Summary: Types of Forces & Bonds
A bond is simply a force of attraction between two or more things. The simplest forces, called fundamental forces, lead to the various types of bonds and all the forces we encounter in the world.
Bonds among Nucleons - STRONG FORCE
This type of bond, within the atomic nucleus, is equally strong in binding proton to proton, proton to neutron, and neutron to neutron. Though it is 100 times stronger than the electromagnetic force it has very short range. It only becomes noticeable at small distances of separation. Only when protons and or neutrons become very close together (like in a nucleus) will this force bind them. The bond involves matter-energy transformation (E=mc2).
Bonds among charges - ELECTROMAGNETIC FORCE
Objects with an electrostatic charge attract or repel each other. This force is referred to as an electrostatic or electromagnetic force. Electrostatic forces are roughly 1010 times stronger than the "weak" force. The matter we interact with every day is controlled and in fact defined by these forces. When these bonds hold the atoms within a single molecule together they are called intramolecular bonds. When these bonds hold molecules to each other they are called intermolecular bonds. The bonds within a molecule are usually stronger than those between molecules, generally, when energy is applied, water will boil before it will separate into hydrogen atoms and oxygen atoms.
Intramolecular Bonds (hold atoms within a molecule)
A. Ionic Bond: The electrostatic attraction between two ions of opposite charge. A common example is the formation of NaCl. One atom of Cl, will pull an electron from an atom of Na, forming two ions, Na+1 and Cl-1. The 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. They form when atoms with very large differences in electronegativities (attraction for an electron) get close to each other.
B. Covalent Bond: When two atoms, both with equally high electronegativies, for example, two chlorine atoms, approach each other, unpaired electrons are swept between the positively charged nuclei.. 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 charges which provide the electrostatic attraction between atoms. Covalent bonds are sub-classified into
- single bonds: in which a single pair of electrons are shared, as in Cl2.
- double bonds: in which two pairs of electrons are trapped between the nuclei, as in O2.
- triple bonds: in which three pairs of electrons are trapped between the two nuclei, as in N2.
- coordinate covalent: bonds in which one atom supplies both of the electrons for the bond as in SO2.
C. Polar Covalent Bond: Two or more 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 (more negative) and another side which is electron poor (more positive). This is a covalent bond with some ionic character and is called polar covalent, as is found, for example, in H2O. The hydrogen atoms having a lower electronegativity (2.1) are more electron poor (+) and the oxygen atom (electronegativity of 3.5) is more electron rich(-).
D. Metallic Bond: Atoms are held together in a solid metal or metal alloy by metallic bonds. Since both atoms have low electronegativities and low ionization potentials, the nuclei of atoms are considered to be positively charged ions occupying evenly spaced (lattice) positions. The valence electrons in the atom which have low ionization potentials are considered to be a free moving "sea" within 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
A. 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 completely transferred from Na to Cl creating a dipole (- at one end + at the other). Dipoles attract each other very strongly so NaCl molecules attach to each other to form giant three dimensional crystals. The high melting and boiling points of ionic crystals indicate the great strength of the ionic bonds. The brittleness and hardness of these compounds is explained by their rigid interlocking structure. When ionic compounds are solid, the electrons are locked in the crystal so that they are extremely poor conductors of electricity. However, when these ionic bonds are loosened by melting, the ions are free to move throughout the compound and conduct electricity well.
B. Dipole-Dipole bonds (Partially Ionic): These bonds form when the intramolecular bonds are polar covalent (uneven sharing). When two BrF molecules approach each other, the partial negative charge near the F of one molecule will attract the partial positive charge near the Br of the second molecule. This creates a weak ionic bond called a dipole-dipole bond. Because these intermolecular bonds require little energy to break, these molecules have low melting and boiling points. Unusually strong dipole-dipole bonds are found in the water molecule, which has a large apparent charge on the H atoms. When H's electron is pulled more toward the oxygen, the H 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. Even though this type of bond is often given its own name, a hydrogen bond, it is still a dipole-dipole bond.
C. Metallic Bonds: The nature of intermolecular bonding in metals is just more of the same kind found in intramolecular metallic bonds. A metallic substance can be thought of as one giant molecule with many nuclei embedded in a sea of electrons. Larger metallic atoms (like gold) tend to have weaker bonds, and lower melting points.
D. Covalent Bonds: Just as in the metallic bonds, the division between inter and intra becomes very fuzzy when atoms form a network compound which is an array of atoms with no defined limit to the number of atoms. There is no easily identified molecule that makes this up, it is just atom upon atom. When carbon atoms come together and form covalent bonds the dense, interlocking tetrahedral structure which results is called diamond. The strength of these bonds is realized when one tries to break a diamond.
E. Induced-Dipole Induced-Dipole Bonding (London Dispersion Forces) Non-polar molecules and atoms of noble gasses are made up of charged particles which do affect each other even though there is no portion of the molecule with more charge than the other. If at any one moment, an excess of electrons happen to be on one side, it will create a temporary dipole which can cause a dipole in a neighboring molecule. At extremely low temperatures, these dipoles can flicker back and forth and hold the neutral molecules together. If the energy of motion (temperature) is low enough on H2 the induced dipole bonding may be strong enough to bond H2 to H2 and form a solid. Keep in mind that ordinarily there should be no attraction between two hydrogen molecules. The same force causes noble gasses to bond with themselves. (He + He forms He2 at low tempertures) These bonds are much weaker than any other molecular bonds.
Bonds among leptons - WEAK FORCE
The "Weak" force is approximately 1030 times stronger than the gravitational force. This force binds the proton to an electron in forming a neutron. This force in action can result in electron capture or in failure results in B-decay.
Bonds among masses - GRAVITATIONAL FORCE
The attraction between any two objects with mass is called the gravitational force. This force is the weakest of all the forces and virtually unnoticeable between small masses. Only when dealing with planet like masses does the force become realistically noticeable. Given enough mass, like the Earth and Sun, it can act at great distances and is responsible for holding entire solar systems together as well as holding us to the earth.
