2.3.1 Ions and the formation of ionic bonds

Losing or gaining electrons turns an atom into an ion

1. An atom holds equal numbers of protons and electrons, so its charges cancel and it is neutral.
2. Removing an electron leaves more protons than electrons, so the particle carries a positive charge.
3. Adding an electron leaves more electrons than protons, so the particle carries a negative charge.
4. Only electrons move, because the protons sit inside the nucleus and stay there.
5. The size of the charge equals the number of electrons lost or gained.
6. A sodium atom that loses one electron becomes ${Na}^{+}$, and a chlorine atom that gains one becomes ${Cl}^{-}$.

Groups 1, 2, 6 and 7 tell you which ion an atom forms

1. For groups $1$, $2$, $6$ and $7$, the group number gives the number of outer-shell electrons.
2. A group $1$ atom loses its single outer electron to form a $1+$ ion.
3. A group $2$ atom loses both outer electrons to form a $2+$ ion.
4. A group $6$ atom gains two electrons to form a $2-$ ion.
5. A group $7$ atom gains one electron to form a $1-$ ion.
6. Either route leaves the ion with a full outer shell, which is why these are the ions that form.
7. Groups $1$ and $2$ hold metals that supply electrons, while groups $6$ and $7$ hold non-metals that accept them.

An ionic bond is the attraction between the ions that form

1. Electron transfer from a metal atom to a non-metal atom produces one positive ion and one negative ion.
2. Opposite charges attract, and that attraction is what holds the two ions together.
3. The transfer makes the ions, and the attraction that follows is the bond itself.
4. Sodium gives one electron to chlorine: $Na\to {Na}^{+}+e^{-}$ and $Cl+e^{-}\to {Cl}^{-}$.
5. Magnesium gives two electrons to oxygen, forming ${Mg}^{2+}$ and $O^{2-}$.
6. The compound as a whole is neutral, because the positive and negative charges inside it balance exactly.

Dot-and-cross diagrams show which atom supplied each electron

1. A dot-and-cross diagram draws the outer-shell electrons of each atom, using dots for one atom and crosses for the other.
2. Dots and crosses both stand for electrons, and the two symbols only record where each electron came from.
3. Draw both atoms with their outer shells first, then move the transferred electrons across.
4. For sodium chloride, chlorine's seven electrons are dots and the single electron from sodium is a cross, giving eight around chlorine.
5. Enclose each ion in square brackets and write its charge outside the top right corner.
6. The sodium ion is drawn with an empty outer shell and a $1+$ charge, because its outer electron has gone.
7. For magnesium chloride, draw one magnesium atom and two chlorine atoms, with one electron moving to each chlorine.

Counting protons, neutrons and electrons in an ion

1. The proton count does not change when an ion forms, so it is still the atomic number.
2. The neutron count is the mass number minus the atomic number, and it does not change either.
3. For a positive ion, subtract the charge from the atomic number to get the electron count.
4. For a negative ion, add the charge to the atomic number to get the electron count.
5. ${Na}^{+}$, with atomic number $11$ and mass number $23$, has $11$ protons, $12$ neutrons and $10$ electrons.
6. $O^{2-}$, with atomic number $8$ and mass number $16$, has $8$ protons, $8$ neutrons and $10$ electrons.
7. Both of those ions hold $10$ electrons, the same number as a neon atom.

2.3.2 Naming and deducing formulae of ionic compounds

Names read positive ion first, then negative ion

1. The positive ion is named first and the negative ion second, so $NaCl$ is sodium chloride.
2. The metal supplies the positive ion, and it keeps its own element name unchanged.
3. The negative ion's name carries an ending that identifies which ion it is.
4. A finished formula carries no charges, because the positive and negative charges inside it cancel.
5. Naming and formula writing are two directions of one step: names give the ions, and charges give the ratio.

The ending -ide marks an ion made from one element

1. An -ide ending shows a negative ion built from a single element.
2. Oxygen forms the oxide ion $O^{2-}$, so $MgO$ is magnesium oxide.
3. The halogens form halide ions: fluoride $F^{-}$, chloride ${Cl}^{-}$, bromide ${Br}^{-}$ and iodide $I^{-}$.
4. ${CaCl}_2$ is therefore calcium chloride, whatever the number of chloride ions in it.
5. Hydroxide, ${OH}^{-}$, is the exception, because it ends in -ide but contains two elements.
6. Oxide and hydroxide are different ions and are never interchangeable in a formula.

The ending -ate marks an ion that also contains oxygen

1. An -ate ending shows a negative ion that contains oxygen alongside another element.
2. Nitrate is ${NO}_3^{-}$, carbonate is ${CO}_3^{2-}$ and sulfate is ${SO}_4^{2-}$.
3. The atoms inside one of these ions are held by covalent bonds and travel as a single unit.
4. The ion keeps both its name and its internal formula inside any compound.
5. ${KNO}_3$ is potassium nitrate, ${CaCO}_3$ is calcium carbonate and ${MgSO}_4$ is magnesium sulfate.

Balancing the charges gives the formula

1. Write the positive ion with its charge and the negative ion with its charge.
2. Find the smallest numbers of each ion that bring the total charge to zero.
3. With $A^{m+}$ and $B^{n-}$, taking $n$ of A and $m$ of B always balances, and the ratio is then simplified if both share a factor.
4. Write each count as a subscript, leaving out a subscript of one.
5. Put brackets round a polyatomic ion whenever more than one of it is needed.
6. ${Al}^{3+}$ with $O^{2-}$ needs two and three, giving ${Al}_2{O}_3$.
7. ${Ca}^{2+}$ with ${OH}^{-}$ needs one and two, giving ${Ca(OH)}_2$.
8. ${Al}^{3+}$ with ${NO}_3^{-}$ needs one and three, giving ${Al(NO}_3{)}_3$.

Checking a finished formula

1. Multiply each ion's charge by its subscript and add the results, which must come to zero.
2. ${Al}_2({SO}_4{)}_3$ gives $(2\times +3)+(3\times -2)=0$.
3. Confirm that each polyatomic ion still carries its own internal formula.
4. Confirm that the ratio cannot be simplified any further.
5. Confirm that no charges have been left on the finished formula.

2.3.3 Giant ionic lattices

An ionic compound is one giant lattice, not a set of pairs

1. A solid ionic compound is a single enormous structure rather than a collection of separate units.
2. Positive and negative ions alternate in a fixed, repeating pattern in three dimensions.
3. The pattern continues for as long as the crystal does, so even a visible grain contains a vast number of ions.
4. The word giant records that the number of ions is not fixed by the formula.
5. The formula gives the ratio of the ions present, not how many of them there are.

Every ion is held by attractions acting in all directions

1. Each positive ion is surrounded by negative ions, and each negative ion by positive ions.
2. In sodium chloride, every ion has six nearest neighbours of the opposite charge.
3. The attractions act in every direction through the structure, not along single lines between chosen pairs.
4. Breaking the structure apart means overcoming very many strong attractions at once.
5. Larger charges pull harder, so the attractions in $MgO$ are stronger than those in $NaCl$.

Describing and drawing a lattice

1. A full description names both the arrangement and the force.
2. The arrangement is a regular, repeating three-dimensional pattern of oppositely charged ions.
3. The force is strong electrostatic attraction between those oppositely charged ions.
4. Saying only that the forces are strong leaves the description unfinished, because it does not name them.
5. A diagram should show many ions in a repeating grid with their charges marked.
6. Drawing sodium chloride as two ions side by side describes a molecule, which an ionic compound does not form.