SCH3U Ionic Bonding & Compounds: Notes and Practice

Atoms are most stable with a full valence shell — usually 8 electrons (the octet rule), matching the electron arrangement of the nearest noble gas. Metals achieve this by losing their few valence electrons; nonmetals achieve it by gaining electrons to fill their shell. When a metal meets a nonmetal, electrons transfer from one to the other, and the resulting oppositely charged ions attract each other — this is ionic bonding. Table salt, the calcium chloride spread on Ontario roads in winter, the calcium carbonate in limestone and the electrolytes in a sports drink are all ionic compounds, and their shared properties all trace back to this one idea.

Why electrons transfer from metals to nonmetals

Look at the periodic trends from the previous lesson. Metals, on the left, have low ionization energies: their one to three valence electrons are far from the nucleus and weakly held, so they are relatively easy to remove. Nonmetals, on the right, have high electronegativity and a strong pull on incoming electrons, because their $Z_{e ff}$ is high and their outer shell is nearly full. Put the two together and the electron moves from the atom that holds it loosely to the atom that pulls hard. Removing an electron from sodium actually costs energy, so why does the reaction happen at all? Because the newly formed ions attract each other so strongly that far more energy is released when they pack together than was spent making them. That is why sodium and chlorine react violently to form salt. The octet rule is a handy shorthand for where the transfer stops: each ion ends up with the electron arrangement of a noble gas.

Fig 1Lewis (electron dot) symbols show only the valence electrons as dots around the element symbol.

A Lewis symbol draws an element's chemical symbol surrounded by dots representing only its valence electrons. Lewis symbols make it easy to see, and predict, how many electrons an atom needs to lose or gain to reach a full octet.

Forming an ionic compound

Metal loses its valence electron(s) → becomes a positively charged cation (smaller than the neutral atom, now matches noble-gas electron count). Nonmetal gains electron(s) → becomes a negatively charged anion (larger than the neutral atom). The cation and anion attract each other electrostatically, and combine in whatever ratio makes the overall compound electrically neutral.

ACation (metal)

Loses electrons
Becomes positively charged
Becomes smaller than the neutral atom
e.g. Na → Na⁺ + e⁻

BAnion (nonmetal)

Gains electrons
Becomes negatively charged
Becomes larger than the neutral atom
e.g. Cl + e⁻ → Cl⁻

Writing ionic formulas

Every ionic compound is electrically neutral: the total positive charge must exactly cancel the total negative charge. That single rule sets the formula. First, find each ion's charge from its group (Group 1 → 1+, Group 2 → 2+, Group 13 → 3+, Group 15 → 3−, Group 16 → 2−, Group 17 → 1−). Next, find the smallest numbers of each ion that make the charges add to zero. Finally, write the metal first with those numbers as subscripts, leaving out any 1. The criss-cross shortcut does the balancing for you: the size of each ion's charge becomes the other ion's subscript. It always gives a neutral combination, but not always the simplest one, so reduce the subscripts to the lowest whole-number ratio at the end. Because an ionic solid is a lattice rather than a molecule, its formula describes a ratio of ions called a formula unit, not a molecule that exists on its own.

Fig 2The criss-cross method: each charge becomes the other ion's subscript, then the ratio is reduced (Mg2O2 becomes MgO).

Ions	Balancing the charge	Formula
$Na^{+}$ and $Cl^{-}$	(+1) + (−1) = 0	NaCl
$Ca^{2 +}$ and $Cl^{-}$	(+2) + 2(−1) = 0	$CaCl_{2}$
$K^{+}$ and $S^{2 -}$	2(+1) + (−2) = 0	$K_{2} S$
$Al^{3 +}$ and $O^{2 -}$	2(+3) + 3(−2) = 0	$Al_{2} O_{3}$
$Mg^{2 +}$ and $N^{3 -}$	3(+2) + 2(−3) = 0	$Mg_{3} N_{2}$
$Ca^{2 +}$ and $O^{2 -}$	(+2) + (−2) = 0	CaO (not $Ca_{2} O_{2}$ )

In every correct formula the positive and negative charges add to zero.

Lattices and the properties they cause

Fig 3Ions pack into a repeating 3-D crystal lattice, held together by electrostatic attraction in every direction.

Ionic compounds don't exist as individual molecules — they form giant, repeating lattices where each ion is surrounded by ions of the opposite charge on all sides. In a sodium chloride crystal, every $Na^{+}$ touches six $Cl^{-}$ ions and every $Cl^{-}$ touches six $Na^{+}$ ions, so no single chloride "belongs" to any one sodium. This extensive 3-D network of attractions is what gives ionic compounds their characteristic properties. Larger charges and smaller ions attract more strongly, which is why magnesium oxide ( $Mg^{2 +}$ with $O^{2 -}$ ) melts at about 2850 °C while sodium chloride melts at 801 °C.

Property	Explanation
High melting/boiling point	Breaking apart the lattice requires overcoming many strong electrostatic attractions at once
Hard but brittle	A shifted layer lines up same-charge ions, which repel and shatter the crystal
Conducts electricity when molten or dissolved (aqueous)	Ions become free to move and carry charge
Does not conduct as a solid	Ions are locked in fixed lattice positions and cannot move
Many dissolve in water	Polar water molecules surround and pull ions out of the lattice (covered later in this unit)

Ionic properties at work

Because molten ionic compounds conduct, they can be split by electricity: aluminum metal is made by passing a current through aluminum oxide dissolved in a molten salt, and sodium metal is made from molten sodium chloride. Road salt works because dissolved ions interfere with ice forming, and a solution of ions in your body (electrolytes such as

Na^{+}

K^{+}

and

Cl^{-}

) carries the electrical signals in your nerves.

Ionic Bonding & Compounds

Key things to remember

Why electrons transfer from metals to nonmetals

Writing ionic formulas

Lattices and the properties they cause

Worked examples

Where marks get lost

Flashcards 10

Practice