Diagonal relationship

Pictorial representation of examples of diagonal relationship.

A diagonal relationship is said to exist between certain pairs of diagonally adjacent elements in the second and third periods of the periodic table. These pairs (lithium (Li) and magnesium (Mg), beryllium (Be) and aluminium (Al), boron (B) and silicon (Si) etc.) exhibit similar properties; for example, boron and silicon are both semiconductors, forming halides that are hydrolysed in water and have acidic oxides.

The organization of elements on the periodic table in to horizontal rows and vertical columns makes certain relationships more apparent (periodic law). Moving rightward and descending the periodic table have opposite effects on atomic radii of isolated atoms. Moving rightward across groups decreases the atomic radii of atoms, while moving down the periods will increase the atomic radii [1].

Similarly, on moving rightward a group, the elements become progressively more covalent[clarification needed], less basic and more electronegative, whereas on moving down a period the elements become more ionic, more basic and less electronegative. Thus, on both descending a period and crossing a group by one element, the changes “cancel” each other out, and elements with similar properties which have similar chemistry are often found – the atomic size[clarification needed], electronegativity, properties of compounds (and so forth) of the diagonal members are similar.

It is found that the chemistry of a first-row (second period) element often has similarities to the chemistry of the second-row (third period) element being one column to the right of it in the periodic table. Thus, the chemistry of Li has similarities to that of Mg, the chemistry of Be has similarities to that of Al, and the chemistry of B has similarities to that of Si. These are called diagonal relationships. (It is not as noticeable after B and Si.)
The reasons for the existence of diagonal relationships are not fully understood, but charge density is a factor. For example, Li+ is a small cation with a +1 charge and Mg2+ is somewhat larger with a +2 charge, so the charge density on each ion is roughly the same.
Using the Li–Mg pair:(under room temperature and pressure)

  • Li and Mg form only normal oxides whereas Na forms peroxide and metals below Na, in addition, form superoxides.
  • Li is the only Group 1 element which forms a stable nitride, Li3N.[2] Mg, as well as other Group 2 elements, also form nitrides.[2]
  • Lithium carbonate, phosphate and fluoride are sparingly soluble in water. The corresponding Group 2 salts are insoluble. (Think lattice and solvation energies).
  • Both Li and Mg form covalent organometallic compounds. LiMe and MgMe2 (cf. Grignard reagents) are both valuable synthetic reagents. The other Group 1 and Group 2 analogues are ionic and extremely reactive (and hence difficult to manipulate).[3]
  • Chlorides of both Li and Mg are deliquescent (absorb moisture from surroundings) and soluble in alcohol and pyridine. Lithium chloride, like magnesium chloride (MgCl2.6H2O) separates out from hydrated crystal LiCl.2H2O.
  • References[edit]

  • ^ See Ebbing and Gammon General Chemistry 10th ed.: Atomic Radius p. 318-319
  • ^ a b Clark, Jim (2005). “Reactions of the Group 2 Elements with Air or Oxygen”. chemguide. Retrieved January 30, 2012. 
  • ^ See Shriver and Atkins Inorganic Chemistry 4th ed.: Li/Mg p. 259; Be/Al p. 274; B/Si p. 288.