Write the Lewis structure, including typical contributions to the resonance structure (where appropriate, allow for the possibility of octet expansion), for (a) hydrogen phosphate ion; (b) bromite ion; (c) perchlorate ion; (d) arsenate ion.

Valence electrons are the electrons that reside in the outermost shell of an atom and are available for bonding. As the 'social butterflies' of atoms, they can interact with valence electrons of other atoms to form chemical bonds.

For instance, while calculating the Lewis structure for the hydrogen phosphate ion \(HPO4^{2-}\), it’s essential to tally up these electrons. Phosphorus (P) brings five to the party, each Oxygen (O) arrives with six, and Hydrogen (H) only has one. Don't forget, since our ion is carrying a 'backpack' of -2 charge, we invite two extra electrons to ensure the balance. That's how we reach a total of 32 valence electrons needing to be arranged in the structure.

The Octet Rule is like the golden rule of atomic stability - atoms strive to have eight electrons in their valence shell, much like most people aim to have a full set of utensils to enjoy a good meal.

When sketching out the basic skeleton for the hydrogen phosphate ion, each Oxygen seeks to fill its electron 'plate' with eight. Since bonding with Phosphorus has already offered them a portion (two electrons), they snatch six more from the 'buffet' of lone pairs. Phosphorus, being the generous sort, may share electrons via double bonds to make sure everyone, including itself, reaches the 'dining' satisfaction of a complete octet.

Picture a dance floor where pairs of dancers (atoms) can switch partners (bonds). Resonance structures are different snapshots of this molecular dance, where electrons 'boogie' between positions without altering the actual arrangement of atoms.

In the case of \(HPO4^{2-}\), after we give everyone their rightful share of electrons, we notice that Phosphorus can still cut a rug with more than one Oxygen - hence, we sketch several possible dance moves (resonance structures). These are not distinct entities but rather a blend, a 'remix', if you will, that better represents the molecule's electron arrangement.

Formal charge is like balancing your checkbook - it helps you keep track of the 'electron economy' to ensure every atom gets its fair share. We calculate by taking the number of valence electrons an atom 'brought' minus the electrons it now 'owns' (half the bonded and all the nonbonded ones).

For our hydrogen phosphate ion friend, ensuring the correct formal charge is crucial for accurately depicting the molecule. With all atoms striving for the lowest possible formal charge, we may need to redistribute the bonding 'wealth', reflecting the ion's true charge, all while making sure our electron account doesn't dip into overdraft!