Introduction to Chemical Bonds

Chemical Bonds Introduction to Chemical Bonds A chemical bond is the force that holds atoms together in compounds Chemical bonds can be classified into different types: metallic bonds (in pure metals), covalent bonds (in molecular substances), and ionic bonds (in ionic substances) Molecular substances —-> Covalent bonds —> lonic substances —->Ionic bonds Bond Energy When a chemical bond forms, the system releases energy called bond energy. This same amount of energy is required to break the bond Bond strength relationship: High bond energy = Strong bond Low bond energy = Weak bond Covalent Bond The covalent bond is a bond in which two atoms, belonging to non-metals, share a pair of electrons. The formation of a bond requires a distance that balances the repulsion between the charges of the same sign of the two atoms and the attraction between the charges of opposite sign between the atoms themselves. Let's imagine that two people (the atoms) must share a rope. If the two people got too close to each other they might collide and move too far apart they could not share the rope. The distance must be balanced. This distance of "equilibrium" is called BIND DISTANCE and represents the distance between the nuclei of two atoms of a chemical species. The electrons not participating in the formation of chemical bonds are called "LONE PAIRS" or "SOLITAR DOUBLES". Between an atom A and a B atom can be formed, at most, three covalent bonds (6 shared electrons in total, i.e. 3 pairs). The number of shared pairs is also called "BIND ORDER" and the higher this number, the higher the total energy of the bond will be and the atoms will be spatially close together. But what happens to atomic orbitals when a chemical bond is formed? They overlap and change to form a "MOLECULAR ORBITAL". A molecular orbital is the region of space occupied by a pair of shared bonding electrons and surrounding both bonded nuclei. It has a different shape and energy (smaller) than the molecular orbitals of the individual isolated atoms. The shape of the molecular orbitals depends on the type of orbitals involved in the binding and the orientation in which they overlap. The latter can be head/head type (or frontal), which generates a molecular bond called σ, or side/ side type (or lateral), which forms a π bond. The bond of type σ (sigma) occurs between two atoms that share one electron each and is formed by the overlapping of the outer orbits. The bond π (pi) is a covalent chemical bond formed by lateral overlapping of two orbitals of appropriate symmetry. The single covalent bond is formed by a bond σ, the double covalent bond has a bond σ and one π and, finally, the triple covalent bond has a bond σ and two bonds π. The single covalent bond is formed by a bond σ, the double covalent bond has a bond σ and one π and, finally, the triple covalent bond has a bond σ and two bonds π. Another important characteristic of covalent bonds is that they can form between equal or different atoms. This implies the existence of two types of covalent bonds in this context: - a covalent bond between equal atoms is called "PURE COVALENT or homopolar" (since there is no difference in electronegativity, a bond between different atoms with a difference in electronegativity < 0.4 can also be defined as such). - the covalent bonds between two different atoms (and with a difference of electronegativity > 0,4 and < 1,7) are called "POLAR COVALENT" bonds, because there is a separation of partial charge between the atoms themselves. It is as if, during a tug-of-war, one group pulls harder than the other but not enough to pull the rope out of its hand and then we will have most of the rope concentrated on one side of the system. In the same way we'll have the electrons more oriented towards the more electronegative atom. Dative Bond The dative (or coordination) bond is a particular type of bond where the doublet of electrons is provided by only one donor atom (which then "donates" one of its lone pairs) while the other atom accepts it since it has an empty orbital. An example is the ammonia NH3 which has an lone pair on nitrogen and can give a dative bond to a hydrogen cation (H+) to form the ammonium cation (NH4+). Ionic Bond Understanding Ionic Bonding An ionic bond forms between atoms with a large electronegativity difference (≥ 1.7), resulting in the transfer of electrons and formation of oppositely charged ions. In ionic compounds, ions don't exist as isolated pairs. Each ion interacts with multiple ions of opposite charge through electrostatic forces, while experiencing repulsion from similarly charged ions. This results in the formation of ordered crystalline structures. The lattice energy is the energy released during crystal formation. It increases with higher ionic charges and smaller ionic radii, leading to stronger crystal structures. lonic compounds typically form between metals and non-metals. They exhibit distinctive physical properties: high melting and boiling points, electrical conductivity only in molten or dissolved states, and crystalline nature making them hard but brittle. Metallic Bond All metals have an ordered three- dimensional crystalline structure, characterized by being formed by cations of the metal itself immersed in a sea of mobile valence electrons (the valence electrons are free to move between the cations and no longer belong to individual atoms). We can imagine the metal bond as a series of fixed and neatly arranged spheres (metal cations), immersed in a tank full of moving liquid (electrons). The metal bond is responsible for the properties of the various metals: electrical conductivity: this phenomenon simply translates into an exchange of electrons between the source of electricity and the sea of electrons, generating a flow of electrons that passes the metal itself. The electrical conductivity is influenced by heat because the energy supplied by heating results in a "vibratory movement" of the cations, around their position of equilibrium, and a greater agitation of the electrons of the metal bond. The set of these chaotic motions, in particular those of the cations, generates a difficulty in the crossing of the electrons, supplied from outside, through the metal, causing a lowering of the electrical conductivity; - thermal conductivity: the presence of the sea of electrons favours thermal conduction, which at microscopic level, results in sudden movements of the electrons; - ductility (ability of a body or material to deform plastically under load before breaking) and malleability (ability of a body or material to be easily deformed and reducible in thin laminiform layers):the metal bond allows cations to flow on electrons, without two cations meeting and repelling each other. This system allows the metal to be processed without causing immediate breakage; - lustre (ability to reflect light): the light radiation that collides on the metal excites the electrons of the metal bond which, immediately afterwards, return to their energetic state releasing light radiation in all directions. This phenomenon causes metals to shine.