Part 1. A Historical introduction to the periodic table and the nuclear physics origin of elements

Doc Brown's Chemistry  

Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Some historical aspects of the development of the concept of the Periodic Table from the work of chemists like Mendeleev are presented, alongside comments about discoveries occurring at the same time as physicists were investigating atomic structure. The modern periodic table is then presented, though the electronic justification for its structure is presented in Part 2. Finally, a brief summary of some examples of how the different elements are formed by the nuclear processes in stars to give our naturally occurring elements. Finally, brief comments with links to more detailed notes about how we extract elements to exploit for our own use.

For non-A level students (c) doc b KS4 Science GCSE/IGCSE Periodic Table notes - including simplified historical comments

INORGANIC Part 1 Historical Introduction page sub-index: 1.1 The early classification of Antoine Lavoisier of 1789 * 1.2  The 1829 work of Johann Döbereiner * 1.3 The work of John Newlands 1864 * 1.4 Dmitri Mendeleev's Periodic Table and Lothar-Meyer graphs of ~1869 * 1.5 A modern Periodic Table based on the electronic structure of atoms * 1.6 Where did the elements come from originally and where do we get the elements from today?

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

 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 'highlight' 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. Its 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.

 1.1 The early classification of Antoine Lavoisier of 1789

Antoine Lavoisier's 1789 classification of substances into four 'element' groups

acid-making elements gas-like elements metallic elements earthy elements
sulphur light cobalt, mercury, tin lime (calcium oxide)
phosphorus caloric (heat) copper, nickel, iron,  magnesia (magnesium oxide)
charcoal (carbon) oxygen gold, lead, silver, zinc barytes (barium sulphate)
  azote (nitrogen) manganese, tungsten argilla (aluminium oxide)
  hydrogen platina (platinum) silex (silicon dioxide)
  • The understanding that an element as a unique atomic 'building block' which could not be split into simpler substances and compound is a chemical combination of two or more elements were not at all understood at the time of Lavoisier.
  • '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. the elements sulphur, phosphorus and carbon and correctly described their oxides as acidic e.g. dissolved in water turned litmus turns red.

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

  • What he described as the 'earthy elements' are of course compounds, a chemical combination of a metal plus oxygen or sulfur (both O and S in case of barium).

  • He didn't have very high temperature smelting technology, or a reactive metal from electrolysis (came in about 1806 onwards)' 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.

    • 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.

Advanced Inorganic Chemistry Page Index and Links


 1.2 The 1829 work of Johann Döbereiner

  • Johann Döbereiner noted that certain elements seemed to occur as 'triads' of similar elements e.g.
    • (i) lithium, sodium and potassium
    • (ii) calcium, strontium and barium
    • (iii) chlorine, bromine and iodine
  • 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.

    • 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.

Advanced Inorganic Chemistry Page Index and Links


 1.3 The work of John Newlands 1864

Newland's 'Law of Octaves' (his 'Periodic Table' of 1864)

H Li Ga B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
Co, Ni Cu Zn Y In As Se
Br Rb Sr Ce, La Zr Di, Mo Ro, Ru
Pd Ag Cd U Sn Sb Te
I Cs Ba, V Ta W Nb Au
Pt, Ir Tl Pb Th Hg Bi Cs
  • 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.

  • 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 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 but 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.

  • Atomic structure history note: 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!

    • 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 tiny hard particles called atoms 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.

      • Dalton considered that a compound is made by joining at least two different elements to form a compound,

      • 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.

      • 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 chemist Jons Berzelius in 1811.

    • In 1876 Goldstein and Jean Perrin in 1895 (1870-1942) 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 he didn't know this!).

      • The electrons were emanating from the negative 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.

Advanced Inorganic Chemistry Page Index and Links


1.4 Dmitri Mendeleev's Periodic Table, 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 or boiling points, specific heat and atomic volume 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, Ar) except for several examples like tellurium (Te, Ar = 127.60) and iodine (I, Ar = 126.90) whose order he reversed because chemically they seemed to be in the wrong vertical column! Smart thinking!

    • Argon (Ar, Ar = 39.95) and potassium (K, Ar = 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'. We now recognise theses dual columns as

  • Advanced Inorganic Chemistry Page Index and LinksHis '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'.

    • 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.

    • 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 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 (GCSE-AS atomic structure notes) to establish the concept of the nucleus and 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 were using to construct their periodic table.

    • Experimentally the 'atomic number' of an element was established by Chadwick in 1920 from beta particle scattering experiments (an atoms electrons deflecting the bombarding beta particle electrons) and from the X-ray spectra results of Moseley in 1913. Moseley 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, Kα) gave a linear plot with the apparent atomic number. However the plot of √Kα against atomic weight (relative atomic mass) gave a zig-zag plot. Therefore finally establishing that the really important 'chemical identity number' 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 GCSE-AS atomic structure 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 the hydrocarbon wax by a sort of snooker ball collision effect. The protons are readily detected and characterised (mass 1, charge +10 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.

Advanced Inorganic Chemistry Page Index and Links


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.

Pd s block 3d to 6d blocks of Transition Metals (Periods 4 to 7), note that the 1st (d1) and 10th (d10) are NOT true transition elements. p block
Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Gp7/17 Gp0/18

1H   Note: (i) H does not readily fit into any group, (ii) He not strictly a 'p' element but does belong in Gp 0/18

2 3Li 4Be The full IUPAC modern Periodic Table of Elements (ZSymbol, z = atomic or proton number) 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba *57-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb 83Bi 84Po 85At 86Rn
7 87Fr 88Ra *89-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Uut 114Fl 115Uup 116Lv 117Uus 118Uuo
Gp 1 Alkali Metals

Gp 2 Alkaline Earth Metals

Gp 7/17 Halogens

Gp 0/18 Noble Gases

Take note of the four points on the right

*57La 58Ce 59Pr 60Nd 61Pm 62Sm 63Eu 64Gd 65Tb 66Dy 67Ho 68Er 69Tm 70Yb 71Lu
*89Ac 90Th 91Pa 92U 93Np 94Pu 95Am 96Cm 97Bk 98Cf 99Es 100Fm 101Md 102No 103Lr

*Horizontal insert in Period 6 of the Lanthanide Metal Series (Lanthanides/Lanthanoids) Z=57 to 71 including the 4f-block series. *Horizontal insert in Period 7 of the Actinide Series of Metals (Actinides/Actinoids) Z=89-103 including the 5f-block series.

  1. Using 0 to denote the Group number of the Noble Gases is historic i.e. when its valency was considered zero since no compounds were known. However, from 1961 stable compounds of xenon have been synthesised exhibiting up to the maximum possible valency of 8 e.g. in XeO4.

  2. Because of the horizontal series of elements e.g. like the Sc to Zn block (10 elements), Groups 3 to 7 and 0 can also be numbered as Groups 13 to 18 to fit in with the maximum number of vertical columns of elements in periods 4 and 5 (18 elements per period 4 and period 5).

  3. This means that 21Sc to 30Zn can be now considered as the top elements in the vertical Groups 3 to 12.

  4. Advanced Inorganic Chemistry Page Index and LinksI'm afraid this can make things confusing, but there it is, classification is still in progress and the notation Group 1 to 18 seems due to become universal.

  5. Elements up to Z = 118 have now been synthesised, if only a few atoms have been identified !

  • 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 113 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 em'?).

    • 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.

      Advanced Inorganic Chemistry Page Index and Links

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

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 0/18: 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)


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 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 AZX, 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

      • 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 ==> 13cC + 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 e.g. the removal of ozone by chlorine atoms.

    • 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

      • 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 examples of nuclear fusion reactions to form heavier elements are quoted in Part 3.4 Where do heavier elements come from?

  • All the elements from atomic numbers 1-92 (H-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 93+ 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.

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

    • The Group 0 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 sulphates, but it can occur as relatively pure sulphur in yellow mineral beds of the element.

  • -

Advanced Inorganic Chemistry Page Index and Links


INORGANIC Part 1 Historical Introduction page sub-index: 1.1 The early classification of Antoine Lavoisier of 1789 * 1.2  The 1829 work of Johann Döbereiner * 1.3 The work of John Newlands 1864 * 1.4 Dmitri Mendeleev's Periodic Table and Lothar-Meyer graphs of ~1869 * 1.5 A modern Periodic Table based on the electronic structure of atoms * 1.6 Where did the elements come from originally and where do we get the elements from today?

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

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