Again, this is a chemistry assignment. The official title is “Compare and Contrast the Physical Properties of Iodine, Graphite, Water and Sodium Chloride. Pay Particular Attention to Their Solid States”. I found writing it was slightly repetitive, but it really helped me revise bonding and structure. I have removed citations as WordPress could not cope with the formatting when I copied this across from Microsoft word.
Iodine, Graphite, Water and Sodium Chloride exhibit all forms of bonding, with the exception of metallic; and all forms of intermolecular forces, with the exception of permanent dipole-dipole excluding water, which does experience permanent dipole-dipole forces. Consequently, all of these substances have a diverse range of physical properties.
Firstly, Iodine is a halogen and, therefore, diatomic. There is a single covalent bond between the two non-metal atoms as each atom shares one electron with one other atom in an iodine molecule. The molecules are then attracted to each other via Van der Waals forces. Iodine is a solid at room temperature with a melting point of 386.9K and a boiling point of 457.6K . Also, iodine is classified as an insulator of both electricity and heat, is soft and breaks easily, and has a very low solubility in water. All of these physical properties are a result of iodine having a molecular crystalline structure, as highlighted in figure 1. Molecular crystals consist of simple covalent molecules held together in a regular pattern by intermolecular forces [Van der Waals for iodine].
Figure 1: the crystalline structure of iodine
Despite being simple molecular, iodine is solid at room temperature because the Van der Waals forces are relatively strong in iodine compared to the Van der Waals forces in bromine or fluorine, for example. With the electron configuration of [Kr] 4d10 5s2 5p5 iodine has 53 electrons. Because these electrons are mobile within the molecule at any one moment they may collect towards one end of the molecule making that end δ- with a high electron density and the other end with a temporarily a low electron density becomes δ+. This temporary polarity of the molecule can cause an induced dipole in a nearby molecule, which will then be positioned with the δ+ end facing the δ- end of the other. In iodine this Van der Waals attraction covers many molecules with the δ+ end always facing the δ- end of another as shown in figure 2 below.
Figure 2: the induced dipole with the δ+ end always the δ– end
All of the halogens experience Van der Waals forces, but only Iodine and Astatine are solids at room temperature. This is because the high numbers of electrons in these atoms allow the difference between the δ+ and δ- ends of the molecules to be large. This relatively large difference between positive and negative cause relatively strong forces of electrostatic attraction between molecules of iodine compared to chlorine, for instance; therefore, larger amounts of energy are required to break the Van der Waals forces in iodine than in chlorine, hence the relatively high boiling point.
What is more, Iodine does not conduct heat or electricity as the diatomic molecular structure ensures that both atoms have a full outer shell of electrons with no spare. As a result, there are no electrons which are free to carry charge and create a current.
Moreover, iodine crystals are soft and break easily because although the atoms in an iodine molecule are joined by strong covalent bonds, the individual molecules are only held together by weak Van der Waals bonding. These weak intermolecular forces of attraction make it easy to separate the molecules from one another. From this it can be hypothesised that iodine has a very low tensile strength; however, no official value has been determined.
Furthermore, Iodine is almost completely insoluble in water because the intermolecular Van der Waals forces interacting with the permanent dipole forces of water would not be very strong compared to only the hydrogen bonds of pure water and so dissolving is not favoured by the molecules. The water molecules would rather remain hydrogen bonded to each other, than to let iodine molecules come between them. As a general rule polar substances dissolve in polar solvents and nonpolar substances dissolve in nonpolar solvents and this explains why iodine, a non-polar substance, is not soluble in water, which is a polar substance. This also explains why iodine is very soluble in organic solvents as they are non-polar. To further explain, the weak Van der Waals forces between iodine molecules can be replaced by slightly stronger Van der Waals forces between the iodine and the organic solvent molecules because this is a more stable arrangement and, therefore, favoured.
In addition, another interesting property of iodine is the way it changes state: straight from its solid crystalline structure to a violet coloured gas. This is process is called sublimation and happens because the Van der Waals forces in iodine are too weak to sustain liquid state under atmospheric pressure. However, when the iodine is heated further, the gas to solid ratio becomes imbalanced and vapour pressure increases until the iodine vapour is under so much pressure it eventually melts at 386.9 K and then boils at 457.6 K.
In contrast, Graphite is an allotrope of carbon with a giant covalent structure. It has very high melting and boiling points, is insoluble in water but is also soft and brittle and conducts electricity, which seems to contradict the general properties of giant covalent substances. This contradiction is a result of each carbon atom only being covalently bonded to three other carbon atoms, which form layers of atoms in a hexagonal arrangement. These sheets of carbon have delocalised electrons [one for every atom] above each layer. As mentioned, graphite is a soft substance despite having a strong covalent structure. This is because the strong covalent bonds only connect the carbon atoms in two dimensions with a 120˚ bond angle of a trigonal planar. These “sheets” of carbon are weakly attracted to each other due to the delocalised electron from each atom, which allows Van der Waals forces to act. As the delocalised electrons move around the sheet, very large temporary dipoles form, which induce opposite dipoles in the sheets above and below. These large areas of δ+ and δ- cause forces of electrostatic attraction between the layers of carbon. Nevertheless, these Van der Waals forces are still relatively weak and can be overcome by low amounts of energy, hence graphite being a soft substance. The ability of the layers of graphite to easily slide over one another without breaking is the reason why graphite is often used as a lubricant [little friction is required to overcome the intermolecular forces]. The Van der Waals forces within graphite are stronger than in iodine as the difference between δ+ and δ- is greater, thus graphite is stronger than iodine.
Figure 3: the structure of graphite
It is reported that graphite’s tensile strength varies between 4.8 and 76 MPa (AZoM, 2016). These tensile strength figures are relatively low. Graphite is weaker than a human femur bone, which has a tensile strength of 130MPa and with an average tensile strength of 40.4 [] graphite has the same tensile strength as pinewood. These statistics could be higher if they were in reference to carbon nanotubes which can be classified as rolled up graphite sheets and are the strongest material known to man. Nonetheless, these low numbers are due to the weak Van der Waals forces between the carbon sheets which are easily overcome.
Likewise, graphite is a brittle substance and this is could be due to crack propagation. The crack propagation theory suggests that as soon as one van der Waals force is overcome [at the surface] the succeeding intermolecular forces will be more easily broken; therefore, a fracture will from between two layers of carbon and the substance will split. This theory would suggest that iodine is also brittle but iodine does not have a giant covalent lattice structure so the two substances are difficult to compare in this way. Another theory is that the covalent bonds along the layers of carbon are directional and so stress across a single layer could cause a fracture easily. Nonetheless, personally, I have doubts about this theory because the bonds along the carbon layer are covalent, extremely strong and require a lot of energy to overcome
Additionally, graphite is a conductor of both electricity and heat due to each carbon atom having a delocalised electron which is free to move throughout the structure and carry charge to create a current. When the carbon atoms bond to the three surrounding carbon atoms, there is one remaining electron in the ‘p’ orbital of each atom. All of the ‘p’ orbitals overlap creating an equal electron density across all of the layers of carbon: this is called sp2 hybridisation. This is why the electrons are called delocalised.
Graphite’s density is 2266kg/m3 which is less than half of iodine’s density, which is 4930 kg/m3. This is because of the differing bond lengths and atomic masses. The covalent bond within an iodine molecule is 267 picometres [pm] and the Van der Waals distance is 430pm; whereas, graphite has 140pm covalent bonds between carbon atoms and 340 pm Van der Waals distance between the carbon layers. This means that there would be more carbon atoms in a volume than there would be iodine atoms. Nevertheless, carbon atoms have an atomic mass of only 12.01 and iodine has the atomic mass of 126.90. Therefore, due to density =mass/volume and iodine having a higher mass of atoms [even with greater distance between particles] iodine is more dense.
Similar to iodine, graphite is insoluble in water. For graphite there would only be weak Van der Waal’s attractions between the carbon atoms and the water molecules in solution; whereas, the carbon atoms are covalently bonded very strongly to one another in plain graphite, which is more stable and therefore preferred by the substances. Exactly like iodine, graphite is non-polar and so will not dissolve in polar water because of the general rule.
Equally, graphite has much higher melting [3823K] and boiling [4098K] points than iodine but, alike iodine, graphite sublimes at atmospheric pressures. Graphite has strong covalent bonds between carbons in the sheets of graphene and despite the weak Van der Waals forces between layers it requires a lot of energy to melt graphite as not only must you overcome the intermolecular forces but the intramolecular forces as well. The carbon phase diagram [figure 4] below explains the conditions required to melt graphite as at 1atm graphite would sublime like iodine instead of evaporate.
Figure 4: the carbon phase diagram
Dihydrogen monoxide, or water, is a polar covalent molecule. Oxygen is much more electronegative than hydrogen [3.5 compared to 2.1] and, as a result, there is dipole-dipole interaction. This is called a hydrogen bond and is much stronger than regular dipole-dipole forces because oxygen atoms in dihydrogen monoxide have spare lone pairs of electrons giving the molecule a bent shape with bond angles of 104.5˚. Also, in water, the hydrogen atoms are electron deprived because as the oxygen is much more electronegative and pulls the shared electrons involved in the covalent bonding towards it, leaving the hydrogen atoms positively charged. What is more, because of the hydrogen atoms small size the protons in the nucleus are left exposed and increase the atoms positive charge. Figure 5 illustrates the hydrogen bond with a dot and cross diagram exhibiting the lone pair of electrons of the oxygen and a stick diagram showing the forces of attraction.
Figure 5: a stick diagram showing the hydrogen bonds and a dot and cross diagram showing the lone pair of electrons of the oxygen
The low melting [273.15K] and boiling [373.15K] points of ice are a consequence of the weak intermolecular forces holding the covalent molecules together. Unlike graphite and similar to iodine only the weak intermolecular forces [Van der Waals in iodine and graphite and hydrogen bonds in ice] need to be overcome in order to melt ice.To continue, solid dihydrogen monoxide [ice] has a low melting and boiling point, is brittle and hard, is a poor conductor of heat and electricity, has a very low tensile strength and most interesting has a low density.
Furthermore, ice is brittle but hard due to the arrangement of the water molecules. Dihydrogen monoxide, can from many structures when solid but the most common is when each oxygen atom is surrounded by four hydrogen atoms, two of which are covalently bonded to the oxygen atom, and the two others [farther away] are hydrogen bonded to the oxygen’s spare electron pairs [this form of ice is known as ice-Ih]. The hydrogen bonds create to a very ordered structure as shown in figure 6. This hexagonal structure which interconnects all of the water molecules with hydrogen bonds is hard because every atom is bonded to another in some way and together they are hard and strong. This also makes ice brittle. Ice is a big structure where everything is strongly bonded to everything else, and nothing can move or flow. If you give this robust structure a burst of energy by hammering it, the only way the energy can be dealt with is by breaking bonds [an endothermic process]; so, the structure will break apart quickly, which is called being brittle. Crack propagation also provides a reason for why ice is brittle, as explained with graphite. Although, ice is described as hard it only has a tensile strength between 0.7 and 3.1 MPa. This is most likely because ice is only hard in relation to other simple molecular substances.
Figure 6: the structure of ice
With regards to solubility, as explained earlier with iodine, because ice is a polar substance it will dissolve in a polar solvent, such as ethanol.Ice is a very poor conductor of electricity and heat. Although it is thought that water can conduct electricity, it cannot. The impurities in water [ions of Mg and Ca, for example] are what move and carry charge to create a current, which is what leads many people to believe water is a conductor. This means de-ionised or distilled pure water cannot conduct electricity. Likewise, even impure ice would still not conduct electricity as the ions within the structure of the ice would not be able to move to carry charge and create a current.
One interesting thing about ice is that it is less dense than its liquid form. Also, 934kg/m3 is very low compared to graphite’s density of 2266kg/m3 and iodine’s density of 4930kg/m3 and this is due to the structure of ice. The three dimensional hexagonal arrangement creates an open structure with many empty spaces, creating a low density. Compared to water, ice’s molecules are spread over a larger area, hence why water has a higher density of 999.8kg/m3. This is simply demonstrated in figure 7.
Figure 7: comparing the structure of ice and water
Sodium chloride has a high melting and boiling point, it is hard and brittle, soluble in water and conducts electricity but only when molten or in solution. All of these properties transpire because of the structure of this ionic substance as explained above.Lastly, sodium chloride is an ionic substance which exists in a giant ionic lattice with electrostatic attraction between positively charged sodium ions and negatively charged chloride ions. Within the lattice each sodium ion is encircled by six chloride ions, and each chloride ion is surrounded by six sodium ions. This is an octahedral arrangement.
Firstly, the high melting point of 1074.15K and the high boiling point of 1686.15K is due to the strong forces of electrostatic attraction between the positively charged sodium ions and negatively charged chloride ions. These ionic bonds require a large amount of energy to overcome; however, covalent bonds are stronger than ionic bonds so graphite still has higher melting and boiling points. Conversely, ionic bonds are significantly stronger than hydrogen bonds and Van der Waals forces, which explains why sodium chloride has higher melting and boiling points than both water and iodine.
Additionally, the strong ionic bonds also explain why sodium chloride is hard [because it is difficult to break the strong forces of electrostatic attraction] and from this theory it can be hypothesized that sodium chloride has a high tensile strength. Continuing, the crystalline lattice structure of this NaCl explains why it is brittle. When the crystal is hit, the stress applied can shift the ion layers slightly resulting in ions of the same charge being next to each and so the like charges repel and the crystal shatters into pieces.
In sight of solubility, Sodium chloride is soluble in water and any other polar solvent. This is because when a solute dissolves in a solvent the ions of the solute separate out and move between the spaces of the solvent [this is endothermic as it is bond breaking of both the ionic bonds and the hydrogen bonds]. The solvent molecules mix with the solute ions and intermolecular forces of attraction between solute and solvent particles form and move the solute particles in the spaces and hold them there [this is exothermic as it is bond making of permanent dipole-dipole forces]. In the case of sodium chloride, the δ+ ends of the water molecules surround the Cl− ions, and the δ− ends of the water surround the Na+ ions. Overall, this is an exothermic reaction, which means the end product will be lower energy, more stable and, therefore, favoured by the particles (Z., 2016). Figure 8 below shows the arrangement of sodium and chloride ions in solution.
Figure 8: the orientation of sodium and chloride ions and water molecules in aqueous sodium chloride
As previously mentioned, sodium chloride is a conductor when molten or in solution. This is because the ions within the ionic lattice are unable to move when sodium chloride is solid. However, when aqueous the ions can move through the water and carry charge to create a current, which is the reason why it is often thought that water can conduct electricity as stated before. Also, when molten the ions become mobile as the ionic bonds are relaxed and allow movement, which enables the ions to carry charge and create a current. Graphite is a much more useful conductor of electricity as is can conduct in its solid state.
Sodium chloride has a density of 2160kg/m3, which is higher than water but slightly lower than graphite and significantly lower than iodine. The bond distance in NaCl is 236pm, which is smaller than the overall bond lengths of graphite and iodine independently. However iodine has a much higher atomic mass than sodium chloride [58.44], which explains those differing densities.
To conclude, the varying strength of the intramolecular bonds covalent and ionic cause differing properties of substances. Also, the diverse strength and size of intermolecular forces account for the many different densities, solubilities and melting and boiling points. Principally, it is the arrangement of electrons and the order of ions within the substance that determines most properties, such as tensile strength and conductivity.
Learning in Parallel 