Bonding

BONDING
Atoms which do not have a complete outer (valence shell) are not stable. Atoms may gain, lose or share electrons and in this way the ions or atoms become linked by **ionic** or **covalent bonds.**

 Ions are formed when **metal atoms** (which lose electrons) react with **non-metal atoms** (which gain electrons). The oppositely charged positive and negative ions are held together by electrostatic attraction to form a 3-dimensional lattice of ions. We say that the ions present in such an **ionic compound** are held together by **ionic bonds.** For example in aluminium oxide the Al3+ and O2 -  ions are held together by ionic bonds and, since the formula is Al2O3 the ratio of Al3+ ions to O2 -  ions in the crystal is 2 : 3.  NOTE: Ionic compounds do not exist as discrete molecules so Al2O3 represents the empirical formula (rather than molecular formula) and shows the ratio of positive and negative ions in the crystal lattice.
 * **Ionic Bonds **

When atoms share electrons they form **molecules** in which the atoms are held together by **covalent bonds**. These molecules only occur when **non-metal atoms** share electrons with other **non-metal atoms** e.g. CO2**,** H2SO4. The shared electrons are called bonding electrons. The formation of a bond (ionic or covalent) is an exothermic process as it leads to a release of energy. Conversely bond-breaking is an endothermic process i.e. energy has to be put in to break the nett attraction between the two atoms (or ions).
 * **Covalent Bonds **

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The atoms linked in a molecule can be shown by drawing a Lewis diagram (or dot diagram), which only shows the valence electrons. e.g. the Cl2 molecule is formed by linking the two Cl atoms, each with their 7 valence electrons. Each Cl atom will need to share 1 extra electron to form a covalent bond.
 * Lewis Diagrams **



Before starting to draw Lewis diagrams, it is essential to count the total number of **valence** electrons. The central atom is identified from symmetry or because it is the least electronegative atom. Then start by linking each atom into the correct structure using a single covalent bond. This bond is equivalent to two electrons i.e. a bonding electron pair. The remaining valence electrons are placed around the outer atoms and then the central atom until all valence electrons are used.

 Other examples of Lewis diagrams are:

 H2O (a total of 8 valence electrons)

 CCl <span style="font-family: Arial,sans-serif;">4 <span style="font-family: Arial,sans-serif; font-size: 11pt;">(a total of 34 valence electrons)

<span style="font-family: Arial,sans-serif; font-size: 11pt;">In some molecules the sharing of one pair of electrons will not result in a complete valence shell (or octet of electrons) for all atoms. It may be necessary to share two pairs of electrons (represented by a double line i.e. =) between 2 atoms. This results in a double covalent bond. Atoms such as oxygen and sulfur which have 6 electrons in their valence shell, may form a double covalent bond when linked to only one other atom. <span style="font-family: Arial,sans-serif; font-size: 11pt;"> e.g. oxygen gas, O2

<span style="font-family: Arial,sans-serif; font-size: 11pt;">Similarly, a nitrogen atom with 5 electrons in the valence shell needs 3 more electrons. If joined to three other atoms each bond will be a single covalent bond but if the atom is linked to only one other atom it will form a triple covalent bond (sharing of 3 electron pairs).
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">Note: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> If the O atom is linked to 2 other atoms (as in H2O) then each bond will only be a single bond.

<span style="font-family: Arial,sans-serif; font-size: 11pt;"> nitrogen gas, N2
 * <span style="font-family: Arial,sans-serif; font-size: 11pt;">NOTE: **<span style="font-family: Arial,sans-serif; font-size: 11pt;"> To draw any Lewis diagram always start by linking the atoms so that they are all connected by a single bond (in the appropriate arrangement). If there are not enough electrons to complete the octet for each atom (except H) then non-bonding pairs will need to be shifted to form a double covalent bond (or triple bond).

//<span style="font-family: Arial,sans-serif; font-size: 11pt;">Exercise: Draw Lewis diagrams of the following molecules. // //<span style="font-family: Arial,sans-serif; font-size: 11pt;">CH4, NH3, H2S, H2CO, PCl3, CS2, CH3OH //