Chemical Symbols & Formulae

Seven specific elements naturally exist as diatomic molecules, meaning their atoms are always found in pairs when in their pure elemental form. These include hydrogen ($H_2$), nitrogen ($N_2$), oxygen ($O_2$), and the halogens fluorine ($F_2$), chlorine ($Cl_2$), bromine ($Br_2$), and iodine ($I_2$).

When writing chemical equations or representing these elements, they must always be written with a subscript of 2 unless they are part of a compound. This physical reality is critical for balancing chemical equations correctly.

These molecules are classified as simple molecules because they consist of a small, fixed number of atoms chemically joined together by covalent bonds.

Simple molecules are typically formed from non-metal atoms that share electrons to achieve stability. The formula reveals the exact count of each atom type involved in the individual molecule.

Visual representations often use space-filling or ball-and-stick models where different colors represent different elements. These diagrams help illustrate the spatial arrangement and connectivity of the atoms.

In these compounds, the atoms are held together by strong internal bonds, but the forces between separate molecules are relatively weak. This distinction is important for understanding physical properties like boiling points.

Ionic compounds form when metal atoms lose electrons to become positive ions (cations) and non-metal atoms gain electrons to become negative ions (anions). Unlike simple molecules, ionic compounds exist in a repeating lattice structure.

Polyatomic ions are groups of atoms that stay together as a single unit and carry an overall electrical charge. Common examples include the carbonate ion ($CO_3^{2-}$) and the sulfate ion ($SO_4^{2-}$).

The most fundamental rule for writing ionic formulae is that the total charge must be zero. The positive charges from the cations must exactly cancel out the negative charges from the anions to create a neutral compound.

The Direct Comparison method works best for ions with equal but opposite charges, such as $Mg^{2+}$ and $O^{2-}$. Since the charges cancel each other out in a 1:1 ratio, the resulting formula is simply the symbols combined without extra subscripts ($MgO$).

The Swap and Drop method is a reliable technique for ions with different numerical charges. The numerical value of the charge on the first ion becomes the subscript for the second element, and vice-versa, while ignoring the plus/minus signs.

When a subscript is needed for a polyatomic ion (to show more than one of that group), the entire ion formula must be placed inside brackets. For instance, in $Mg(OH)_2$, the brackets show there are two hydroxide groups for every one magnesium ion.

When naming compounds, the metal always comes first. The name of the non-metal is modified to end in '-ide' if it is a simple ion (e.g., oxygen becomes oxide, chlorine becomes chloride).

If the compound contains a polyatomic ion that includes oxygen, the name typically ends in '-ate' (e.g., sulfate, carbonate), unless it is the hydroxide ion ($OH^-$).

Always verify your final formula by multiplying the subscripts by the ion charges. The sum of (count of cation $\times$ cation charge) + (count of anion $\times$ anion charge) must equal exactly zero.