University:
University of Alberta
Course:
CHEM 101 | Introductory University Chemistry I
Academic year:

2024
Views:

417

Pages:

4
Author:

Neena N.

Metallic bonding Metals make up the left-hand side (s-block) and middle (d-block) of the periodic table. They contain low numbers of valence (outer shell) electrons, low ionisation energies and low electronegativity resulting in all metals losing electrons easily to be overall positively charged. Metallic bonding is the electrostatic attraction of positively charged cations and delocalised electrons. It is the attraction between the positively charged cations and the negatively charged delocalised electrons that forms the bond This structure is described as a regular, three-dimensional lattice of positive ions, that occupy fixed positions in the lattice. This is a metallic lattice. The metallic bonding lattice is simple and therefore results in limitations of this model. The limitations include: 1. The different bond strengths of metals can't be shown. 2. Properties such as lustre and magnetism and differences in electrical conductivity can't be shown. 3. They don't show the movement of the sea of delocalised electrons. The electrons come from the valence shells (outer shells). For example, a sodium atom (Na = 11) has an electronic configuration of 1s22s22p63s¹. This means it has one valence electron, as can be seen in the 3s subshell. This valence electron becomes delocalised with the other sodium atoms to give the overall delocalised sea of electrons. The sodium atom will form a sodium ion with a +1 charge, as it now has one more proton in its nucleus than electrons in the shells. It will have 11 protons and 10 electrons. A metal that has more valence electrons would produce a more negative sea of delocalised electrons and a greater positive charge on the cation, leading to a greater attraction. This leads to a stronger metallic bond and so contributes to higher melting and boiling points of metals. An aluminium atom (Al = 13) has an electronic configuration of 1s22s22p63s 3 This means it has three valence electrons, as can be seen in the 3s subshell. The valence electrons become delocalised with the other aluminium atoms to give the overall delocalised sea of electrons Transition metals have more valence electrons in the 3d and 4s subshells, and so have greater attractions and higher melting and boiling points The ability of metals to conduct electricity comes from the sea of delocalised electrons. These delocalised electrons move freely whilst carrying charges which gives the metal the conductive properties. One of the key factors that differentiates metallic bonding from other types of bonding is that each metal atom is bonded to a set number of other atoms. This forms a constant rigid structural pattern. For example, magnesium is always surrounded by 12 other magnesium atoms and sodium is always surrounded directly by 8 other sodium atoms. Bond strength is also influenced by the radius of the atom. It is important that you understand that these metallic bonds do not involve metal cations. Metallic bonding that occurs involves metal atoms does not involve any cations. Whilst the electron(s) are not attached directly to their original atom, they are delocalised and shared between the other atoms and so have not been lost Therefore, when describing metals in their structures, it is written in atomic form e.g. Na not Na+.