Valence electrons and ionic compounds — AP Chemistry
1. Identifying Valence Electrons and Lewis Dot Notation ★★☆☆☆ ⏱ 3 min
Valence electrons are electrons occupying the highest principal energy level $n$ of an atom, and are the electrons that participate in chemical bonding. For main group elements (s- and p-block, groups 1A-8A), only electrons in the outermost s and p sublevels count as valence; d and f electrons in lower energy levels are never counted, even if they appear after the outermost s orbital in condensed configurations. A quick shortcut: the number of valence electrons equals the element's group number for main group elements.
Lewis dot notation is a standard convention to represent valence electrons: draw the element's chemical symbol, then add one dot per valence electron, placing one dot on each of the four sides before pairing any dots. This makes it easy to see how many electrons an atom will lose or gain to form an ion.
Exam tip: If you are asked to count valence electrons for a main group ion, add one electron for each negative charge and subtract one for each positive charge, always starting from the neutral atom count.
2. Predicting Ionic Charges and the Octet Rule ★★☆☆☆ ⏱ 2 min
Metals (left of the metalloid staircase on the periodic table) have low ionization energy, so they lose all their valence electrons to form positively charged cations. For main group metals, the charge of the most stable cation equals the group number: 1A = +1, 2A = +2, 3A = +3. Nonmetals (right of the staircase) gain electrons to fill their valence shell, forming negatively charged anions. The charge of the most stable main group anion is $-(8 - \text{group number})$: 7A = -1, 6A = -2, 5A = -3. Transition metals are an exception: they form multiple stable cations with different charges, so their charge cannot be predicted from group number alone.
Exam tip: When asked for the most stable ion, always default to the octet rule prediction for main group elements; do not leave the ion neutral or give a non-standard charge unless explicitly prompted.
3. Writing Formulas for Neutral Ionic Compounds ★★☆☆☆ ⏱ 3 min
All stable ionic compounds are electrically neutral, meaning the total positive charge from cations equals the total negative charge from anions, for a net charge of zero. The criss-cross method is a simple technique to get the correct formula:
- Write the cation first, then the anion, with their correct charges.
- The absolute value of the cation charge becomes the subscript for the anion, and the absolute value of the anion charge becomes the subscript for the cation.
- Reduce the subscripts to the lowest whole number ratio by dividing by their greatest common factor.
- Enclose polyatomic ions in parentheses if their subscript is greater than 1, to indicate the subscript applies to the entire ion.
Exam tip: Always check for common factors after criss-cross; for example, $\text{Ca}_2\text{O}_2$ must be reduced to $\text{CaO}$, which is the correct formula for calcium oxide.
4. Lattice Energy Trends from Coulomb's Law ★★★☆☆ ⏱ 3 min
Per Coulomb’s law, the magnitude of the electrostatic force between two charged particles is proportional to:
E \propto \frac{q_1 q_2}{r}
where $q_1$ and $q_2$ are the charges of the two ions, and $r$ is the interionic distance (sum of the two ionic radii). This gives two key trends: (1) lattice energy increases as the product of the ion charges increases, and (2) for the same charge product, lattice energy increases as interionic distance decreases. Charge product has a much larger effect on lattice energy than radius, so always compare charge first.
Exam tip: On FRQ, you must explicitly reference Coulomb’s law and compare both charge (first) and radius (second) to earn full justification points; vague statements about "stronger bonding" are not enough.
5. AP-Style Concept Check ★★★☆☆ ⏱ 3 min
Common Pitfalls
Why: Students confuse total electrons outside the noble gas core with the highest $n$ definition of valence electrons used by AP Chemistry
Why: Students mix up the order from electron transfer diagrams that show nonmetals gaining electrons first
Why: Students do not recognize that the subscript applies to the entire polyatomic ion, not just the last atom
Why: Students stop after applying the criss-cross method and do not check for common factors
Why: Students prioritize size over charge, but charge has a much larger effect on electrostatic force
Why: Students extend the main group charge rule to all elements