ORDINARY MATTER #3

Greetings to the community and the entire world. Still on ordinary matter physics; today I will be discussing the different types of bonding in solids and liquids and also on crystal structures.

IONIC BONDING

This is found in many crystalline materials that consist of at least two different types of atom, such as sodium chloride- common salt. Atoms try to fill their outer shell of electrons. For example, one type of atom, the sodium atom in the case of sodium chloride, loses an electron and the other, chlorine, gains an electron. Once the sodium has lost an electron it becomes positively charged and is said to be a positive ion: by gaining an electron the chlorine becomes a negative ion. The crystal of sodium chloride must be electrically neutral everywhere. The only way this is possible is for the sodium and chloride ions to alternate in nearest-neighbour directions.

Sodium and fluorine atoms undergoing a redox reaction to form sodium fluoride. Sodium loses its outer electron to give it a stable electron configuration, and this electron enters the fluorine atom exothermically. The oppositely charged ions – typically a great many of them – are then attracted to each other to form a solid.

^{Wdcf, CC BY-SA 3.0}

Electrical forces act over a relatively long distance, so bonding is not directional but it is strong. It is difficult for ions of opposite charge on different planes to slip past each other, so ionic crystals are not plastic. They usually fracture before reaching their elastic limit.

Ionic crystals are poor conductors of electricity as neither the ions themselves nor individual electrons are free to move. When put into water, ionic crystals dissolve: the water reduces the electrical forces between the ions and so the ions can separate. They now move around and can conduct electricity (they move, carrying charge).

COVALENT BONDING

In this case, atoms share, rather than exchange, electrons in the region between the two atoms. Unlike the ionic bond, the covalent bond is highly directional. Molecules produced by such bonding retain a definite shape.

Oxygen, O₂, is a covalently bonded molecule. Both oxygen atoms would like two more electrons in their outer shells. They achieve this by each providing two electrons for sharing (making a total of four shared electrons).

Covalent bonding can lead to highly extended 'molecules'. Diamond is such an example. Carbon has four electrons in its outer shell and by sharing a further four electrons, one from each of four other carbon atoms, it forms a continuous structure. The four bonds are symmetrically arranged in space and are called tetrahedral bonds.

A covalent bond forming H2 (right) where two hydrogen atoms share the two electrons

^{Jacek FH, CC BY-SA 3.0}

Silicon and germanium also have four electrons in their outer shells (carbon, silicon and germanium are all from Group IV of the Periodic Table), and so they bond in the same way. The arrangement of the outer electrons of silicon and germanium gives rise to their semiconducting property. This property has been the basis of the development of semiconductor devices and the advance of much of the electronics industry. Unlike ionic compounds, covalent materials do not conduct in aqueous solution.

METALLIC BONDING

In metals, each atom has one or more outer delocalised electrons that are free to move between atoms. The metal can be pictured as a giant structure of positively charged ions in a sea of electrons. We might expect these ions to repel each other and fly apart. Not so. The interaction between the electrons and the ions sets up a counteracting attractive force. Why this should happen was not properly understood until physicists came up with a new theory called quantum physics.

The fact that the electrons can move freely around accounts for the good electrical conductivity of metals. The interaction between the large number of ions and the 'sea' of electrons is not directional. The ions, though, stay fixed, often close packed together. If a tensile or shearing force is applied to the metal, it is easy for the planes of ions to move past each other. So metals exhibit plasticity.

VAN DER WAALS BONDING

Even neutral atoms attract each other weakly. The atoms are composed of both positively charged protons in their nuclei and moving electrons in a surrounding cloud that constantly changes shape. In a pair of like atoms, the positive and negative charges from one atom and the positive and negative charges from the other produce a small instantaneous force. Although the force changes as the electrons move, it averages over time to produce an attraction that is called van der Waals bonding and occurs between all atoms and molecules.

Gecko climbing a glass surface as a result of Van der Waals force.

^{w:User:Lpm, CC BY-SA 3.0}

CRYSTAL STRUCTURES

Adding together layers of spheres is similar to adding together layers of atoms. The final structure will depend on how we build up the layers and whether these layers themselves are made up of identically sized spheres or spheres of different sizes.
Substances that are built from an ordered stacking of atoms are called crystals. These often possess flat faces that match regular planes of atoms. Most crystals can be cleaved. This means that a hard blow with a sharp edge will split the crystal along a plane to create a pair of flat surfaces. Because they can be cleaved, diamonds and other gemstones have attractive, shiny surfaces and are valuable for jewellery.

SIMPLE CUBIC STRUCTURE

Let us continue to think of atoms as hard spheres. They can be arranged so that they touch in a single layer as a square array. Note that there are four atoms round a hollow. Next we can arrange another layer of spheres vertically above the first, and so on as in the figure. Such a three-dimensional arrangement is called simple cubic.

The primitive and cubic close-packed (also known as face-centered cubic) unit cells

^{Owen Graham, Public Domain}

Taking eight atoms from the arrangement in the figure and reducing their size, their centres are shown joined by some lines. Note that each atom will also be at the corners of eight cubes, one and a further seven cubes. Alternatively, looking at the figure with the atoms (spheres), it is clear that the part of an atom at each corner is an eighth of an atom.

CLOSE PACKING STRUCTURES

The spheres in the figure are not packed together as closely as possible. The figure below shows a different arrangement of spheres in a single layer. Here, a second row of touching atoms is displaced relative to the first one. The third row is a repeat of the first row. Comparing with the previous figure, note that there are only three atoms at a hollow but that the closest separation of centres of the atoms remains the same, equal to the diameter of the atom. The result is that in this figure there are more atoms per unit area.

Illustration of the close-packing of equal spheres in both hcp (left) and fcc (right) lattices

^{en:User:Twisp, Public Domain}

Building close packed structures: hcp and fcc

We can build a regular close packed structure on the layer of the figure above in two different ways, as follows:

Method 1

We call our first layer of close packed spheres (atoms) layer A. Notice that each sphere touches six others (nearest neighbours) and is surrounded by six hollows. Now we start to build up a three-dimensional structure. We place the next plane of spheres into hollows. Notice that, even if we make the next plane of spheres close packed, we can only centre spheres in half the hollows. At this stage it doesn't matter which half-set of hollows we use. We label this second layer B.

Next, we add the third layer of spheres and again we can only fill half the hollows. This time the choice makes a difference.

By filling one set of hollows, we can place all the spheres in this third layer vertically above spheres in the first layer. So we label this layer as another A layer. Note the hollows in layer B of the figure in which the third-layer spheres are placed.
We can picture a hexagonal cell made up within these three layers, showing the conventional picture of the hexagonal cell with the spheres (atoms) shrunk to a small size. It is easy to see that we can continue with this ABABAB… arrangement of the layers. We call this pattern a hexagonal close-packed (hcp) structure.

Hexagonal close packed (hcp) unit cell

^{Dornelf, CC BY-SA 3.0}

Method 2

Let us return to packing our third layer and look again at the figure. We already have layers A and B. We now choose the alternative half-set of hollows to the set we chose previously. The spheres of the third layer no longer lie vertically above the spheres of layer A, nor do they lie above the spheres of layer B. So we label this third layer C. We go on to consider a fourth layer. Again there is a choice of hollows, but we choose to place the spheres in hollows such that they lie vertically above the spheres in layer A. This produces another A type layer. Packing on upwards produces an ABCABCABC...pattern. The figure below shows such layers packing in a pyramidal form starting from a triangular base. The structure is still close packed and the arrangement is called face-centred cubic (fcc) close packing.

But why cubic? As the name suggests, a face-centred cube has atoms at the corners of a cube and atoms at the centres of each face. The close packing structure that I have described can be related to a series of face-centred cells by tilting the close packed planes. It is tricky to see but you should be able to work out the relationship in the figures below. The close packing planes link up three diagonally related corners of a cube. Try to identify how an atom in such a plane will be surrounded by six nearest neighbours within the same plane, and by three atoms in each of the two adjacent planes. All the outer surfaces of the pyramid shown are of close packed planes of atoms.

Face-centered cubic crystal structure

^{Daniel Mayer and DrBob, CC BY-SA 3.0}

BODY-CENTRED CUBIC STRUCTURE

Another common crystal structure is called body-centred cubic (bcc) because there is an atom at the centre of the cube (that is, at the centre of its 'body') in addition to the atoms at the corners. Iron has this structure up to 800 °C, at which temperature it changes to fcc. This is an example of a solid changing its phase but remaining a solid. Heating iron to above 800 °C and quenching it in water prevents the iron from returning to a bcc structure. The face-centred form of iron has different properties from the body-centred form - in particular it is very brittle. Iron containing carbon as an impurity gives steel. Such variations in the properties of iron are immensely important for the science of metallurgy and for the iron and steel industry.

I will pause here for now, and discuss in my next post, how to determine the structures of crystals. I will afterwards explain the behavior of liquids and gases. But till then, I remain my humble self, Hive account@emperorhassy.