There are three forms of substances exist: liquid, solid and gas in which solid is one of the major state of matter. The solid structure bonds the atoms together by different chemical physical attraction. Chemical bonds are the result of interactions of electrons by various forces of attraction. This attraction can hold atoms together in a stable arrangement. Atoms may transfer or share atoms to form molecules and compounds. When atoms bond together by chemical bonds, they will become more stable. Different types of chemical bonds determine solids properties, such as melting point, conductivity and solubility (Lister and Renshaw, 2000).
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Electron transferring form ionic bonds while electron sharing and joining molecules form covalent bonds. Linus Pauling came up with a scale, a value for each element called the electronegativity (E.N.) value. Each element has different desire wanting, ability to attract electrons. The strong electron attractions were given high values and some atoms have very low ability to attract were given a low value. These values are relative-comparison values and have no units. The value of difference electronegativity between two atoms less than 1.8 are defined as covalent bonds while the value of difference electronegativity between two atoms more than 2.0 are defined as ionic bonds.
This essay will describe four types of solid structures: ionic and metallic which contains of unit cell, giant covalent which is held by network and simple molecular which are small molecules with weak forces of attraction.
Ionic Structure
First of all, ionic bonding commonly exists in crystal solid structures. Ionic bonding electrons are transferred from mental atoms to non-metal atoms which result in each ion obtaining a full outer shell of electrons to become stable. Commonly, metals form cation by losing valence electrons while non-metals form anions by gaining valence electrons. Sodium chloride (NaCl) is a well known ionic compound. Sodium loses one electron from its outer shell while the chlorine gains one electron to fill its outer most shell. When sodium (Na+) ions bond with chloride (Cl-) ions they form common table salt, sodium chloride (NaCl) (Lister and Renshaw, 2000).
In addition, Sodium chloride (NaCl) is made up of giant lattice of ions. There are a large amount of sodium ions and chloride ions packed together which depends on how big the crystal is. Figure 1 (Adapted Steinberg, 2000) shows how does a bit of sodium chloride lattice arranged.
Figure 1: Ionic Bonding in Sodium Chloride (Adapted from Steinberg, 2000)
Each sodium ion is at a centre surrounded by 6 chloride anions. Each chloride ion is also in the centre, it is also surrounded by 6 sodium cations. So sodium chloride is described 6 co-ordinated. The pattern in this way will be repeated countless times in sodium chloride crystal and ensure the maximum stability in sodium chloride. Because when each ion is touched by 6 opposite charged ions, there is more attraction between the ions which makes the structure more stable (Clark, 2010). Figure 2 (Adapted from Clark, 2010) shows clearly the unit cell of sodium chloride.
Figure 2: Unit Cell of Sodium Chloride (Adapted from Clark, 2010)
Commonly, the atoms arrange in a regular way, but sometimes this is not the case. All metal atoms consist of a lot of ‘crystal grains’ which are regions of regularity. At grains boundaries atoms become inconformity.
There are various properties in ionic compounds. Firstly, ionic compounds have a high melting point and boiling point. There are strong attractions between the positive and negative ions which take a lot of energy to overcome them. Secondly, one of the main properties of ionic compounds is they conduct electricity when molten, because when the compound is in the liquid state, the ions can carry the charge freely. Ionic compounds are brittle which resulting from an applied stress. The ions will be moved sufficiently to make contact between ions. Ions of the same charged are brought side-by-side leading to repulsion forces within the crystal. Many ionic compounds dissolve in the water. Water molecules have unbonded electrons, called lone pairs. They attract positive ions and negative ions in the compounds form dative bonds and polar water forms electrostatic attractions between the ions. Water molecules also produce energy by hydration to break up the lattice and reduce their attraction (Clark, 2010).
Covalent bonds
Atoms sharing electrons form covalent bonds which the electronegativity difference between two atoms are less than 1.8.
Giant Covalent Structure
In giant covalent structures, all atoms are bonded together by covalent bonds which directly have an influence on specific atoms in a regular extended network. The electrostatics forces hold other structures together to act equally in all directions. Diamond is best example of giant covalent structure (Lister and Renshaw, 2000).
Diamond has a tetrahedral shape in three-dimensions, with four covalent bonds from each carbon atom. According to VSEPR theory which can explain the levels of repulsion between chemical bonds (Lane, 2009), to reduce the repulsion four pairs of covalent bonds form and covalent bonds repel each other equally. There are no intermolecular forces of attraction only, between the carbon atoms.
Figure 4 (Adapted from Chemcases, 2010) shows the structure of diamond
Figure 4 shows the structure of diamond (Adapted from Chemcases, 2010).
Diamond is well known as the hardness natural substance, and results from four strong covalent bonds giving diamond a strong rigid extended and continuous structure, so diamond just can be cut by other diamond. In addition, diamond has a very high melting point of around 4000K, because to break up the four strong covalent bonds requires a lot of energy. Next, the electrons which are held tightly are not free to move, resulting in the diamond having neither conduct electricity nor heat. At Last, diamond is insoluble in water because the covalent bonds are much stronger than the attraction of water molecules (Clark, 2000).
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Simple Molecular Structure
The simple molecular substances are non-metal compounds which are made up of atoms by strong covalently bond and relatively weak intermolecular forces. Water is taken as a typical type of the simple molecular structure (Lower, 2010). In water, each hydrogen atom is bound to the central oxygen atom by a pair of sharing electrons. Oxygen has four free electrons in its outer second level to form two lone pairs to reduce the repulsions between clouds of negative charge, leaving two of the outer electrons shared with hydrogen to form covalent bonds. This rule leads to a tetrahedral structure in which the angle between electron pairs is 104.5°. In addition, water molecules forms hydrogen bonding.
Oxygen likes electrons more than hydrogen and electrons spend more time near the oxygen, resulting in oxygen is part negative charge while the hydrogen is positive charge. Electrons are distributed leading to water form a negative structure of hydrogen bonds. Dipole-dipole attractions occur between two water molecules, due to its polar nature (Lane, 2009). Figure 5 (Adapted from Google, 2010) shows the structure of water molecule.
Figure 5: Hydrogen Bonding Between Water Molecules (Adapted from Google, 2010)
As the weak forces exist in simple molecular substances, they are not very dense or strong nor solid resulting in it having a low melting point and boiling point. Simple molecular substances can easily become gas or liquid. There are no ions existing in these substances, so they are insoluble in the water and can not conduct electricity (Lower, 2010).
Metallic Structure
Metals are giant structures which hold the atoms together by metallic bonding transferring the electrons. All elements of metal can easily lose electrons forming positive ions which are in a freely moving ‘sea’ and electrons. How many electrons have been lost by each metal determine the number of electrons in the sea (Lister and Renshaw, 2000).
Most ions in metals pack as close as they can. However, sodium in Group one is more open and less dense and forms a unit called the body-centred cubic (BCC) structure which is a common packing geometry for some metals. This structure is not a close-packed arrangement, just 68% of the space being filled (Lister and Renshaw, 2000). There is an atom located in the centre of a cube which is surrounded by eight other spheres. Figure 3 (Adapted from Clark, 2007) shows the structure of sodium: a coordination of eight and a unit cell containing two atoms.
Figure 3: The Structure of Sodium (Adapted from Clark, 2007)
There are several properties about metals. Most of metals tend to have a high melting point and boiling point due to the strength of the metallic bonding. The strength of metallic bonding is not only different from metal to metal, but also depends on the number of electrons which each atom can delocalise in the sea of electrons and by the packing. Transition metals have a high melting point and boiling point because they form a unit called hexagonal close packing. Group one such as Sodium is an exception and has a low melting point and boiling point, because it only has one electron to contribute to the bond and it is 8 co-ordinated which can not form strong enough bonds as other metals. Secondly, metals are good conductors of electricity. The electrons in the ‘sea’ are free to move throughout the structure even cross the grain boundaries. The metallic bonding still exists as long as atoms are touching each other. In addition, these electrons are also responsible for the high thermal conductivities of metals. Electrons of the metals can pick up heat energy which is transferred to the rest of metal by moving electrons (Clark, 2007).
In conclusion, electrons are transferred forming ionic bonding in ionic crystal solid structure. As giant lattice of ions exist in ionic solids, the ionic solids not only have a high melting point and boiling point but also conduct electricity well. Metal hold the atoms together by metallic bonding which determines that metal compounds have a high melting point and boiling point and are good conductors of electricity and heat. Although both of giant covalent structures and simple molecular structures have electron sharing covalent bonding, they have different properties. Strong covalent network bonding is involved in giant covalent structures holding the atoms together, with the result that giant covalent compounds are easily to melt and can not conduct electricity. In contrast, there are some weak intramolecular forces in simple molecular structures which lead to this kind of structure having a low melting point and boiling point, and unable to conduct electricity.
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