Other Types of Bonding, Chemistry tutorial


In the preceding chapter, we have studied about the 2 major significant kinds of bonding: ionic bonding and covalent bonding. Both of such are eventually driven via the want that atoms have to be bounded through a whole shell of electrons. They attain this via correspondingly either increasing or losing, or sharing, one or more electrons.

Intermolecular attractions are attractions between 1 molecule and a neighboring molecule. The forces of attraction which hold an individual molecule together (for instance, the covalent bonds) are known as intramolecular attractions. Such 2 words are so confusingly similar that it is safer to abandon one of them and never utilize it. The term "intramolecular" will hence, not be utilized again in this chapter.

All molecules experience intermolecular attractions, although in several cases those attractions are extremely weak. Even in a gas like hydrogen, H2, if we slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together eventually to form a liquid and then a solid.

In hydrogen's case, the attractions are so weak that the molecules have to be cooled to 21K (-252°C) before the attractions are sufficient to condense the hydrogen as a liquid. Helium's intermolecular attractions are even weaker - the molecules will not stick mutually to form a liquid until the temperature drops to 4 K (-269°C). 

Metallic bonding

A metal is an element that can simply lose up to 3 electrons, therefore forming positive ions. The lost electrons though come together to form a joined 'sea' or 'cloud' of electrons. Metals consist of a lattice of positive ions existing in a freely moving 'sea' of 'cloud' of electrons that attach them together. Electrons in the 'sea' don't belong to specific metal atoms and move easily throughout the assembly. Metals are well known to be solids (except for Mercury). The bonds between metals can loosely be  explained as  covalent bonds  (due to sharing electrons), except that the metal atoms do not just share electrons through 1, 2, 3 or 4 neighbours, as in covalent bonding, but by many atoms. The structure of the metal is found out by the fact that each atom tries to be as close to as many other atoms as possible. This is shown here for one typical metal structure (assumed, for instance, through iron at several temperatures):

318_The Structure of a Metal.JPG

Fig: The Structure of a Metal

Contrast this through the structure of diamond seen previously. Because the electrons are shared with all the neighbours, it is quite easy for the electrons in metals to move around. If each "shared" electron shifts one atom to the right or left, this leads to a net shift of charge. This occurs quite easily in metals, but much less so in ionic solids, or covalent ones, where the electrons are rigidly associated with either a particular atom or ion, or a particular pair of atoms. It is because electrons can move around so easily inside metals that the latter conduct electricity.

Explaining the Physical Properties of Metals

This strong bonding generally consequences in dense, strong materials through high melting and boiling points.

Metals are good conductors of electricity since such 'free' electrons carry the charge of an electric current whenever a potential difference (voltage!) is applied across a piece of metal. 

Metals are as well good conductors of heat. This is in addition due to the free moving electrons; the 'hot' high kinetic energy electrons shift around freely to move the particle kinetic energy more proficiently to 'cooler' atoms. 

Typical metals as well have a silvery surface but remember this might be simply tarnished via corrosive oxidation in air and water. 

Unlike ionic solids, metals are extremely malleable; they can be eagerly bent, pressed or hammered into shape. The layers of atoms can slide over each other with no fracturing the structure (Figure). The cause for this is the mobility of the electrons. When planes of metal atoms are 'bent' or slide the electrons can run in between the atoms and sustain a strong linking situation. This can't take place in ionic solids. 

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Fig: Alloy Structure

Figure (1) illustrates the regular arrangement of the atoms in a metal crystal and the white spaces demonstrate where the free electrons are (yellow circles really positive metal ions).

Figure (2) illustrates what take places when the metal is stressed via a strong force. The layers of atoms can slide over each other and the bonding is sustained as the mobile electrons keep in contact through atoms, so the metal continues intact BUT a different shape.

Figure (3) illustrates an alloy mixture. It isn't a compound but a physical mixing of a metal plus at least one other material (one can be another metal for example Ni, a non-metal for instance C or a compound of carbon or manganese, and it can be bigger or smaller than iron atoms). Many alloys are generated to provide a stronger metal. The occurrence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and decreases the 'slip ability' of one layer next to another. The consequence is a stronger harder less malleable metal.

The main point about using alloys is that we can make up, and try out, all sorts of dissimilar compositions until we find the one that best suits the needed purpose

Hydrogen bonding

In covalent bonds, the electrons are allocated, so that each atom obtains a filled shell. When the distribution of electrons in molecules is considered in detail, it becomes apparent that the 'sharing' isn't always perfectly 'fair': often, one of the atoms gets "more" of the shared electrons than the other does.

This take places, in particular, when atoms these as nitrogen, fluorine, or oxygen bond to hydrogen. For instance, in HF (hydrogen fluoride), the structure can be illustrated through the following "sharing" picture:

2138_Structure of Hydrogen Fluoride.jpg

Fig: Structure of Hydrogen Fluoride

Though, this structure doesn't tell the whole truth about the distribution of electrons in HF. certainly, the subsequent, 'ionic' structure as well respects the filled (or empty) shell rule:

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Fig: Structure of HF

In reality, HF is explained through both such structures, so that the H-F bond is polar, through each atom bearing a small positive (δ+) or negative (δ-) charge. When 2 hydrogen fluoride molecules come close to each other, the like charges attract each other, and one gets a 'molecule' of di-hydrogen fluoride as revealed:

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Fig: Hydrogen Bonding in Hydrogen Fluoride

The weak 'bond' between the F atom and the H is termed a Hydrogen Bond, and is shown here as the dotted green line. Hydrogen bonds can as well take place between oxygen atoms and hydrogen. One of the most significant kinds of hydrogen bonds is of this type, and is the one occurring in water. As conversed for HF, the electrons in H2O molecules aren't evenly 'shared': the oxygen atom has more of them than the hydrogen atoms. As a consequence, oxygen has a (partial) negative charge, and the 2 hydrogen atoms have a positive charge. When we have 2 water molecules close to another, a hydrogen atom on one of the molecules is attracted to the oxygen of the other molecule, to provide a dimer. The structure of this dimer is following here:

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Fig: Three-Dimensional Structure of Water

As we know that how the oxygen, hydrogen, and oxygen atoms included in the hydrogen bond are arranged more or less in a straight line. This is the termed geometry for hydrogen bonds, and clarifies why only one hydrogen bond can be shaped in the water dimer.

Upon going to 3 water molecules, it is now possible to form numerous hydrogen bonds. This is given here:

In liquid water or ice, many water molecules are seal to each other, and they shape dense networks of hydrogen bonds. In ice, the arrangement of the water molecules through respect to each other is regular, whereas in water, it is random. Figure illustrates a typical arrangement of water molecules similar to what we might discover in the liquid:

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Fig: Hydrogen Bonding in Water

Such bonds are weaker than typical covalent or ionic bonds, but however strong sufficient to make molecules that can hydrogen bond much more 'sticky' by respect to each other than are other covalent molecules through otherwise alike properties. For instance, the molecular mass of water is 18, and that of nitrogen is 28, yet nitrogen is a gas down to approximately degrees centigrade, while water is a liquid up to 100 degrees.

Hydrogen Bonds in Biology

The cells of living things are made up of many different sorts of molecule. Two significant classes of molecule are proteins and nucleic acids. In both of such molecules, parts of the (extremely huge) molecules are included in hydrogen bonds through other parts of the similar molecules. This is extremely significant in establishing the structure and properties of such molecules. Hydrogen bonding plays significant role in the structure of DNA (Deoxyribonucleic acid), one of the most vital nucleic acids, and explaining the vital role of hydrogen bonding.

Other intermolecular forces

Intermolecular forces are forces of attraction or repulsion which act between neighbouring particles: atoms, molecules or ions. They are weak compared to the intramolecular forces, the forces that keep a molecule together. For instance, the covalent bond current within HCl molecules is much stronger than the forces present between the neighbouring molecules, which exist when the molecules are sufficiently close to each other. Hydrogen bonding is an instance of intermolecular forces.

London dispersion forces

London dispersion forces (LDF, as well recognized as dispersion forces, London forces, instantaneous dipole-induced dipole forces) is a kind of force acting between atoms and molecules. They are part of the van der Waals forces. The LDF is named after the German-American physicist Fritz London.

The LDF is a weak intermolecular force arising from quantum induced instantaneous polarization multi-poles in molecules. They can hence act between molecules with no permanent multi-pole moments. London forces are exhibited through non-polar molecules because of the correlated movements of the electrons in interacting molecules. Because the electrons from different molecules start 'fleeing' and avoiding each other, electron density in a molecule becomes redistributed in proximity to another molecule, (see  quantum mechanical theory of dispersion forces). This is frequently explained as formation of 'instantaneous dipoles' that attract each other. London forces are present between all chemical groups and usually represent the main part of the whole interaction force in condensed matter, even though they are usually weaker than ionic bonds and hydrogen bonds. This is the only attractive intermolecular force present between neutral atoms (for instance a noble gas). With no London forces, there would be no attractive force between noble gas atoms, and they would not exist in liquid form.

London forces become stronger as the atom or molecule in question becomes larger. This is due to the raised Polarizability of molecules through larger, more dispersed electron clouds. This trend is exemplified by the halogens (from smallest to largest: F2, Cl2, Br2, I2).  Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. The London forces as well happen to stronger through superior amounts of surface contact. Greater surface area means closer interaction between dissimilar molecules.

Dipole-dipole interactions

The non-symmetrical distribution of charge within a molecule polarizes the molecule into positive and negative poles such that there exist electrostatic interactions between close molecules called a permanent dipole-permanent dipole attraction. Hydrogen chloride, HCl, is made up of a positively charged end and a negatively charged end such that there is charge interaction forming a weak bond between the hydrogen atom of one hydrogen chloride molecule and the chlorine end of another molecule.

Trichloromethane (chloroform), CHCl3, is another instance of molecules by dipole-dipole attraction where polar molecules are held together more strongly than non-polar molecules of comparable mass. Dipole-dipole forces are:

Stronger intermolecular forces than dispersion forces 

Occur between molecules that have permanent net dipoles (polar molecules), for instance, dipole-dipole interactions take place between SCl2 molecules, PCl3 molecules and CH3Cl molecules. If the permanent net dipole sinside the polar molecules consequences from a covalent bond between a hydrogen atom and either fluorine, oxygen or nitrogen, the consequential intermolecular force is termed to as Hydrogen Bonding.

Established if the incomplete positive charge on one molecule is electrostatically attracted to the partial negative charge on a neighbouring molecule. 

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