atomic No. Z Ac Actinium 89 Al Aluminium/Aluminum 13 Am Americium 95 Sb Antimony 51 Ar Argon 18 As Arsenic 33 At Astatine 85 Ba Barium 56 Bk Berkelium 97 Be Beryllium 4 Bi Bismuth 83 Bh Bohrium 107 B Boron 5 Br Bromine 35 Cd Cadmium 48 Cs Caesium/Cesium 55 Ca Calcium 20 Cf Californium 98 C Carbon 6 Ce Cerium 58 Cl Chlorine 17 Cr Chromium 24 Co Cobalt 27 Cn Copernicium 112 Cu Copper 29 Cm Curium 96 Ds Darmstadtium 110 Db Dubnium 105 Dy Dysprosium 66 Es Einsteinium 99 Er Erbium 68 Eu Europium 63 Fm Fermium 100 Fl Flerovium 114 F Fluorine 9 Fr Francium 87 Gd Gadolinium 64 Ga Gallium 31 Ge Germanium 32 Au Gold 79 Hf Hafnium 72 Hs Hassium 108 He Helium 2 Ho Holmium 67 H Hydrogen 1 In Indium 49 I Iodine 53 Ir Iridium 77 Fe Iron 26 Kr Krypton 36 La Lanthanum 57 Lw Lawrencium 103 Pb Lead 82 Li Lithium 3 Lv Livermorium 116 Lu Lutetium 71 Mg Magnesium 12 Mn Manganese 25 Mt Meitnerium 109 Md Mendelevium 101 Hg Mercury 80

Chemical Symbol Element name Atomic No. Z Mo Molybdenum 42 Mc Moscovium 115 Nd Neodymium 60 Ne Neon 10 Np Neptunium 93 Ni Nickel 28 Nh Nihonium 113 Nb Niobium 41 N Nitrogen 7 No Nobelium 102 Og Oganesson 118 Os Osmium 76 O Oxygen 8 Pd Palladium 46 P Phosphorus 15 Pt Platinum 78 Pu Plutonium 94 Po Polonium 84 K Potassium 19 Pr Praseodymium 59 Pm Promethium 61 Pa Protactinium 91 Ra Radium 88 Rn Radon 86 Re Rhenium 75 Rh Rhodium 45 Rg Roentgenium 111 Rb Rubidium 37 Ru Ruthenium 44 Rf Rutherfordium 104 Sm Samarium 62 Sc Scandium 21 Sg Seaborgium 106 Se Selenium 34 Si Silicon 14 Ag Silver 47 Na Sodium 11 Sr Strontium 38 S Sulphur/Sulfur 16 Ta Tantalum 73 Tc Technetium 43 Te Tellurium 52 Ts Tennessine 117 Tb Terbium 65 Tl Thallium 81 Th Thorium 90 Tm Thulium 69 Sn Tin 50 Ti Titanium 22 W Tungsten 74 U Uranium 92 V Vanadium 23 Xe Xenon 54 Yb Ytterbium 70 Y Yttrium 39 Zn Zinc 30 Zr Zirconium 40

No elements synthesised or named beyond Z = 118 so far!

WHAT NEXT? PLEASE NOTE GCSE Level periodic table notes are on separate webpages Advanced Level Inorganic Chemistry Periodic Table Index: Part 1 Periodic Table history Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr AND important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots All 11 Parts have their own sub-indexes near the top of the pages Group numbering and the modern periodic table The original group numbers of the periodic table ran from group 1 alkali metals to group 0 noble gases (= group 8). To account for the d block elements and their 'vertical' similarities, in the modern periodic table, groups 3 to group 0/8 are numbered 13 to 18. So, the p block elements are referred to as groups 13 to group 18 at a higher academic level, though the group 3 to 0/8 notation is still used, but usually at a lower academic level. The 3d block elements (Sc to Zn) are now considered the head (top) elements of groups 3 to 12.

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