Lewis Electron Accounting for Ions

Introduction

The Lewis structure of an ion requires an electron budget adjusted for charge. Add one electron for a -1 ion, or subtract one for a +1 ion, relative to the neutral-atom total. A correct diagram uses that budget and shows the final net charge in brackets.

Core explanation

For NH₄⁺, nitrogen contributes five valence electrons and four hydrogens contribute four, giving nine from neutral atoms. The +1 charge means one fewer electron, so the ion has eight valence electrons. Four N–H single bonds use all eight; nitrogen has no lone pair in this Lewis picture. Put the structure in brackets with + outside. Four bonds around N give formal charge +1 on N and zero on each H, matching the overall charge.

For OH⁻, oxygen contributes six and hydrogen one, then the -1 charge adds one: total eight. One O–H bond uses two electrons, leaving six as three lone pairs on O. Oxygen's formal charge is 6 - 6 nonbonding - 1 assigned from the bond = -1. Brackets carry the - charge. The ion is not a neutral OH radical; charge changes the electron count and diagram.

For NO₃⁻, nitrogen contributes five, three oxygen atoms contribute eighteen, and the negative charge adds one, making twenty-four valence electrons. A valid common Lewis contributor has N central, one N=O and two N–O single bonds, with lone pairs filling O octets. Formal charges are +1 on N, -1 on each single-bonded O and 0 on the double-bonded O, totaling -1. There are three equivalent resonance contributors depending on which O is double-bonded. The real ion is not switching among three fixed structures.

The electron budget is necessary but not sufficient. One can place twenty-four electrons in an incorrect connectivity or violate octet expectations. Check the likely skeleton, bond count, each atom's electron environment and formal charges. Hydrogen is terminal and follows a duet, not an octet. Some central atoms can be electron deficient or hypervalent in useful Lewis representations, so applying octet as an absolute rule can mislead.

When adding electrons for an anion, do not assign the extra electron to a particular atom before drawing a plausible structure. The net charge belongs to the ion; formal charges in a particular Lewis contributor are bookkeeping assignments. Resonance can delocalise electron density beyond the single drawn placement.

Step-by-step reasoning

1. Sum neutral-atom valence electrons. 2. Add electrons equal to negative charge or subtract for positive charge. 3. Choose connectivity and allocate bonds from the budget. 4. Fill terminal lone pairs, then complete the central atom where appropriate. 5. Recount electrons, formal charges and the bracketed net charge.

Visual explanation

Make three electron ledgers: NH₄⁺ as 5 + 4 - 1 = 8; OH⁻ as 6 + 1 + 1 = 8; NO₃⁻ as 5 + 18 + 1 = 24. Beside each, show bonds and lone-pair dots summing to the ledger total.

Real-world analogy

A budget tells how many coins can be distributed, but not which purchase is best. The valence-electron total limits Lewis structures; connectivity, octets and formal charge decide how to allocate those electrons sensibly.

Real-world example

Polyatomic ions such as ammonium and nitrate appear in ionic salts. Their internal atoms are covalently connected, while the whole ion interacts electrostatically with counterions in a crystal. Lewis electron accounting explains the internal ion structure without turning the salt into isolated neutral molecules.

Why?

Why subtract an electron for a cation? A +1 ion has one fewer electron than its neutral-atom sum while retaining the same nuclei. The adjusted budget is a direct expression of its charge.

Common misconception

“A -1 ion means every atom has one extra electron.” Only the total species has one extra electron relative to neutral atoms. Its distribution can be shared, and formal charges depend on the chosen Lewis contributor.

Worked example

Draw CO₃²⁻. Carbon contributes four, three O contribute eighteen, and 2- adds two: twenty-four valence electrons. Put C in the centre with three O around it. A common contributor has one C=O and two C–O⁻ bonds. Carbon formal charge is 0, double-bonded O is 0, and each singly bonded O is -1; total -2. Three equivalent contributors place C=O at any oxygen, supporting equivalent C–O bonds in the resonance hybrid.

Quick check

1. How many valence electrons does NH₄⁺ have in its Lewis budget? Answer: Eight: 5 from N plus 4 from H minus 1 for positive charge.

Exam focus

Show the charge adjustment explicitly and place brackets around ions. Recount all bonds and lone pairs. Sum formal charges to the net charge, but do not mistake formal-charge locations for measured charge density.

Advanced insight

Resonance structures share a fixed nuclear arrangement and electron total. Electron density can be distributed over several atoms, so no one formal-charge diagram is a full picture of the ion's quantum state.

Summary

Lewis accounting for ions starts from neutral-atom valence totals, adds for negative charge and subtracts for positive charge. A sound structure then uses all electrons, respects plausible bonding and has formal charges summing to the bracketed ion charge.

Practice questions

1. How many valence electrons are in OH⁻? Answer: Eight. 2. How many are in NO₃⁻? Answer: Twenty-four. 3. What charge appears outside brackets for carbonate? Answer: 2-, because CO₃²⁻ is a dianion. 4. Why can nitrate have three Lewis contributors? Answer: The double bond can be assigned to any one of three equivalent oxygens while keeping the same atom connectivity and electron count.