Extra Bonding Notes

Chemical bonds form to lower the energy of the system, the components of the system become more stable through the formation of bonds. Inert gases (Noble gases) are stable due to a full valence shell of electrons (ns², np⁶). Bonds form to enable elements to attain this 'noble gas configuration'. Ionic bonds involve the transfer of electrons from a metal to a nonmetal, forming positive and negative ions that are attracted to each other. Covalent bonds ~8 formed by overlap (or combination) of atomic orbitals of each element to form a molecular orbital. Metallic bonds can be seen as each metal donating electron(s) to a common "sea" of electrons which are shared by all the ions within the solid.

Hydrogen bonding: The oxygen-hydrogen bond is polar, oxygen being more electronegative element. The molecule is therefore polar (the molecule is not linear but has a bent, V, shape). This is extenuated by the two lone pairs of electrons on the oxygen atoms. One end of the molecule is partially negative while the two hydrogen atoms become partially positive. The molecules of water are attracted to one another, with the slightly positive hydrogens attracted to the negative 'ends' (the oxygen of other water molecules). This intermolecular attraction is termed hydrogen bonding and acts almost like a glue holding the molecules of water together.

If we hook hydrogen up to an atom that is very good at attracting electrons (like N, O, or F), the hydrogen end of the bond becomes very positively charged and the other atom becomes negatively charged (polar). Remember that hydrogen is the smallest atom on the periodic table, it is possible for two molecules to get very close together. The combination of high polarity and close approach result in the interaction being particularly strong.

Coulombic forces: attractive (or repulsive) forces between charged particles. The force of attraction between No opposite charges is proportional to the magnitude of their charges divided by the square of the distance between them.

Hydrogen bonding produces adhesion, cohesion, surface tension, anal capillary action.

In polar molecules, the presence of oppositely charged ends produces forces of attraction between these molecules that are greater than the forces of attraction between similar nonpolar molecules. These forces affect the properties of the polar substances. Some of the effects include an increase in the boiling and melting points and higher heats of vaporization and fusion. These greater intermolecular forces also are likely to account for the lower vapor pressure of substances having polar molecules. Also, polar molecules have a greater attraction for ions.

It is the polarity of its molecules that makes water a good solvent. Many ionic compounds readily dissolve in water. The H‑O bonds in water are polar because O atoms have a higher electronegativity than H atoms.

If the water molecule were linear, it would be nonpolar because the bonds in a linear water molecule would be symmetrically placed. The lack of symmetry in electron distribution gives a partial negative charge to the end of the molecule at the greatest distance from both hydrogen atoms. At the opposite end, there is a partial positive charge.

The hydrogen atom in a hydrogen bond is in effect bonded to two atoms, more weakly to one than the other. It is bonded covalently to an atom within its own molecule and, through the hydrogen bond, to an atom in a neighboring molecule. The strength of the hydrogen bond increases with the degree of electronegativity of the atom bonded to the hydrogen. The strength decreases with an increase in the size of the bonded atom. For example, nitrogen and chlorine atoms have nearly the same electronegativies. However, the hydrogen bond between hydrogen in one molecule and nitrogen in an adjacent one is much stronger than the bond between hydrogen and an adjacent chlorine atom. This is because nitrogen atoms are much smaller than chlorine atoms. The negative charge of the electrons in the nitrogen atom is concentrated into a smaller volume, and it therefore exerts a greater attraction for the proton of the hydrogen atom in a neighboring molecule. In fact, the bond between the hydrogen atom and an adjacent chlorine atom is not considered to be a hydrogen bond because it does not have a sufficiently great force of attraction. In any molecule where a hydrogen atom is bonded to one of the small, highly electronegative atoms F, O, or N, hydrogen bonding is likely to occur.

Hydrogen bonding also explains why some substances have unexpectedly low vapor pressures, high heats of vaporization, and high melting points. In order for vaporization or melting to take place, molecules must be separated. An input of energy is required to break hydrogen bonds between molecules and thus break down the larger clusters of molecules into separate molecules. As with the boiling point, the melting point of water is abnormally high when compared with the melting points of the hydrogen compounds of the other elements from Group 16. These substances are chemically similar but have no appreciable hydrogen bonding.

Hydrogen bonding has an effect on the crystal structure of ice. X ray studies show that the three dimensional structure caused by hydrogen bonding gives the ice crystals a crystalline stricture with many hexagonal openings. This open structure accounts for the low density of ice. When ice is melted, hydrogen bonds are broken. Then the open structure is destroyed, and molecules move closer together Therefore, the liquid phase of water has a greater density than the solid phase. For most substances, the solid phase has a greater density than the liquid phase. .

Hydrogen bonds are mainly responsible for the coiled shape of protein molecules. Hydrogen bonds are found in nucleic acids DNA and RNA. The bonds hold together the double helix structure in the DNA.

Metallic Bonding

Most metals have only one or two valence electrons and low ionization energies. Their valence electrons are not tightly bound to the atom, but seem to move easily from one atom to another. They can be considered a part of the whole metal crystal. These mobile electrons exert an attractive force ors t he positive ions, helping to fix their positions. The ease with which the valence electrons move within the crystal distinguishes the metallic bond from the ionic or covalent bond.

1. Metals are good conductors of heat and electricity because of the mobility of their valence electrons.

2. The binding action of the electrons is the basis for the hardness of metals. Softer metals have weaker binding forces.

3. The high luster of metals is the result of the uniform way in which the valence electrons absorb and re‑emit light energy that strikes them.

4. The malleability (ability to be flattened into thin foil without breaking), ductility (ability to be stretched into a thin wire without breaking), and sectility (ability to be cut into sections without shattering) result from the fairly uniform attraction between the electrons and the ions. The ions can change position, or `flow' in the `sea' of valence electrons. Metals can be flattened or stretched into a wire because the electrons and ions can move into other positions without breaking up the essential structure. The attraction between electrons and ions continues even while forces are applied that change the shape of the metal.