Compare and contrast the structure and bonding in ionic, covalent and metallic species. Relate the structure and bonding to the properties of the species. Ionic, covalent and metallic bonding are all various types of bonding found between atoms, but there are big differences between how they work, and what atoms they bond. Ionic bonding is when a metal and a non-metal are stuck together by an electrostatic bond. An ion is formed when electrons are transferred from one atom to another.
As electrons have a negative charge, the atom that loses the electrons becomes positive, and the atom that gains electrons becomes negative.
Atoms do this so that they can have a full outer shell of electrons. Electrostatic bonds then hold these positive and negative ions together very strongly. When atoms are bonded like this it is called ionic bonding – bonding of ions. An example of an ionic bond is sodium chloride. Sodium only has one electron in its outer shell and chlorine has 7. The sodium atom loses and electron to become Na+ and the chlorine atoms gains this electron and becomes Cl-.
Electrostatic bonds then hold these oppositely charged ions together. Ionic crystals are giant lattices of ions.
The structure is called giant because it is made up of the same basic unit repeated over and over. The structure of ionic compounds decides their physical properties. Ionic compounds conduct electricity when they are molten or dissolved, but not when solid. This is because ions in liquid or aqueous form are free to move around and can therefore conduct electricity.
In a solid the ions are fixed in place by the strong bonds. Ionic compounds also have high melting and boiling points. The giant ionic lattices are held together by very strong electrostatic forces and so a large amount of energy is required to overcome these forces, which leads to very high melting and boiling points. Sodium chloride is an example of an ionic compound, and its melting point is 801oC. Ionic compounds do however often dissolve in water. This is because water molecules are polar. This basically means that due to electronegativity the oxygen atom in water is slightly negative and the hydrogen atoms are slightly positive. These polar water molecules pull the ions away from the lattice and cause them to dissolve.
Another type of bonding is called covalent bonding. This is where atoms share electrons with each other to gain a full outer shell or to the nearest noble gas configuration. These bonds are formed when orbitals each containing one electron, overlap. This forms a region in space where an electron pair can be formed and new molecular orbitals are formed. The greater the overlap, the stronger the bond. Covalent bonds can bond atoms that are different or the same and are formed between non-metals only. For example, two hydrogen atoms each shares one electron so that each atom feels like it has a full outer shell. One type of covalent bonding is dative covalent bonding. This is where one atom donates both electrons to a bond. Unlike ionic bonds, covalent bonds have very low melting and boiling points because instantaneous dipole-induced dipole (ID-ID) forces are weak; they increase as molecules get a larger surface area.
As the ID-ID forces are weak, little energy is required to separate molecules from each other, which leads to low melting and boiling points. However, some melting and boiling points are higher than expected for a given mass because you can get additional forces of attraction. Covalently bonded molecules are usually insoluble in water as the polar water molecules are more attracted to each other than the molecular substance. They also do not conduct electricity as there are no charges involved that are free to move. Giant covalent structures however have very different properties. Giant molecular structures have a huge network of covalently bonded atoms and are sometimes called macromolecular structure. Diamond (a form of carbon) and silicon are two examples or giant covalent structures. The reason carbon and silicon atoms can form these structures is because they can each form four strong covalent bonds.
In diamond for example, each carbon atom is covalently bonded to four other carbon atoms and the atoms arrange themselves in tetrahedral shapes to for a crystal lattice structure. Because of this, giant covalent structures are extremely hard and strong. Diamond is often used as tips for drills and saws. Another quality is that vibrations travel easily through the still lattice and so make it a good conductor. These structures also have very VERY high melting and boiling points. Diamond actually sublimes (changes straight from a solid to a gas) at over 3800 K. They can’t conduct electricity because all the outer electrons are held in covalent bonds with no charge and won’t dissolve in any solvent. This is because atoms are more attracted to their neighbours in the lattice than to solvent molecules.
Metallic bonding is another type of bonding that occurs between metals. Metals cannot achieve the stable electron arrangement of a noble gas through ionic or covalent bonding because they don’t have enough outer shell electrons to do so. Metal elements exist as giant metallic lattice structures. In metallic lattices, the electrons in the outermost shell of the metal atoms are delocalises – they are free to move. This leaves positive metal ions which are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons. This is metallic bonding. The melting points of metals are generally high because of the strong metallic bonding, with the number of delocalised electrons per atom affecting the melting point. The more there are the stronger the bonding will be and therefore the higher the melting and boiling points.
The size of the metal ions and the lattice structure also affect the melting point. As there are no bonds holding specific ions together, the metal ions can slide over each other when the structure is pulled, so metals can be shaped and are ductile. Metals are also good conductors as the delocalised electrons can pass kinetic energy to each other through vibrations. However, as vibrations increase with temperature is can cause the electrons to no longer move freely though the metal lattice and this is why conductivity of a metal drops with increasing temperature. Metals are also good conductors because the delocalised electrons can carry a current. Finally metals are also insoluble, except in liquid metals, because of the strength of the metallic bonds.
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