History of Periodic Table

What you need to know

Reflections and Exam tips

The quest for a systematic arrangement of the elements started with the discovery of individual elements.

The Law of Triads

German chemist Johann Dobereiner (1780-1849) grouped elements based on similarities.

Law of Triads - the middle element in the triad had atomic weight that was the average of the other two members.

For example:

Calcium (atomic weight 40), strontium (atomic weight 88), and barium (atomic weight 137) possess similar chemical prepares.  Dobereiner noticed the atomic weight of strontium fell midway between the weights of calcium and barium: Ca (40),  Sr  (88),  Ba (137) =    (40 + 137) ÷ 2 = 88.

The law of triads worked for alkali metal triad (Li/Na/K) and the halogen triad (Cl/Br/I) but couldn't be applied to all other elements.

Law of Octaves

English chemist John Newlands (1837-1898), having arranged the 62 known elements in order of increasing atomic weights, noted that after interval of eight elements similar physical/chemical properties reappeared. 

Law of Octaves - elements exhibit similar behaviour to the eighth element following it in the table.

Newlands was the first to formulate the concept of periodicity in the properties of the chemical elements.

Problem of law of octaves

1. The positions of some pairs of elements are reversed when ordered by mass (K and Ar).
2. Not all elements had been discovered at the time and Newlands left no spaces for undiscovered ones.
3. Some groups contained elements with differing properties.

Mendeleev's Periodic Table

In 1869, Russian chemist Dimitri Mendeleev (1834-1907) proposed arranging elements by atomic weights and properties (Lothar Meyer independently reached similar conclusion but published results after Mendeleev). 

Mendeleev's table exhibited similarities not only in small units such as the triads, but showed similarities in an entire network of vertical, horizontal, and diagonal relationships.

Mendeleev’s ordered the elements by their relative atomic mass.

In order to make similar elements line up in the same group:

1. He swapped the positions of certain pairs of elements (e.g. Ar and K, I and Te).
2. He also had to leave gaps in certain places, e.g. between gallium and arsenic
3. He predicted the properties of the missing elements and was proved correct in each case.

Arranging the elements according to increasing atomic numbers and not atomic masses eliminated some of the inconsistencies associated with Mendeleev's table.

Group 1 (alkali metals) elements

Properties of alkali metals

1. Soft, reactive metals with low density which must be stored under oil
2. Have one electron in the outer shell
3. Exhibit metallic bonding in which the outer electron from each atom is lost into a delocalised ‘sea of electrons’ free to move between a lattice of metal ions
4. Relatively low melting points as the solid lattice is held together by the electrostatic forces between the metal cations (1+) and a single delocalised electron (1-) per atom
5. Good conductors of heat and electricity due to the ‘free’ electrons (lost from the outer shell) which can ‘flow’ between the metal ions in the solid lattice.

Reactions with halogens

Alkali elements react with non-metals to form white, soluble, crystalline ionic compounds which have high melting points due to the strong attraction between the metal cations (1+) and the negatively charged non-metal anions.

Reaction with water

React with water to form hydrogen gas and the metal hydroxide which dissolves in water to give an alkaline solution.

2 Na(s) + 2 H₂O(l) → 2 NaOH(aq) + H₂(g)

Reactivity

Alkali elements are more reactive going down the group as the outer electron (further from the nucleus) is less strongly attracted to the positive nucleus and hence more easily lost.

Group 7 (halogens) elements

Properties of halogens

1. Atoms have seven electrons in the outer shell
2. Share an electron with another atom to complete the outer shell and form diatomic molecules (X₂)
3. Poor conductors of heat and electricity as all electrons are held tightly
4. Low melting and boiling points (due to weak intermolecular forces of attraction which increase down group)
5. Change in state as you go down group - fluorine and chlorine (gases), bromine (liquid) to iodine and astatine (solids)
6. Coloured molecules F₂ (pale green); Cl₂ (yellow); Br₂ (red); I₂ (dark grey, sublimes to give purple vapour).

Reactivity

React with metals (gaining an electron in the outer shell) to form ionic compounds which have high melting points due to the strong attraction between the metal cations (positive) and the halide anions (-1).

React with non-metals (sharing one electron in the outer shell) to form covalent molecular compounds which have low melting points due to the weak intermolecular forces of attraction.

Trend in reactivity

Halogens become less reactive going down the group as the outer electrons (further from the nucleus) are less strongly attracted to the positive nucleus and hence additional electrons are less easily gained.

More reactive halogens will displace non-metals of lower reactivity from their salts

Cl₂(aq) + 2KBr(aq) → Br₂(aq) + 2 KCl(aq)

Transition metals

1. Electrons are being added to the d – orbital 
2. Exhibit metallic bonding in which more than one electron from each atom is lost into a delocalised ‘sea of electrons’ free to move between a lattice of metal ions
3. Have high melting points as the solid lattice is held together by strong electrostatic forces between the metal cations and sea of delocalised electrons
4. Hard, tough and strong because of the large forces holding the lattice together
5. Good conductors of heat and electricity due to the ‘delocalised’ electrons which can ‘flow’ between the metal ions
6. Malleable and ductile as layers of metal ions slide over each other in the lattice without disrupting the attractions which hold the lattice together

Reactivity of transition metals

1. Less reactive than group 1 metals (because their outer electrons are less easily lost), reacting much more slowly with air and water
2. Form ionic compounds with various charges or oxidation states
3. In general, form coloured compounds
4. Produce insoluble hydroxides with characteristic colours when reacted with sodium hydroxide.

Uses of transition metals

1. Used in construction due to their strength and resistance to corrosion
2. Used in the electrical industry due to their high conductivity
3. Used in the pigment industry (paints e.t.c) due to their coloured compounds
4. Used as catalysts due to their variable oxidation state