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Doc Brown's Chemistry Clinic

 The Periodic Table 

 Part 1 "Historical Introduction" 

 Advanced Level Inorganic Chemistry Revision Notes

(e.g. UK Advanced Level Chemistry GCE-AS-A2-IB US K12 grades ~11-12)

GCSE Atomic Structure Notes  *  GCSE Periodic Table notes  *  EMAIL query?comment

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 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?

Advanced Periodic Table Index * Part 2 Electronic structure of atoms : Spectroscopy and the H spectrum : Ionisation energies * Part 3 Period 1 survey : 1. Hydrogen : 2. Helium : Summary of  Period 1 : heavier element formation-stellar nuclear fusion * Part 7 s-block metals Groups 1/2 Alkali/Alkaline Earth Metals * Part 10 3d-block Sc-Zn and Transition Metals * Part 11 Group and Series data summaries and links to periodicity plots *

 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

• 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.

 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.

  • 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.

 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.

• 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.

  • Atomic structure history note: In 1876 Goldstein 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 negative electrons (but he didn't know this) emanating from the negative electrode and being accelerated towards the positive anode. He was also 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 did not have any structure but just combined in various ways to make different compounds.

 1.4 Dmitri Mendeleev's Periodic Table of 1869

• It was published simultaneously in 1869 with the work of Lothar Meyer 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 and specific heat values were plotted against 'atomic weight' - what  we now call relative atomic mass. do atomic volume curve

• 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 because 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

• His '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.

  • 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, K_α) gave a linear plot with the apparent atomic number. 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' 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).

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

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

• With are knowledge the modern Periodic Table 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.

    

GENERAL REVISION NOTES

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