Bond is the word we use when we speak of what holds things together. It is useful to organize the universe into the different types of bonds which hold it together.

Bonds among masses

A force of gravity equal to Gm₁m₂/d² exists between any two masses (”m₁“, “m₂“) in the universe when those masses are separated by a distance “d”. This force is, at least compared with other forces we will consider, very weak, but if the masses are large enough, like with planets, it can hold them together, as in a solar system, and it can hold you, pretty much, on the earth. There’s no point in jumping. You’ll just fall back down.

Bonds among charges

Coulomb’s Law expresses the force (equal to kq₁q₂/r²) which exists between two charged particles (”q₁“,”q₂“) separated by a distance, “r”. This force is referred to as an electrostatic force. Electrostatic forces are roughly 10²⁵ times as strong as gravitational forces. So atoms and molecules rarely jump up off the surface of the earth as people do. Most of nature is locked up pretty tightly by these forces. Atoms bond with each other to form molecules because atoms are composed of charged particles, protons and electrons, which attract each other. These bonds, within the molecule, are intramolecular. The nature of the distribution of these charges determines how molecules will bond with each other to form liquids and solids. These bonds, between molecules, are intermolecular. Since bonds within molecules are usually stronger than those between molecules, we will consider those first. After all, when energy is applied, most liquids will boil before they decompose.

Intramolecular Bonds

1. Ionic Bond: The pulling power of an atom for electrons from an outside source, called the electronegativity, increases diagonally across the periodic table, from a low at Francium (excluding the noble gases) to a high at Fluorine. If the two atoms involved in bonding have very different electronegativities, for example, Na and Cl, one atom, Cl, will pull an electron from Na, forming two ions, Na⁺¹ and Cl^-1. The two ions will attract each other and form an ionic bond. Ionic bonds are bonds between ions and are always formed as the result of a gain and loss of electrons. But it is the electrostatic attraction between the two ions which is responsible for the bond.

2. Covalent Bond: When two atoms, both with equally high electronegativies, for example, two chlorine atoms, approach each other, unpaired electrons with opposed spin are swept between the positively charged nuclei, as described by the Valence Bond Model of Covalent Bonding. This electron pair, shared evenly between the atoms, pulls the positive nuclei together to form a covalent bond. Although both atoms are electrically neutral, it is still the distribution of proton and electron charges which provide the Coulombic attraction between atoms. Covalent bonds are sub-classified into

(a) single bonds: in which a single pair of electrons, one from each atom, are shared, as in Cl₂.

(b) double bonds: in which two pairs of electrons are trapped between the nuclei, as in O₂.

(c) triple bonds: in which three pairs of electrons are trapped between the two nuclei, as in N₂.

(d) coordinate covalent bonds in which one atom supplies both of the electrons for the bond as in SO₂.

3. Polar Covalent Bond: Often, two atoms with relatively high electronegativities will form covalent bonds, but because the electronegativities are different, the electron pair(s) is not shared equally. This creates a side of the molecule which is electron rich (with a partial negative charge) and another side which is electron poor (with a partial positive charge). This is a covalent bond with some ionic character and is called polar covalent, as is found, for example, in H₂O. The hydrogen atoms (electronegativity is 2.1) are electron poor and the oxygen atom (electronegativity is 3.5) is electron rich. A dipole is a pair of separated, opposite charges. The greater the dipole (measured as a dipole moment) the greater the ionic character of the covalent bond.

4. Metallic Bond: Atoms are held together in a solid metal or metals (alloy) by a metallic bond. Since both atoms have low electronegativities and low ionization potentials, the nuclei of atoms are considered to be positively charged ions occupying lattice positions. Those electrons in the atom which have low ionization potentials are considered to be free to move through the lattice, giving metals the ability to conduct heat and electricity. However the bond between two atoms in the metal is still electrostatic, between the positive nuclei and the loosely held negative electrons.

Intermolecular Bonds

1. Ionic Bonds: Once an ionic bond is formed, as in Na⁺¹ + Cl^-1, the resulting molecule is, of course, highly polar. In NaCl, an electron is said to have been completely transferred from Na to Cl creating a full dipole. Dipoles attract each other very strongly and so NaCl molecules readily attach to each other to form giant three dimensional crystals. The high melting and boiling points of ionic crystals attest to the great strength of these ionic bonds. The brittleness and hardness of these compounds is explained by their rigid interlocking structure. In the solid state, the electrons are locked tightly in the crystal so that solid ionic compounds are extremely poor conductors of electricity. However, when these rigid bonds are loosened by melting, the ions are free to move throughout the compound and conduct electricity well.

2. Dipole-Dipole bonds (Partially Ionic) When two BrF molecules approach each other, the partial negative charge of the F of one molecule will attract the partial positive charge of the Br of the second molecule. This creates a kind of an ionic bond, but a rather weak one, called a dipole-dipole bond. Such molecules have low melting and boiling points because these intermolecular bonds require little energy to rupture. However particularly strong dipole-dipole bonds are found in the water molecule, which has a large dipole moment. This type of bond is often given its own name, a hydrogen bond, however it is still a dipole-dipole bond. When H’s electron is pulled partially away, the resulting nucleus has no other electrons to shield its positive charge and small radius, resulting in an especially strong interaction of dipoles. Naturally, water is already a liquid at room temperature, and is easily solidified, a sign of the strength of those hydrogen bonds.

3. Metallic Bonds: As in ionic compounds, the nature of intermolecular bonding in metals is just more of the same type found in intramolecular bonds. In fact, some consider a metallic substance to be one giant molecule with countless nuclei imbedded in a sea of electrons. Proton-electron attraction in metals decreases as the size of the atom decreases, since the outer “free” electrons are unable to get as close to the positively charged nuclei. Thus, large metallic atoms tend to have weaker bonds, and lower melting points.

4. Covalent Bonds When zillions of carbon atoms come together and form covalent bonds using sp³ hybridized orbitals, the dense, interlocking tetrahedral structure which results in called a network compound. It is characterized by a seemingly unending array of carbon atoms which we call diamond. The strength of these bonds may be appreciated when one tried to break a diamond.

5. Induced-Dipole Induced-Dipole Bonding (London Dispersion Forces) Non-polar molecules are still made up of charged particles which do affect each other. If at any one moment, an excess of electrons appear on one side, it will create a temporary dipole which can induce a dipole in a neighboring molecule. At extremely low temperatures, these dipoles can flicker back and forth and serve to hold the neutral molecules together. He₂ may be formed at very low temperatures and its formation is described with this type of bond. Clearly, these bonds are much weaker than others previously described.

Bonds among quarks

As strong as electrostatic forces are, at least compared with gravitational forces, the strong force within the atomic nucleus is a thousand times greater! This type of bond, within the atomic nucleus, is equally strong between proton-proton, proton-neutron, and neutron-neutron. The atomic nucleus owes its incredible stability to the strength of this force. The bond operates only over extremely tiny distances and involves matter-energy transformation (E=mc²).

Bonds among fermions

A different kind of force, called the weak force is used to describe, for example, bonding between the proton and electron in the neutron with the simultaneous formation of a pion (p + e^- ? n + v). This weak force interaction always involves the interaction among four different fermions.

One is inclined to feel just a bit nervous about one’s understanding of the bonds among quarks and fermions. This may be because familiarity with the concepts of mass and charge from the age of 3 years old has made us forget their essential mystery.

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