Valence Electrons and Ionic Compounds价电子与离子化合物

Why is sodium always +1 and oxygen always −2?

Table salt is NaCl, not $\ce{NaCl2}$; magnesium oxide is MgO, not $\ce{Mg2O}$. Whether two elements bond, and in what ratio, depends on the interactions between their valence electrons and nuclei. This lesson applies U1.5 (configurations) and U1.7 (ionization energy, electronegativity) to compounds: from an element’s place in the periodic table, predict its ion charge and the formulas of its compounds.

Valence electrons set the ion charge

Main-group elements usually form ions with the same configuration as a noble gas:

Why exactly these charges? Coulomb’s law again:

• Metals lose their valence electrons. Valence electrons are far from the nucleus and shielded by the core, so removing them takes a moderate amount of energy. Going further means removing core electrons, and the ionization energy jumps several-fold (the successive ionization energies of U1.7). So Na stops at +1 and Mg at +2.
• Nonmetals gain electrons. Added electrons go into the same unfilled shell and feel a fairly large effective nuclear charge. Once that shell is full, another electron would have to start a new shell farther out and well shielded, where it isn’t attracted enough. So O stops at −2 and Cl at −1.

Transition-metal charges can’t be read straight from the group number; iron, for example, forms both $\ce{Fe^2+}$ and $\ce{Fe^3+}$ (U1.5).

Elements in the same group form similar compounds

Elements in one group have the same number of valence electrons, so they form ions with the same charge and compounds with the same formula pattern:

• All the group 2 chlorides are $\ce{MCl2}$: $\ce{MgCl2}$, $\ce{CaCl2}$, $\ce{SrCl2}$, $\ce{BaCl2}$. The fluorides $\ce{MgF2}$ and $\ce{BaF2}$ follow the same pattern.
• Bromine has 7 valence electrons, like chlorine, so its chemistry resembles chlorine’s: NaBr is like NaCl, and HBr like HCl.
• Covalent compounds follow suit: carbon’s 4 valence electrons bond to 4 halogen atoms in both $\ce{CCl4}$ and $\ce{CBr4}$, and neither burns.

The reason must be the valence electrons (or a low ionization energy, meaning electrons are easily lost), not “similar size” or “similar mass”. Elements in a group differ a lot in size and mass: Mg’s covalent radius is 141 pm and Ba’s 215 pm, and their atomic masses differ by a factor of more than 5.

Predicting the oxides of rubidium

Rubidium is in group 1 with the valence configuration 5s¹, so it forms $\ce{Rb+}$. Charge balance with $\ce{O^2-}$ predicts the oxide $\ce{Rb2O}$. When the alkali metals actually burn in oxygen, the products vary: lithium gives $\ce{Li2O}$, sodium mainly the peroxide $\ce{Na2O2}$, and potassium, rubidium, and cesium mainly superoxides such as $\ce{RbO2}$; rubidium peroxide, $\ce{Rb2O2}$, also exists. In every one of these, the metal is +1; what changes is the anion: oxide $\ce{O^2-}$, peroxide $\ce{O2^2-}$, or superoxide $\ce{O2-}$. The valence electrons fix the metal ion’s charge, and that doesn’t change.

Reactivity and first ionization energy

The alkali metals react with water more and more violently from Li to Cs (lesson 2.3.6). The explanation AP accepts: each step down adds a shell, so the valence electron is farther from the nucleus and the first ionization energy falls (Li 520, Na 496, K 419, Rb 403, Cs 376 kJ/mol), making the electron easier to lose. “Potassium is heavier” and “potassium is less dense than sodium” are not reasons.

Ionic character and electronegativity difference

The larger the difference in electronegativity between two elements, the more the bonding electrons shift to one side and the more ionic the bond. As lesson 2.3.5 said, a difference above about 1.7 is usually taken to mean mostly ionic bonding (a rule of thumb). CsF has a difference of 3.19 and is almost entirely ionic; $\ce{AlCl3}$ differs by only 1.55 and has considerable covalent character.

An ionic compound is a huge array of alternating cations and anions with strong Coulombic attraction between neighboring opposite charges. Melting it means overcoming those attractions, so ionic compounds generally have high melting and boiling points (NaCl melts at about 800 °C). Larger ion charges and smaller ions mean stronger attraction, a point U2.3 develops.

Pick a metal and a nonmetal to see the ions and the formula; then swap in other elements from the same groups and watch whether the formula pattern changes:

Example 1: Predicting formulas

Predict the formulas of the compounds formed by (1) calcium and nitrogen, and (2) aluminum and sulfur.

Example 2: Why do BaF₂ and MgF₂ have the same formula pattern?

Magnesium fluoride is $\ce{MgF2}$. Predict the formula of barium fluoride and explain it in terms of atomic structure.

Common mistakes

Quick check

AP-style free response

Summary

• Main-group ion charges come from the valence electrons: group 1 +1, group 2 +2, group 13 +3, group 15 −3, group 16 −2, group 17 −1. Write formulas by charge balance.
• Elements in the same group have the same number of valence electrons and form compounds with the same pattern ($\ce{MgF2}$ and $\ce{BaF2}$, $\ce{CCl4}$ and $\ce{CBr4}$); the reason is the valence electrons, not size or mass.
• Alkali metals get more reactive down the group because the valence electron is farther away and the first ionization energy is lower.
• A larger electronegativity difference means more ionic bonding; the strong Coulombic attraction between ions gives ionic compounds high melting and boiling points.