Why do ionic compounds shatter instead of bending like metals?

When a force shifts one layer of ions in the crystal, ions of the same charge suddenly line up next to each other. Like charges repel strongly, and the crystal splits along that plane. In a metal there are no alternating charges — the delocalized electrons hold the positive cores together no matter how the layers slide — so metals deform instead of breaking.

How do I know whether a compound is ionic or covalent?

Check the elements. A metal combined with a nonmetal is almost always ionic, and any compound containing a polyatomic ion such as nitrate, sulfate, or ammonium is ionic. Two nonmetals together share electrons and form covalent molecules. The underlying reason is electronegativity difference: a large gap favors electron transfer, a small gap favors sharing.

Why is it wrong to call NaCl a molecule?

There is no separate particle containing exactly one sodium and one chlorine. In the crystal each sodium ion is surrounded by six chloride ions and each chloride by six sodium ions, extending through the whole solid. The formula gives the one-to-one ratio of ions, which is why it is called a formula unit rather than a molecular formula.

How do I get the charge of a transition metal in a compound?

Work backward from the anion, since the compound must be neutral. In FeCl_3, three chloride ions supply -3, so the single iron must be Fe^3+. If you are given a name, the Roman numeral states it directly: iron(III) chloride means Fe^3+. Naming conventions get their own treatment later in this unit.

CHEM-4.1

Ionic Bonds & Ionic Compounds

Learn how metals and nonmetals form ions, how to write balanced ionic formulas, and why ionic crystals melt high, shatter, and conduct only when melted or dissolved.

What you'll do in this lesson

A voice-first session with the Crimsora tutor on Ionic Bonds & Ionic Compounds, then targeted practice and FRQs — with the tutor adapting to where you get stuck.

Start Ionic Bonds & Ionic Compounds free

What this lesson covers

Table salt survives a 700 degree Celsius oven without melting, shatters if you hit a crystal with a hammer, and turns pure water into something that will light a conductivity bulb. All three behaviors come from the same source: sodium and chlorine did not share electrons, they traded them, and the resulting charged particles locked into a rigid three-dimensional grid.

In this lesson you will predict what charge an atom takes when it forms an ion, combine cations and anions into formulas that come out electrically neutral, and then use the picture of a crystal lattice to explain melting point, brittleness, and conductivity. The reasoning chain runs one direction the whole way: position on the periodic table gives charge, charge gives formula, and lattice structure gives properties. Once you can run that chain, ionic compounds stop being memorization.

Electron Transfer and Predictable Ion Charges

An ionic bond is the electrostatic attraction between oppositely charged ions produced when one or more electrons transfer from a metal atom to a nonmetal atom. Metals have low ionization energies, so they give electrons up; nonmetals have high electron affinities, so they take them. Neither atom is being generous — both end up with a noble-gas electron configuration, which is a lower-energy arrangement.

For main-group elements the charge is predictable from the group number, because the group tells you the number of valence electrons.

Group	Example	Ion formed	Electrons lost or gained
1	Na	$\text{Na}^{+}$	loses 1
2	Mg	$\text{Mg}^{2+}$	loses 2
13	Al	$\text{Al}^{3+}$	loses 3
15	N	$\text{N}^{3-}$	gains 3
16	O	$\text{O}^{2-}$	gains 2
17	Cl	$\text{Cl}^{-}$	gains 1

Notice the pattern: metals on the left lose enough electrons to empty the outer shell, and nonmetals on the right gain enough to fill it. A chloride ion,

\text{Cl}^{-}

, has the same electron count as argon but still has 17 protons, so it is not argon — it is a chlorine atom carrying an extra unit of negative charge.

Two places students slip. First, gaining electrons makes an ion more negative, not more positive; the sign follows the charge added, not the word "gain." Second, most transition metals have more than one possible charge (

\text{Fe}^{2+}

and

\text{Fe}^{3+}

\text{Cu}^{+}

and

\text{Cu}^{2+}

), so you cannot read their charge off the periodic table. For those you need information from the compound itself or from a Roman numeral in the name.

Writing Neutral Formulas

A correct ionic formula obeys one rule: the total positive charge must exactly cancel the total negative charge, so the compound is electrically neutral overall. Everything else is bookkeeping.

The fastest method is the crossover: write the cation and anion with their charges, then use the magnitude of each charge as the subscript on the other ion. Aluminum and oxygen give

\text{Al}^{3+}

and

\text{O}^{2-}

, so the 3 becomes the subscript on O and the 2 becomes the subscript on Al, producing

\text{Al}_2\text{O}_3

. Check it:

2(+3) + 3(-2) = 0

.

Always reduce the subscripts to the lowest whole-number ratio. Magnesium and oxygen crossed over would give

\text{Mg}_2\text{O}_2

, which reduces to

\text{MgO}

. The formula names a ratio, not a count of atoms in a particle.

Polyatomic ions — groups of covalently bonded atoms carrying an overall charge, such as

\text{NO}_3^{-}

\text{SO}_4^{2-}

\text{CO}_3^{2-}

\text{PO}_4^{3-}

, and

\text{NH}_4^{+}

— behave as single units. If you need more than one of them, wrap the whole ion in parentheses before adding the subscript: calcium nitrate is

\text{Ca(NO}_3)_2

, not

\text{CaNO}_{32}

and not

\text{CaN}_2\text{O}_6

.

The most common error is changing the charge of an ion to make a formula work. Charges are fixed by the periodic table; only subscripts are yours to adjust. If a formula will not balance, you have the wrong charge, not the wrong element. A second frequent error is writing

\text{Na}_2\text{Cl}

because sodium "has one extra electron" — one

\text{Na}^{+}

already cancels one

\text{Cl}^{-}

exactly, so the formula is

\text{NaCl}

The Crystal Lattice and the Formula Unit

There is no such thing as an NaCl molecule. In solid sodium chloride, every

\text{Na}^{+}

is surrounded by six

\text{Cl}^{-}

ions and every

\text{Cl}^{-}

is surrounded by six

\text{Na}^{+}

ions, repeating in all directions through the entire crystal. A grain of salt is one giant assembly of ions held together by attractions that reach in every direction at once.

Because of that, the formula

\text{NaCl}

describes a formula unit: the simplest whole-number ratio of ions in the lattice, not a discrete particle you could isolate. This is the single biggest conceptual difference between this lesson and the covalent bonding you meet elsewhere in the unit, where

\text{H}_2\text{O}

really does name one separate molecule.

The energy released when gaseous ions assemble into one mole of a lattice is called lattice energy, and its size follows Coulomb's law: attraction grows as charges grow and shrinks as the distance between ion centers grows.

F \propto \frac{q_1 q_2}{r^2}

\text{MgO}

, built from

2+

and

2-

ions, has a far larger lattice energy than

\text{NaCl}

, built from

1+

and

1-

ions of similar size. Among compounds with the same charges, smaller ions sit closer together and bind more strongly:

\text{LiF}

beats

\text{KI}

. This one relationship lets you rank melting points without looking anything up, and it explains why magnesium oxide is used to line furnaces while sodium chloride is used on food.

Explaining the Physical Properties

Every classic property of an ionic compound traces back to the lattice.

High melting and boiling points. Melting requires breaking enough of the electrostatic attractions to let ions move past one another. Those attractions are strong and there are many per ion, so it takes a lot of thermal energy. Sodium chloride melts near 801 degrees Celsius; magnesium oxide, with doubled charges, melts near 2852 degrees Celsius.

Brittleness. Ionic crystals are hard but shatter under a sharp blow. Struck hard enough, one plane of ions slides by one ion-width. That shift puts

+

next to

+

and

-

next to

-

, and the sudden repulsion splits the crystal along a clean plane. Metals, by contrast, bend, because their delocalized electrons tolerate the same slippage.

Conductivity. Charge only flows if charged particles can move. In the solid, ions are locked in place, so a salt crystal does not conduct. Melt it or dissolve it in water and the ions become mobile, and the liquid conducts well. A substance that conducts when molten or dissolved is called an electrolyte.

State	Are ions mobile?	Conducts?
Solid crystal	No, fixed in lattice	No
Molten liquid	Yes	Yes
Dissolved in water	Yes, ions separated	Yes

A complete explanation always names the mechanism, not just the fact. "Ionic compounds conduct when dissolved" restates the observation; "dissolving separates the lattice so ions are free to move and carry charge" explains it. When a question asks why, your answer should mention ions, charges, and either the strength of attraction or the ability of ions to move.

Recognizing Ionic Compounds and Common Traps

Before applying any of this, confirm the compound is actually ionic. The reliable signal is a metal bonded to a nonmetal, or any compound containing a polyatomic ion such as ammonium or sulfate. A large difference in electronegativity between the two elements is the underlying reason; roughly speaking, differences above about 1.7 point toward electron transfer rather than sharing.

Some pitfalls worth naming directly. Ionic compounds are neutral, so you never write a charge on the finished formula —

\text{CaCl}_2

, not

\text{CaCl}_2^{2+}

. Subscripts of 1 are never written. Parentheses appear only when you need more than one polyatomic ion, so it is

\text{NaNO}_3

but

\text{Mg(NO}_3)_2

.

Watch out for the assumption that a bigger subscript means a stronger bond. In

\text{Na}_2\text{O}

the subscript 2 exists because sodium is only

1+

while oxide is

2-

; it says nothing about bond strength. Strength comes from charge magnitude and ion size, so

\text{MgO}

with no subscripts at all is bound far more tightly than

\text{Na}_2\text{O}

.

Finally, solubility and ionic character are different questions. Many ionic compounds dissolve readily in water because polar water molecules surround and stabilize the separated ions, but plenty do not — calcium carbonate and silver chloride are ionic and nearly insoluble. Insolubility does not make a compound covalent. The test for ionic character is what the elements are and whether electrons transferred, not whether the solid disappears in a beaker.

Key terms

Ionic bond.: The electrostatic attraction between oppositely charged ions formed when electrons transfer from a metal atom to a nonmetal atom.
Cation.: A positively charged ion, formed when an atom loses one or more electrons; metals form cations.
Anion.: A negatively charged ion, formed when an atom gains one or more electrons; nonmetals form anions.
Formula unit.: The simplest whole-number ratio of ions in an ionic compound, such as $\text{NaCl}$ ; it is a ratio, not a discrete molecule.
Crystal lattice.: The repeating three-dimensional array of alternating cations and anions that makes up an ionic solid.
Lattice energy.: The energy released when gaseous ions come together to form one mole of an ionic solid; it increases with larger ion charges and smaller ion radii.
Polyatomic ion.: A group of covalently bonded atoms carrying an overall charge, such as $\text{SO}_4^{2-}$ or $\text{NH}_4^{+}$ , that acts as a single ion in a formula.
Electrolyte.: A substance that conducts electricity when molten or dissolved in water because its ions are free to move.

Worked example

Aluminum reacts with sulfur to form an ionic compound, and separately with the sulfate ion to form another. Write the formula for each, then predict which of aluminum sulfide or sodium chloride has the higher melting point and explain why in terms of lattice structure.

Step 1 — Find the ion charges. Aluminum is in group 13, so it loses 3 electrons to give

\text{Al}^{3+}

. Sulfur is in group 16, so it gains 2 electrons to give

\text{S}^{2-}

.

Step 2 — Balance the charge for aluminum sulfide. The lowest common multiple of 3 and 2 is 6, so you need two

\text{Al}^{3+}

ions (

+6

total) and three

\text{S}^{2-}

ions (

-6

total). The formula is

\text{Al}_2\text{S}_3

. Check:

2(+3) + 3(-2) = 0

. The subscripts 2 and 3 share no common factor, so no reducing is needed.

Step 3 — Handle the polyatomic version. Sulfate is

\text{SO}_4^{2-}

, the same

2-

charge as sulfide, so the ratio is again two to three. Because you need more than one sulfate unit, it must go in parentheses:

\text{Al}_2(\text{SO}_4)_3

. Writing

\text{Al}_2\text{S}_3\text{O}_{12}

would be wrong, since it hides the fact that sulfate is a single bonded unit.

Step 4 — Compare melting points. Melting means supplying enough energy to pull ions out of their fixed positions in the lattice. Coulomb's law says attraction scales with the product of the charges,

q_1 q_2

. In

\text{NaCl}

that product is

(1)(1) = 1

; in

\text{Al}_2\text{S}_3

it is

(3)(2) = 6

. The much larger charges in aluminum sulfide produce far stronger electrostatic attractions throughout the lattice, so

\text{Al}_2\text{S}_3

has the higher melting point. Measured values agree: sodium chloride melts near 801 degrees Celsius, aluminum sulfide above 1100 degrees Celsius.

Practice questions

A solid white compound does not conduct electricity, but after it is melted the liquid conducts well. Which statement best explains this behavior?

The compound is covalent, and melting breaks its molecules into atoms.
The compound is ionic; its ions are locked in a lattice when solid but become mobile when molten.
Melting converts the neutral compound into ions for the first time.
The compound is metallic, and heating releases its delocalized electrons.

Answer: The compound is ionic; its ions are locked in a lattice when solid but become mobile when molten.

Conduction requires mobile charge carriers. The ions already exist in the solid — the compound was built from them — but they are held in fixed lattice positions, so no charge can flow. Melting overcomes the electrostatic attractions enough to let ions move, and the mobile ions carry current. The idea that melting creates ions is a common wrong answer; the ions were there all along, they simply could not move.

Write the formula for the compound formed between calcium and phosphorus, and for the compound formed between calcium and the phosphate ion. Explain why one requires parentheses and the other does not.

Answer: $\text{Ca}_3\text{P}_2$ and $\text{Ca}_3(\text{PO}_4)_2$ .

Calcium is in group 2 and forms

\text{Ca}^{2+}

; phosphorus is in group 15 and forms

\text{P}^{3-}

. Balancing to a total of six positive and six negative charges gives three calcium ions and two phosphide ions:

\text{Ca}_3\text{P}_2

. Phosphate,

\text{PO}_4^{3-}

, carries the same

3-

charge, so the ratio is again three to two — but phosphate is a single covalently bonded unit, and you need two of them. Parentheses show that the subscript 2 applies to the entire

\text{PO}_4

group, giving eight oxygen atoms in all. Without parentheses,

\text{Ca}_3\text{PO}_{42}

would claim forty-two oxygen atoms attached to one phosphorus.

Magnesium oxide and sodium fluoride have nearly the same distance between ion centers, yet magnesium oxide melts at a much higher temperature. Explain why.

Answer: Magnesium oxide is built from $2+$ and $2-$ ions while sodium fluoride is built from $1+$ and $1-$ ions, so the electrostatic attraction in the magnesium oxide lattice is about four times stronger and requires far more energy to break.

Coulomb's law,

F \propto \frac{q_1 q_2}{r^2}

, has two variables. The problem tells you

r

is essentially the same for both, so the difference must come from charge. For

\text{MgO}

the charge product is

(2)(2) = 4

; for

\text{NaF}

it is

(1)(1) = 1

. Stronger attractions throughout the lattice mean more thermal energy is needed before ions can break free and flow, which is exactly what a higher melting point measures. A complete answer names both the charges and the connection between attraction strength and the energy needed to melt.

FAQ

Why do ionic compounds shatter instead of bending like metals?: When a force shifts one layer of ions in the crystal, ions of the same charge suddenly line up next to each other. Like charges repel strongly, and the crystal splits along that plane. In a metal there are no alternating charges — the delocalized electrons hold the positive cores together no matter how the layers slide — so metals deform instead of breaking.
How do I know whether a compound is ionic or covalent?: Check the elements. A metal combined with a nonmetal is almost always ionic, and any compound containing a polyatomic ion such as nitrate, sulfate, or ammonium is ionic. Two nonmetals together share electrons and form covalent molecules. The underlying reason is electronegativity difference: a large gap favors electron transfer, a small gap favors sharing.
Why is it wrong to call NaCl a molecule?: There is no separate particle containing exactly one sodium and one chlorine. In the crystal each sodium ion is surrounded by six chloride ions and each chloride by six sodium ions, extending through the whole solid. The formula gives the one-to-one ratio of ions, which is why it is called a formula unit rather than a molecular formula.
How do I get the charge of a transition metal in a compound?: Work backward from the anion, since the compound must be neutral. In $\text{FeCl}_3$ , three chloride ions supply $-3$ , so the single iron must be $\text{Fe}^{3+}$ . If you are given a name, the Roman numeral states it directly: iron(III) chloride means $\text{Fe}^{3+}$ . Naming conventions get their own treatment later in this unit.

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