Dot & Cross Diagrams

Representing Ionic Bonds

• Step 1: Identify the metal and non-metal. The metal will lose electrons to become a cation, and the non-metal will gain them to become an anion.
• Step 2: Draw the ions in square brackets. This indicates that the charge is distributed over the entire ion.
• Step 3: Place the charge (e.g., $+, 2-, 3+$) at the top right-hand corner outside the bracket. Ensure the metal ion is shown with an empty outer shell (or the new full shell) and the non-metal with a full octet containing both its original electrons and the ones it gained.

Representing Covalent Bonds

• Step 1: Draw the atoms with their shells overlapping. The shared electrons must be placed within the intersection of the circles.
• Step 2: Ensure each shared pair represents one single covalent bond. Double bonds require four electrons (two pairs) in the overlap, and triple bonds require six electrons (three pairs).
• Step 3: Include lone pairs (non-bonding electrons) on the outer shells of the atoms to ensure the octet rule is satisfied where applicable.

• Check the Group Number: Always use the Periodic Table to verify the number of valence electrons. For example, Group 16 elements must have 6 valence electrons in their neutral state.

• Count the Total Electrons: After drawing, count the electrons for each atom. In covalent molecules, the shared electrons count toward the total for both atoms involved in the bond.

• Bracket Placement: In ionic diagrams, never forget the square brackets. Omitting them or placing the charge inside the circle is a frequent cause of lost marks.

• Hydrogen Exception: Remember that Hydrogen only needs two electrons to fill its shell ($1s^2$), not eight.

• Inner Shells: A common mistake is drawing all electron shells. Unless specifically asked, only draw the outermost shell to keep the diagram clear and accurate.

• Octet Expansion: Students often struggle with Period 3 elements (like Sulfur or Phosphorus) which can 'expand the octet' and hold more than 8 electrons. Do not force an octet if the chemical formula implies more bonds (e.g., $SF_6$).

• Electron Deficiency: Some atoms, like Boron or Aluminum, can be stable with fewer than 8 electrons in certain molecules. These are known as electron-deficient species.