Molecular bonds are formed in case of those elements or compounds whose electron configuration is such that little transfer takes place between atoms. These bonds are formed as a result of weak van der Waals forces of attraction which exist between various atoms. These forces are due to the electrostatic attraction between the nucleus of one atom and the electrons of the other. This is largely but not completely neutralized by the electrostatic repulsion of the nucleus of one atom by the nucleus of the other.
The resultant weak attraction between the two atoms is called van der Walls force. These bonds are made possible largely because the electrons of adjacent atoms in a molecule tend to repel each other. As the electrons rotate around their nuclei, they end to keep in phase for a hydrogen molecule.
The result is that the molecule has a small fluctuating net charge on each end and acts as an oscillating dipole. The hydrogen molecule is instantaneously charged negatively on the right end and positively on the left. This fluctuating charge on one molecule tends to interact with the fluctuating charge on a neighbouring molecule, resulting in a net attraction.
The strength of the bond depends on the ease with which one atom can influence the other. Molecules of inert gases which consists of single atoms, are held together by dispersion forces when the gases are solidified. In many organic solids the most important bonding forces between molecules are of this type. In the covalent bonded hydrogen chloride HCl molecule the net effect of the electron-sharing process is to give the chlorine atom a slightly negative charge while the hydrogen atom has a corresponding positive charge.
The charges are actually very small being 0. The spacing of the atoms is 1. Adjacent HCl molecules therefore, attract each other by means of the electrostatic attraction between their oppsitively charge ends. The attraction is small compared with that between ions because the charge on an ion is at least equal to that of one electron, 4. Dipole bonds are much weaker than bonds, but at the same time they are considerably stronger than disperson bonds. Some other materials subject to dipolar bonding, and their dipole moments in cm x esu are —.
The hydrogen bond might be considered as a special type of dipole bond, but one that is considerably stronger. It occurs between molecules in which one end is a hydrogen atom.
The one electron belonging to the hydrogen atom is fairly loosely held, and if the adjacent atom in the molecule is strongly electro-negative, it may keep all the electrons around itself, leaving the hydrogen atom in effect a positive ion.
A strong permanent dipole is created that can bond to other similar dipoles with a force near that involved in the ionic bond. A good example of hydrogen bonding is water or ice. In water the hydrogen and oxygen atoms are held by covalent bonds in a configuration as shown in Fig. It is seen that this special form of dipole bond is connected as the dipole moment of the H-0 bond rather than that of the molecule as a whole. It is possible to have mixed bonding. In fact bonding between atoms in many materials cannot be classified as one of the four ideal types, i.
The metals gradually change from pure metallic bonding as in sodium to the less perfect metals such as tellurium and arsenic, finally reaching the pure covalent bonding of carbon in the form of diamond. Starting at diamond, we find graphite, benzene rings and high polymers. Ultimately, we reach the rare gases and pure van der Waals bonding. Engineering , Solids , Bonds , Bonds in Solids. Structure of Metals and Alloys Metallurgy. Glass: Structure, Constituents and Properties Engineering.
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For example, in sodium chloride, the negatively charged chloride ions are much larger than the positively charged sodium ions. As a result, the structure of sodium chloride is a little different than you may have been led to expect.
This is illustrated in the following figure:. The type of crystal structure of a particular ionic compound frequently depends on the ratio of the sizes of the anion and cation. It was explained that metals are good conductors of electricity and heat, have high malleability bendability , high ductility can be made into wires , and are shiny.
What we never explained was why metals exhibit these properties. As it turns out, the properties of metals stem from the nature of metallic bonds. One very simple model used to explain bonding in metals is referred to as the electron sea theory. In the electron sea theory, the cations in a metallic solid remain in stationary crystalline positions while the valence electrons from each metal are free to wander throughout the entire solid:.
Figure This theory does a good job of explaining the properties of metals. Because electrons are able to move freely throughout the entire solid, metals are excellent conductors of electricity. The high mobility of electrons also causes metals to conduct heat because they do a good job of dispersing energy. Because metal nuclei can move from place to place without causing bonds to be broken, metals are both malleable and ductile. Though the electron sea theory accurately describes the properties of metals, it glosses over how the electrons are able to wander freely throughout the solid.
After all, didn't we spend a great deal of time learning about how electrons exist within orbitals? Think back to when we discussed hybrid orbitals Bonding and Structure in Covalent Compounds.
When an s-orbital and three p-orbitals overlap, they form four sp3 orbitals. If an s-orbital and two p-orbitals overlap, they form three sp2 orbitals. In a metal, a similar thing happens. However, unlike covalent compounds, where only a few orbitals mix, all of the metal atoms mix their atomic orbitals s-, p-, and d-orbitals together to form a huge number of orbitals known as "molecular orbitals.
Eventually, these orbitals become so close in energy that they form a conducting band in which electrons can easily jump from one orbital to another. The molecular orbital theory that explains electron motion in metals is also referred to as "band theory. Because there are more molecular orbitals in the conduction band than there are electrons, it doesn't take much energy to raise an electron from a filled orbital to an empty one with higher energy.
When these electrons jump to empty orbitals, they are able to move freely around the metal. As a result, it's easy for electrons to jump to empty orbitals where they are free to move around the solid.
There are typically several conducting bands in metals. One band called the "s-band" is caused by an overlap between all of the s-orbitals in the metal. The other bands are called the "p-band" and "d-band" because they result from the overlap of p- and d-orbitals, respectively. Because these bands overlap in energy, they behave as one large, partially filled band:.
Other elements are sometimes added to metals to give them desired properties such as hardness, durability, or strength. The resulting material is referred to as an alloy. An alloy is a metallic material in which several elements are present. The elements added to a pure metal to form an alloy are selected to maximize a desired property. Network atomic solids are formed when many atoms are bonded together covalently to form one gigantic molecule.
Unlike regular covalent molecules that are generally small, network atomic solids may grow quite large. One common example of a network atomic solid is a diamond:.
As a result, diamonds can be thought of as being very large covalent molecules. Network atomic solids have a wide number of varying properties. They are usually hard, owing to the strong bonds between neighboring atoms?
They also tend to have high melting and boiling points due to the very strong covalent bonds. Network atomic solids are frequently brittle because a small movement of atoms in the crystal tends to disrupt the network of covalent bonding. Aside from diamonds, some common network atomic solids are quartz SiO2 , graphite, and silicon.
When elements form network atomic solids, their atomic orbitals s- and p-orbitals overlap to form conducting bands in the same way that metals do. However, there is one major difference in the structure of the bands between metals, nonmetals, and metalloids.
Whereas the s- and p-bands overlap in metals to form a giant conducting band, they don't for either nonmetals or metalloids, making it difficult for either to conduct electricity:.
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