Bonding: Ionic, Covalent and Metallic
Why Atoms Form Bonds
Atoms are most stable when their outer electron shell is full. Most atoms do not naturally have a full outer shell, so they bond with other atoms to achieve one. There are three main types of chemical bond: ionic bonding, covalent bonding and metallic bonding. All chemical bonds involve electrostatic forces — the attraction between opposite charges. The type of bond that forms depends on whether the bonding atoms are metals or non-metals.
Ionic Bonding
Ionic bonding occurs between a metal and a non-metal. The metal atom transfers one or more electrons to the non-metal atom. The metal loses electrons and becomes a positively charged ion (cation). The non-metal gains electrons and becomes a negatively charged ion (anion). The oppositely charged ions attract each other strongly — this electrostatic attraction is the ionic bond.
For example, in sodium chloride (NaCl): sodium (electronic configuration 2, 8, 1) loses one electron to become \(\text{Na}^+\) (2, 8). Chlorine (2, 8, 7) gains one electron to become \(\text{Cl}^-\) (2, 8, 8). Both now have full outer shells. In magnesium oxide (MgO): magnesium (2, 8, 2) loses two electrons to become \(\text{Mg}^{2+}\), and oxygen (2, 6) gains two electrons to become \(\text{O}^{2-}\). In an ionic compound, vast numbers of ions pack together in a regular three-dimensional giant ionic lattice, where each ion is surrounded by ions of opposite charge.
In ionic bonding, electrons are transferred from metal to non-metal. The ions formed have opposite charges and attract each other. The compound forms a giant ionic lattice. You can use the charges of the ions to work out the formula — the charges must balance to zero overall.
Covalent Bonding
Covalent bonding occurs between non-metal atoms. Instead of transferring electrons, atoms share one or more pairs of electrons. Each shared pair of electrons forms one covalent bond. By sharing, both atoms achieve full outer shells. A single bond involves one shared pair; a double bond involves two shared pairs (e.g. \(O_2\), \(CO_2\)); a triple bond involves three shared pairs (e.g. \(N_2\)).
Examples: in water (\(H_2O\)), the oxygen atom forms a single covalent bond with each of two hydrogen atoms. In chlorine gas (\(Cl_2\)), two chlorine atoms share one pair of electrons. In carbon dioxide (\(CO_2\)), the carbon forms a double bond with each oxygen atom. Covalent compounds range from small simple molecules (\(H_2O\), \(NH_3\), \(CH_4\)) to huge giant covalent structures such as diamond and graphite.
Dot-and-cross diagrams show the outer-shell electrons of each atom involved in bonding. Use dots (•) for electrons from one atom and crosses (×) for electrons from the other. For ionic bonding: draw the ions in separate sets of square brackets, show each ion’s charge, and show where the transferred electrons now sit. For covalent bonding: overlap the outer shells and show shared pairs in the overlap region. Common examples to practise: \(H_2\), \(Cl_2\), \(HCl\), \(H_2O\), \(NH_3\), \(CO_2\), NaCl, \(MgO\). These diagrams are models — they show electron arrangements but do not show the 3D shape, relative atom sizes, or bond lengths.
A dative covalent bond (also called a coordinate bond) is a special type of covalent bond in which both electrons in the shared pair come from the same atom. This occurs when one atom has a lone pair of electrons it can donate to another. An example is the ammonium ion (\(NH_4^+\)) formed when ammonia (\(NH_3\)) donates a lone pair of electrons to a hydrogen ion (\(H^+\)). The resulting bond is identical to any other covalent bond once formed; the distinction is only in how it was formed. Dative covalent bonding is not on the OCR Gateway course — treat it as background reading.
Metallic Bonding
In metals, atoms are arranged in a regular lattice structure. The outer-shell electrons leave their parent atoms and become delocalised — they are free to move throughout the whole structure. This creates a lattice of positive metal ions surrounded by a “sea” of delocalised electrons. The strong electrostatic attraction between the positive ions and the negative delocalised electrons holds the metal together. This is metallic bonding.
Metals conduct electricity because the delocalised electrons are free to move and carry charge through the structure. The same delocalised electrons allow metals to conduct heat. Metals are malleable (layers of positive ions can slide without breaking the bond, because the electron sea adjusts).
Predicting Bond Type from the Periodic Table
You can use an element’s position in the periodic table to predict what type of bond it will form. Metals are on the left and tend to lose electrons; non-metals are on the right and tend to gain or share electrons. A metal and a non-metal together → ionic bond. Two or more non-metals → covalent bond. A pure metal or alloy → metallic bond. Group number gives the outer electron count, and period number gives the number of shells — both are important in working out which atoms will bond and in what ratio.
The overall charge in an ionic compound must be zero. Group 1 metals form +1 ions; Group 2 metals form +2 ions. Group 6 non-metals form −2 ions; Group 7 non-metals form −1 ions. To work out the formula: find the lowest number of each ion that makes the charges cancel. For example, calcium (\(Ca^{2+}\)) and chloride (\(Cl^-\)): you need two \(Cl^-\) ions to balance one \(Ca^{2+}\), giving \(CaCl_2\). Aluminium (\(Al^{3+}\)) and oxide (\(O^{2-}\)): need 2 \(Al^{3+}\) and 3 \(O^{2-}\) for a total charge of zero, giving \(Al_2O_3\).