Part 5 Metallic Bonding — structure and properties of metals. General introduction to nanoscience and commonly used terms explained. Revision notes on the theory of giant covalent bonding, which type of elements form giant covalent structures? Part 4. Large Covalent Molecules and their Properties. Macromolecules — giant covalent networks and polymers. Because covalent bonds act in a particular direction i. Its a good idea to have some idea of where the elements forming giant covalent structures are in the periodic table. The black zig—zag line 'roughly' divides the metals on the left from the non—metals on the right of the elements of the Periodic Table.
Chemical bonding comments about the selected elements highlighted in white.
The non—metallic elements carbon and silicon form giant covalent structures Materials that consist of giant covalent structures are solids with very high melting points and usually physically hard materials not graphite. All of the atoms in these structures are linked to other atoms by strong covalent bonds in specific directions eg a grain of sand silica is one giant molecule!
The Relationship between Molecular Structure and Bond Energy
These substances usually have an extended 3D network of strong covalent bonds. These bonds must be overcome to melt or boil these giant covalent substances and this requires very high temperatures to give the particles sufficient kinetic energy to weaken the bonds and cause melting of the substance. Diamond and graphite forms of carbon and silicon dioxide silica are examples of giant covalent structures.
You should be able to recognise giant covalent structures from diagrams showing their bonding and structure. You also need to be able to explain the properties of giant covalent substances in terms of their molecular structure. Most giant covalent structures don't have freely moving charged particles like ions or electrons to carry an electric current, so they are poor conductors of electricity. The structure of the four allotropes of carbon diamond, graphite, graphene and fullerenes , silicon and silicon dioxide silica DIAGRAMS It is possible for many atoms to link up to form a giant covalent structure or lattice.
The structures of giant covalent structure are usually based on non—metal atoms like carbon, silicon and boron. The atoms in a giant covalent lattice are held together by strong directional covalent bonds and every atoms is connected to at least 2, 3 or 4 atoms. What you might call 'atomic networking'! This very strong 3—dimensional covalent bond network or lattice gives the structure great thermal stability e.
This is because it takes so much thermal kinetic energy to weaken the bonds sufficiently to allow melting. The covalent bonds are very directional, giving rise to a strong and fixed network that we call a giant covalent structure. This gives them significantly different properties from the small simple covalent molecules see simple molecular substances.
This is illustrated by carbon in the form of diamond an allotrope of carbon. Pure silicon , another element in Group 4, has a similar structure. NOTE: Allotropes are different forms of the same element in the same physical state. They occur due to different bonding arrangements and so diamond , graphite below and fullerenes below are the three solid allotropes of the element carbon.
Oxygen dioxygen , O 2 , and ozone trioxygen , O 3 , are the two small gaseous allotrope molecules of the element oxygen. Sulphur has three solid allotropes, two different crystalline forms based on small S 8 molecules called rhombic and monoclinic sulphur and a 3rd form of long chain —S—S—S— etc.
A relatively large amount of energy is needed to melt or boil giant covalent structures because strong chemical bonds must be broken and not just weakening intermolecular forces as in the case of small covalent molecules like water. Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes. They are usually poor conductors of electricity because the electrons are not usually free to move as they are in metallic structures and they are NOT made up of ions.
All the valency bonding electrons are tightly held and shared by the two atoms of any bond, so in giant covalent structures they are rarely free to move through the lattice and not even when molten either, since these giant molecular covalent structures do NOT contain ions. Also, because of the strength of the bonding in all directions in the structure, they are often very hard, strong and will not dissolve in solvents like water.
The bonding network is too strong to allow the atoms to become surrounded by solvent molecules Silicon dioxide silica, SiO 2 has a similar 3D structure and properties to carbon diamond shown below and also pure silicon itself. Diamond is an allotrope of carbon - a distinct molecular form of solid carbon 1st allotrope of carbon discussed. Note: Allotropes are defined as different physical forms of the same element in the same physical state, so the atoms in allotropes are identical, but bonded in different ways, as you will see with the four forms of carbon.
In diamond every carbon atom is strongly linked to four other carbon atoms by strong directional covalent bonds giving a very three dimensional 3D strong lattice.
Van der Waals Forces Definition
This results in a very rigid strong structure. Theoretically in a diamond crystal all the carbon atoms are linked together. The result is a very pure crystal structure with a high refractive index that gives diamonds quite a sparkle as light passes through it. The strength and hardness of carbon in the form of diamond enables it to be used as the 'leading edge' on cutting tools, the hardness is derived from the very strong rigid three—dimensional carbon—carbon bond network. Diamond also has a very high melting point because of this very strong giant covalent lattice in which every carbon atom is strongly bonded to four other carbon atoms see diagram above on right.
It takes a lot of energy to break overcome the carbon-carbon bonds and cause melting. The more energy needed, the higher the melting point. Under very high pressure, diamond will melt at a staggering o C! The strong bond network in diamond and graphite and silica prevents these materials from dissolving in any conventional solvent. Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances is given in a section of the Energetics Notes.
Pure elemental silicon not the oxide has the same molecular structure as diamond and similar properties, though the 3D giant covalent bond network is not as strong, so elemental silicon is not as high melting as carbon i the form of diamond. The molecular diagram is the same for Si n , where n is a huge number! Silicon melts at o C, and has poor electrical conductivity no free electrons or ions and won't dissolve in any solvent. The silicon in the transistors of electronic devices is 'doped' with other elements to increase its electrical conductivity.
Many naturally occurring minerals are based on —O—X—O— linked 3D structures where X is often silicon Si and aluminium Al , three of the most abundant elements in the earth's crust. Silicon dioxide 'silica' is found as quartz in granite igneous rock and is the main component in sandstone — which is a sedimentary rock formed the compressed erosion products of igneous rocks. Looking at the diagram on the right of silica, each silicon atom black blobs forms four strong covalent bonds with the linking oxygen atoms yellow blobs.
You can also see from the diagram that there are two oxygen atoms to every silicon atoms giving the empirical formula SiO 2. Again like diamond, theoretically all the atoms in a silica crystal are linked together by a strong 3D covalent bond network giving a strong rigid structure. It takes a lot of energy to break overcome the strong silicon-oxygen bonds in the giant covalent lattice of silicon dioxide silica , hence the high melting point of o C.
Therefore Silica SiO 2 is a very hard substance with a very high melting point and won't dissolve in any solvent. There are no free electrons so silicon dioxide doesn't conduct electricity. Many more minerals that are hard wearing, rare and attractive when polished, hold great value as gemstones , but sand is also mainly silica, but not quite as valuable! Graphene is a single sheet of graphite, therefore it is another form of the non-metallic element carbon.
You can consider it as a 3rd allotrope of carbon to be described here. Graphene is a 2D nanomaterial because each layer is only one carbon atom thick made up of a huge number of linked hexagonal rings of carbon atoms. Like graphite, graphene conducts electricity because delocalised electrons can run through the layer. Because of the strong carbon-carbon bonding, it is a very strong light-weight material with a higher tensile strength than steel. They consists of a 3D connected arrangement of pentagonal, hexagonal or heptagonal rings like graphite.
Alternating pentagonal rings of carbon atoms allow curvature of the spherical surface, in fact curved sufficiently to form their characteristic 'football' or 'rugby ball' shapes and elongated closed tubes. Some fullerenes have rings of seven carbon atoms, again to allow curvature of the surface of the hollow sphere of the 'bucky ball'.
This means fullerenes are 'hollow' molecules in which other molecules can fit in.
The Additivity of the Energies of Normal Covalent Bonds.
Buckminster Fullerene C 60 ,the first to be discovered, is shown on the right and the bonds form a pattern like a soccer ball. Others are oval shaped like a rugby ball. It is a black solid insoluble in water. All of the fullerenes are hollow with the hexagonal and pentagonal rings of carbon atoms forming the surface. These 'molecular size' fullerene particles behave quite differently to a bulk carbon materials like graphite or diamond. Fullerenes are NOT considered giant covalent structures and are classed as simple molecules.
Fullerenes do dissolve in organic solvents giving coloured solutions e.
Fullerene molecules may be used for drug delivery into the body, they can enclose or 'cage' another molecule e. Fullerenes are used as lubricants, the molecules readily slide over each other. Fullerenes can be part of a catalyst composite, catalysts can be attached to their large surface area. Fullerenes in the form of carbon nanotubes can be used for reinforcing composite materials, eg sports equipment like tennis rackets.
More on nanotubes below.
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Carbon nanotubes are basically long cylindrical fullerenes. Carbon nanotubes have a high length to diameter ratio Carbon nanotubes Breaking a noncovalent bond involves modifications of the molecular structure on a nanometer scale: this requires breaking and rearranging many van der Waals, hydrogen, or ionic bonds as well as stretching covalent bonds. Weak bonds are noncovalent bonds corresponding to the typical interaction between two small molecules such as ions. In water, salt dissociates into positively-charged sodium ions and negatively-charged chlorine ions.
Although water significantly screens reduces the electrostatic force between ions, the ions nevertheless experience a weak electrostatic bonding force. The strength of this bond is sufficiently low that molecular Brownian motion thermal jostling often bumps one ion out of its binding with its counterpart. We just discussed the ionic nature of weak bonds, but there exist other kinds of weak interactions like hydrogen bonding or the hydrophobic interaction.
All these interactions have in common that their strength is comprable to that of thermal molecular agitation. The paradigm of such a weak force is the base-pairing interaction which holds together complementary DNA bases and involves either two or three hydrogen bonds.
Hydrogen bonding in ionic liquids - Chemical Society Reviews (RSC Publishing)
The strength of a weak bond is given by its typical energy which is on the oder of k B T. T, measured in units of degrees Kelvin, is K at room temperature. We have said that Brownian motion is a big issue at the molecular level, and at the cellular level on the scale of microns , this phenomenon is still quite strong.
If, for instance, we consider an E. If we imagine a plane separating the bacteria into two equal parts, one could argue that the same number of particules should hit both sides of the plane producing a perfect cancellation of this molecular bombardment and thus no net motion. This arguement is true but only statistically, that is if we average on a huge number of molecular collisions. Consider what occurs in a tiny slice of time: Nr is the number of molecules pushing the bacteria to the right, Nl is the number of molecules pushing it to the left.
In other words, they are not perfectly equal for any given tiny slice of time. N is proportional to the cross-section of the object: the larger the object, the more water molecules will bombard it. The small mismatch between Nr and Nl will produce a force and thus a displacement of the bacteria in a random direction. This random force is called the Langevin force, and its strength is thus related to the size of the object. Note, however, that this strength must be compared to that of the viscous drag experienced by the object in the fluid.
For a micron-size object such as an E. It is interesting to compare this number with the weight of E. Imagine pulling the poor bacteria out of the water and weighing it with a scale, its typical weight will be on the order of 0. To get an idea of the violence of Brownian motion at the scale of a bacteria, imagine that you are swimming in a pool where every second you recieve a random blow whose strength is comparable to your own weight!
Just above these random bombardment forces lie the entropic forces that result from a reduction of the number of possible configurations of the system consisting of the molecule e. Upon stretching, the molecular entropy is reduced so that at full extension there is but one configuration left: a straight polymer linking both ends.
To reach that configuration, work against entropy has to be performed, i. Entropic forces are rather weak. How strong is a covalent bond? Grandbois, MK. Beyer, M. Rief, H. Clausen-schaumann, HE. Gaub, Science p. Florin, VT.