Model electron transfer, lattice formation, charges, and formula units for ionic compounds. Predict formulas using ion charges and polyatomic ions.
A lesson of Chemistry, in Chemical Bonding and Molecular Structure, about 35 minutes.
By the end of this lesson, you will be able to predict when a transfer happens rather than a share, write a correct formula by balancing charge, name the compound, and explain every one of its properties from the lattice.
Sodium has one valence electron it barely holds. Chlorine is one short of a full
outer level and pulls hard. The electron moves, and the two resulting ions are
held together by ordinary electrostatic attraction between opposite charges.
An ionic bond is not a bond between two atoms. It is the attraction between
ions, and it acts in every direction at once, which is why an ionic compound is a
three-dimensional lattice rather than a collection of molecules.
No prefixes. "Magnesium oxide", never "magnesium monoxide". The charges fix
the ratio, so there is nothing to disambiguate, and this is exactly what
distinguishes ionic naming from covalent naming in the nomenclature lesson.
Transition metals need a Roman numeral, because they form more than one
charge:
The numeral is the metal's charge, not the number of anions. Work it out
backwards from the anions: in $\mathrm{FeCl_3}$ the three chlorides carry $-3$
altogether, so the single iron must be $+3$.
With a polyatomic ion, use its name unchanged:
$$\mathrm{Ca(NO_3)_2} = \text{calcium nitrate}$$
Sort It: Does It Need a Roman Numeral
Worked Example 2: Naming With a Numeral
$$\mathrm{Fe_2O_3}$$
The Lattice Explains Every Property
Ionic compounds are not made of molecules. They are a repeating three-dimensional
grid of alternating ions, and every characteristic property comes from that one
structure.
Very high melting points. Melting means breaking apart a lattice held by
strong attractions in every direction at once. $\mathrm{NaCl}$ melts at $801°C$;
sugar, held by weak forces between separate molecules, melts at $186°C$.
Hard but brittle. Hard because the lattice resists deformation. Brittle for
a subtle reason: hit it hard enough to slide one layer by a single ion's width
and suddenly like charges face each other. The attraction becomes repulsion and
the crystal splits cleanly along that plane. Metals, whose electrons are shared
communally, just slide and bend instead.
Conduct when molten or dissolved, never when solid. Conduction needs charged
particles that can MOVE. In the solid the ions are locked in place; melt it or
dissolve it and they are free, and the substance conducts. That test distinguishes
ionic from covalent more reliably than anything else.
Often soluble in water. Water molecules are polar, so their negative ends
surround the cations and their positive ends surround the anions, pulling the
lattice apart ion by ion.
"Formula unit", not "molecule". $\mathrm{NaCl}$ names the simplest ratio in
the lattice; there is no discrete $\mathrm{NaCl}$ particle anywhere in a salt
crystal.
Worked Example 3: Predicting Properties
A white solid melts at $995°C$ and conducts electricity when molten but not
when solid.
Common Mistakes
Mistake 1: Leaving charges in the formula
$\mathrm{NaCl}$, not $\mathrm{Na^+Cl^-}$. The compound is neutral.
Mistake 2: Not reducing the ratio
$\mathrm{Mg_2O_2}$ is $\mathrm{MgO}$.
Mistake 3: Omitting brackets on a polyatomic ion
$\mathrm{Ca(NO_3)_2}$, not $\mathrm{CaNO_6}$.
Mistake 4: Using prefixes in an ionic name
"Magnesium oxide", never "magnesium monoxide".
Mistake 5: Reading the Roman numeral as a count
It is the metal's charge, worked out from the anions.
Where This Goes Next
Two questions about ionic bonding deserve honest answers, and both take you
somewhere larger.
Why do these compounds form at all? Removing sodium's electron COSTS energy,
and chlorine gaining one releases less than that cost. On its own the transfer is
uphill. What pays for it is the lattice energy, the enormous release when
billions of ions assemble into an ordered crystal. The bond is worth forming
because of the lattice, not because of the transfer, and a single isolated
$\mathrm{Na^+Cl^-}$ pair would not be stable.
Where does ionic stop and covalent start? Nowhere sharply. Every bond
transfers some charge and shares some, and $1.7$ is a line drawn across a
continuum. $\mathrm{AlCl_3}$ has a difference of $1.5$ and behaves covalently,
melting at $190°C$ and dissolving in organic solvents, because aluminum's small
highly charged ion pulls electron density back off the chloride. That is called
polarization, and Fajans' rules predict when it happens: small, highly charged
cations with large soft anions drift toward covalent behavior.
So the categories are useful and approximate, exactly like the states of matter
and the physical-chemical line. Real substances sit on a spectrum, and the ones
in the middle are usually the interesting ones.
Check Your Understanding
Five of them, no instructions this time.
Summary
Ionic bonding is transfer, and the electrostatic attraction acts in every
direction, producing a lattice rather than molecules.
Balance the charges to get the formula, then reduce, and use brackets for
multiple polyatomic ions.
Name the metal, then the nonmetal as -ide, with no prefixes and a Roman
numeral for transition metals.
Every property comes from the lattice: high melting point, hard but brittle,
and conducting only when the ions can move.