There are three main kinds of bonds: electrovalent, covalent and metallic. Each and every bond kind gives characteristic properties to the compounds which are made up of.
As electrovalent bonding is between metals and non-metals, covalent bonding is between non-metals. In the creation of electrovalent and covalent bonds valence electrons play very significant roles and each and every valence shell of the bonded atoms accomplish inert gas stable configuration. For metal-metal bond, the valence electrons are very few that the electrons sharing to accomplish electron octet is not possible. Electrovalent bonds can't be made as metals tend to lose electrons and not accept them. A negatively charged metal ion is not possible.
How then do we elucidate bonding in the metallic solids? How can we describe the fact that metallic solids are good conductors of electricity and heat? What are the main differences in the structures of ionic and metallic solids? The questions above will be answered here. We shall as well describe the origin of intermolecular forces which hold covalent molecules altogether in the mass sample and account for their special properties.
Metals generally have 1, 2 or 3 electrons in their valence shells.
Li (2, 1) or 1s2 2s1
Na (2, 8, 1) or 1s2 2s2 2p6 3s1
Mg (2, 8, 2) or 1s2 2s2 2p6 3s2
Al (2, 8, 3) or 1s2 2s2 2p6 3s2 3p1
In the metallic bonding, each and every metal atom contributes its valence electrons to form a cloud or sea of delocalized electrons such electrons don't belong to any specific metal atom however will circulate freely via the metal lattice. The electrostatic attractions between the positive cores which form the metal lattice and the sea or cloud of electrons comprise the metallic bond. The above description of metallic bonding means that the lattice makes a single large crystal. This accounts for the high strength of metals. There is no direction to metallic bond and therefore the metallic lattice can be distorted simply by drawing and hammering. Metals are ductile and malleable. The free moving electrons conduct electricity and heat by their movement. The strength of metallic bond based on the attraction of the electron cloud to the positive cores in the metal lattice. The metallic bond strength rises with the number of valence electrons each metal contribute to the electron cloud. Now, take the illustration of:
Mg (2, 8, 2) 1s2 2s2 2p6 3s2
Na (2, 8, 1) 1s2 2s2 2p6 3s1
Sodium is a softer metal than magnesium as for sodium only one valence electron per atom however for magnesium two electrons are contributed per atom to the electron cloud.
Obeying the above argument compare the strength of the metallic bonding in magnesium by that in aluminium (13).
For metals in the similar group of the periodic table, metallic strength reduces down the group. The raise in atomic size down the group is not accomplished by any increase in electron cloud strength. These listed properties of metals are illustrated by the metallic bonding just described.
a) High tensile strength.
b) Conductor of heat and electricity.
c) Malleable and ductile.
d) High density.
e) Solid at room temperature except mercury that is a liquid at room temperature.
f) Usually have shiny surface (lustre)
Uses of metals:
The utilization of a metal for a specific application based on a number of factors that comprises cost, durability, availability and performance. The table presents the uses of certain metals and the reasons for such uses.
Table: Uses of metals
Reason for use
Excellent conductor of electricity and very ductile
Coating tin cans
The ionic and covalent bonds stand for very strong interactions between the atoms in a compound. In addition to such bonds there are other weaker attractive forces which exist between the atoms and molecules. The existence of such weak attractive forces describes a number of physical properties of certain compounds. Because such forces are generally between molecules they are termed as intermolecular forces. For illustration: Van der Waal's forces, dipole-dipole attractions and hydrogen bonding.
Van der Waal's forces:
Van der Waal's forces exist even between the uncombined atoms and non-polar molecules. A non polar molecule is one in which the electron pair for bonding is equivalently shared by the atoms comprised in the bond formation.
Illustrations of non-polar molecules are N2, Cl2, H2, O2 and so on, that is, covalent bond between the two identical atoms is a non polar bond. Non polar bond might as well exist between unlike atoms if they contain the same electronegativity, for illustration CO2.
The movement of electrons around an atom can lead to the momentary shift of more electrons to one side of the molecule than the other. All through this shift an imbalance in charge exists by one side of the molecule slightly positive and the other slightly negative. The positive end will attract the negative end of the other molecule close to it. This attraction constitutes a bond. This attractive force might be strong however because it is for a short time its effect is usually very small.
The magnitude of this force rises with increasing number of electrons. This force is present among all molecules atoms and ions. Its effect can be extremely large if there are lots of electrons in the molecules or atoms. Let's take the case of halogens (that is, Group VII elements) fluorine, chlorine are gases, bromine is a liquid whereas iodine is a solid. Keep in mind, that all of them exist as diatomic molecules and are only bonded altogether by van der waal forces; Van der Waal's forces are attractions between molecules that take place due to the formation of temporary dipoles in all the molecules. The huge number of electrons in bromine and Iodine permits for substantial cohesive force between bromine and iodine molecules making bromine liquid and iodine solid at room temperature. Van der Waal's forces are at times termed as induced dipole-induced dipole attraction.
The Covalent bonding between atoms of various elements will outcome in a polar bond. The shared electron pair will be more in the control of the more electronegative atom. Let's take the illustration of HCI. Chlorine is more electronegative than hydrogen. The shared pair of electron is controlled more via Chlorine. The chlorine end of molecule will be slightly negative and the hydrogen end slightly positive example: HCl, Hδ+, - Clδ-. The positive end of one hydrogen chloride molecule will attract the negative end of the other molecule. This is dipole-dipole attraction. Via dipole-dipole interactions are not as substantial as full ion-ion interactions; they are stronger than the Van der Waal's forces.
The table below provides the normal boiling points of covalent compounds. Dipole interactions are just around one percent as strong as covalent and ionic bonds.
Table: Boiling points of covalent compounds
Boiling point (K)
The given are significant observations from the table.
1) The boiling points of non-polar compounds are extremely as low compared to polar compounds.
2) The boiling point rises by molecular mass for polar and non-polar compounds.
Boiling point H2 < Boiling point O2 Boiling point HCl < Boiling point HBr
3) The boiling point of H2O is unusually high compared to those of other hydrides example: HCl, HBr and H2S.
Hydrogen bonding is a special kind of dipole-dipole interaction that takes place when hydrogen is bonded to very small electronegative elements such as N, F and Oxygen. In combination with these small electronegative elements, hydrogen carries a substantial positive charge. The attraction of this positive end with the negative end of the other molecule will comprise a strong bond. This bond is the hydrogen bond. Hydrogen bond is around 5 to 10 times stronger than ordinary dipole-dipole interaction. This is not as strong as ordinary covalent bonds between the atoms in a compound.
Hydrogen bonding is mainly responsible for water being a liquid at room temperature instead of a gas. Hydrogen bonding illustrates the high boiling point of water as compared to hydrogen sulphide. Hydrogen bonding describes why hydrofluoric acid is a weaker acid than the hydrochloric acid.
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