Chemical Bond
From MyMCAT
A chemical bond is the physical process responsible for the attractive interactions between atoms and molecules, and that which confers stability to diatomic and polyatomic chemical compounds. In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. Molecules, crystals, and diatomic gases—indeed most of the physical environment around us—are held together by chemical bonds, which dictate the structure of matter.
Bonds vary widely in their strength. Generally covalent bond and ionic bonds are often described as "strong", whereas hydrogen bonds and Van der Waals forces are generally considered to be "weak". Care should be taken because the strongest of the "weak" bonds can be stronger than the weakest of the "strong" bonds.
Overview
Chemical bonds, for the sake of simplicity, are classically assigned characteristics of two major types: covalent and ionic.
In a pure covalent type bond valence electrons are shared; the bond forms as valence electrons become attracted to the region of space between two nuclei. In this region, negatively-charged electrons experience attraction from the positively-charged protons from more than one nucleus. In turn, the nuclei are stabilized in position by the pull from these shared electrons.
In a pure ionic type bond one or more outer electrons donated (not shared as in covalent bonds) from one atom to another. This transfer causes the donating atom to assume a net positive charge, and the other to assume a net negative charge the atoms thus become positive or negatively charged ions. The bond then results from electrostatic attraction between these ionized atoms.
Valence bond theory
(Valence bond theory likely beyond scope of MCAT). Valence bond theory proposes that a chemical bond forms when two valence electrons work or function to hold two nuclei together. The six rules for valence bond theory pertaining to the shared electron bond, are:
- 1. The electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.
- 2. The spins of the electrons have to be opposed.
- 3. Once paired, the two electrons cannot take part in additional bonds.
- 4. The electron-exchange terms for the bond involves only one wave function from each atom.
- 5. The available electrons in the lowest energy level form the strongest bonds.
- 6. Of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie in the direction of the concentrated orbital.
Strong chemical bonds
| Typical bond lengths in pm and bond energy in kJ/mol. Bond lengths can be converted to Ångström|Å by division by 100 (1 Å = 100 pm). Data taken from<ref>Bond Lengths and Energies</ref>. | ||
| Bond | Length (pm) | Energy (kJ/mol) |
|---|---|---|
| H — Hydrogen | ||
| H–H | 74 | 436 |
| H–C | 109 | 413 |
| H–N | 101 | 391 |
| H–O | 96 | 366 |
| H–F | 92 | 568 |
| H–Cl | 127 | 432 |
| H–Br | 141 | 366 |
| C — Carbon | ||
| C–H | 109 | 413 |
| C–C | 154 | 348 |
| C=C | 134 | 614 |
| C≡C | 120 | 839 |
| C–N | 147 | 308 |
| C–O | 143 | 360 |
| C–F | 135 | 488 |
| C–Cl | 177 | 330 |
| C–Br | 194 | 288 |
| C–I | 214 | 216 |
| C–S | 182 | 272 |
| N — Nitrogen | ||
| N–H | 101 | 391 |
| N–C | 147 | 308 |
| N–N | 145 | 170 |
| N≡N | 110 | 945 |
| O — Oxygen | ||
| O–H | 96 | 366 |
| O–C | 143 | 360 |
| O–O | 148 | 145 |
| O=O | 121 | 498 |
| F, Cl, Br, I — Halogens | ||
| F–H | 92 | 568 |
| F–F | 142 | 158 |
| F–C | 135 | 488 |
| Cl–H | 127 | 432 |
| Cl–C | 177 | 330 |
| Cl–Cl | 199 | 243 |
| Br–H | 141 | 366 |
| Br–C | 194 | 288 |
| Br–Br | 228 | 193 |
| I–H | 161 | 298 |
| I–C | 214 | 216 |
| I–I | 267 | 151 |
| S — Sulfur | ||
| C–S | 182 | 272 |
Strong chemical bonds are intramolecular forces, which hold atoms together in molecules. The larger the difference in electronegativity, the more an electron is attracted to a particular atom involved in the bond, and the more "ionic" properties the bond is said to have ("ionic" means the bond electron(s) are unequally shared). The smaller the difference in electronegativity, the more covalent properties (full sharing) the bond has.
Covalent bond
Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. Bonds within most organic compounds are described as covalent.
Polar covalent bond
Polar covalent bonding is intermediate in character between a covalent and an ionic bond.
Ionic bond
Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity difference. There is no precise value that distinguishes ionic from covalent bonding but a difference of electronegativity of over 2.0 is likely to be ionic and a difference of less than 1.5 is likely to be covalent.<ref> Template:Cite book </ref> Ionic bonding leads to separate positive and negative ions. Ionic charges are commonly between −3 to +3.
Bent bonds
Bent bonds, are bonds in strained or sterically hindered molecules whose binding orbitals are forced into a banana-like form. Bent bonds are often more susceptible to reactions than ordinary bonds.
Aromatic bond
In aromatic compounds bonds which occur in planar rings of atoms where the Huckel's rule (4n+2 rule) determines whether ring molecules would show extra stability.
In benzene, the model aromatic compound, 18 bonding electrons bind 6 carbon atoms together to form a planar ring structure. The "bond order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.
Metallic bond
In metallic bonding, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. Because of delocalization or the free moving of electrons, it leads to the metallic properties such as conductivity, ductility and hardness.
Intermolecular bonding
There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point) of a substance.
Permanent dipole to permanent dipole
A large electronegativity difference between two strongly bonded atoms within a molecule causes a dipole to form (a dipole is a pair of permanent partial charges). Dipoles will attract or repel each other.
Hydrogen bond
In some ways this is an especially strong example of a permanent dipole. However, in the hydrogen bond, the hydrogen proton comes closer to being shared between target and donor atoms. Hydrogen bonds explain the relatively high boiling points of liquids like water, ammonia, and hydrogen fluoride, compared with their heavier counterparts in the same periodic table column.
Van der Waals (Instantaneous dipole to induced dipole)
Instantaneous dipole to induced dipole, or van der Waals forces, are the weakest, but also the most prolific—occurring between all chemical substances. Imagine a helium atom: At any one point in time, the electron cloud around the (otherwise neutral) atom can be thought to be slightly imbalanced, with momentarily more negative charge on one side. This is referred to as an instantaneous dipole. This dipole, with its slight charge imbalance, may attract or repel the electrons within a neighbouring helium atom, setting up another dipole. The two atoms will be attracted for an instant, before the charge rebalances and the atoms move on.
