Mistake Master
Valence electrons and ionic compounds
Valence electrons decide how an atom bonds. When a metal hands electrons to a nonmetal, both become charged ions that snap together — and the formula that results is set by one rule: the charges must balance.
§1
Valence electrons decide the chemistry.
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Only the outermost electrons — the valence electrons — take part in ordinary bonding. Their count, which for a main-group element matches its group, sets what ions an atom tends to form.
Many main-group atoms gain or lose valence electrons to reach the electron count of the nearest noble gas — a stable, filled-shell arrangement. Metals on the left have few valence electrons and tend to lose them, becoming positive cations; nonmetals on the right are a few electrons short and tend to gain them, becoming negative anions. This "noble-gas-like" pattern is a reliable guide for the main group, but it is a trend, not a universal law — transition metals and some other ions settle on charges that do not leave a simple noble-gas configuration.
When a metal hands electrons to a nonmetal, the resulting cations and anions attract each other and lock into an ionic compound. What matters for the formula is not the noble-gas story but the charges themselves: the compound must come out electrically neutral.
§2
Building an ionic formula.
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A neutral compound has no leftover charge. That single requirement fixes the formula — you just balance the pluses against the minuses.
- Find each ion's charge. For main-group elements, the common charge follows the group: group 1 forms 1+, group 2 forms 2+, group 16 forms 2−, group 17 forms 1−, and so on.
- Combine so the charges cancel. Take enough of each ion that the total positive charge equals the total negative charge. The compound as a whole must be neutral.
- Read the counts as subscripts. The number of each ion needed becomes its subscript. A quick shortcut: the magnitude of one ion's charge becomes the other ion's subscript (the "criss-cross"), then reduce to the simplest ratio.
- Check neutrality and simplify. Confirm the charges sum to zero, and reduce subscripts to their smallest whole-number ratio (Ca²⁺ with O²⁻ is CaO, not Ca₂O₂).
Notice the formula comes entirely from charge balance, not from atoms "sharing" or "completing octets." The octet idea helps predict a main-group ion's charge; neutrality is what turns those charges into a formula.
§3
The pieces you'll meet.
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Quick reference card. Charge is the currency; neutrality is the rule.
§4
Worked example: the formula of aluminum oxide.
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Question. Aluminum (group 13) and oxygen (group 16) form an ionic compound. What is its formula?
Step 1 — find the charges. Aluminum loses 3 valence electrons to form Al³⁺. Oxygen gains 2 to form O²⁻.
Step 2 — balance the charges. The charges (+3 and −2) don't cancel one-to-one. The least common multiple is 6: you need two Al³⁺ (giving +6) and three O²⁻ (giving −6).
Step 3 — write the subscripts. Two aluminums and three oxygens gives Al2O3. (Shortcut: criss-cross the charge magnitudes — the 2 from oxygen and the 3 from aluminum — to the opposite subscripts.)
Step 4 — check neutrality. 2(+3) + 3(−2) = +6 − 6 = 0. Neutral, and already in the simplest ratio. ✓
Sanity check. Nothing about "octets" told us the subscripts — the +3 and −2 charges did. If you'd guessed AlO by pairing them one-to-one, the charge would be +1 left over, which no neutral compound can have.
§5
3 mistakes that cost real points.
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"Every ion ends up with a full noble-gas shell."
That's a solid guide for main-group ions — sodium to Na⁺, oxygen to O²⁻ — but it isn't universal. Transition metals commonly form more than one ion (iron as Fe²⁺ or Fe³⁺), and plenty of stable ions don't land on a tidy noble-gas configuration. Treating "reach an octet" as an absolute law leads to wrong charges outside the main group.
Fix. Use the noble-gas pattern to predict main-group ion charges. For transition metals, take the charge as given (often from a Roman numeral) rather than forcing an octet.
"Just write the two elements side by side."
Aluminum and oxygen is not "AlO." The formula is whatever makes the charges cancel, and +3 with −2 needs a 2:3 ratio — Al2O3. Pairing ions one-to-one only works when their charges are equal in size. Skipping the charge balance produces formulas that aren't neutral and can't exist.
Fix. Always balance the total + against the total −. Criss-cross the charge magnitudes to subscripts, then reduce to the simplest ratio.
"All ionic compounds are held together about equally."
The strength of an ionic lattice follows Coulomb's law: it grows with the product of the ion charges and shrinks with the distance between ions. Magnesium oxide (2+ with 2−) is bound far more strongly than sodium chloride (1+ with 1−), which is why MgO melts near 2850 °C and NaCl near 800 °C. Ignoring charge and ion size hides these large differences.
Fix. Compare lattices by charge first, then size: higher charges and smaller ions mean a stronger lattice and a higher melting point.
§6
Skill Check.
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Ten scenarios. Pick the chips that match your answer, then check. A scenario marks complete the first time every part is right. Progress saves on this device.