Properties of Period 3 elements and their oxides

In a nutshell

Across Period 3 the elements change from metals to non-metals, and their oxides change with them: from ionic, basic oxides (Na₂O, MgO), through amphoteric Al₂O₃, to covalent, acidic oxides (SiO₂, P₄O₁₀, SO₂, SO₃).

Almost every question in this topic comes back to one skill: name the structure and bonding of the oxide, then use it to explain its melting point and how it behaves with water, acids and bases.

Assumed knowledge: Periodicity, Bonding, Group 2, the alkaline earth metals.

Core content

Reactions of sodium and magnesium with water

Sodium reacts vigorously with cold water to form sodium hydroxide and hydrogen:

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

What you see:

• The sodium floats, melts into a ball and moves around on the surface.
• It fizzes (hydrogen is released) and eventually disappears.
• A colourless solution forms, which is strongly alkaline (pH 13 to 14).

Magnesium reacts only very slowly with cold water, forming magnesium hydroxide:

Mg(s) + 2H₂O(l) → Mg(OH)₂(s) + H₂(g)

Mg(OH)₂ is only sparingly soluble, so the solution is only weakly alkaline (about pH 9 to 10).

Magnesium reacts rapidly with steam, and the product is the oxide, not the hydroxide:

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

What you see: a bright white light (flame) and a white solid (MgO). The hydrogen is colourless, so you do not see it, and there is no effervescence because the reaction is not happening in a solution.

Reactions of the elements with oxygen

All six elements Na to S burn in oxygen to form oxides. You need these equations and the observations for Na, Mg, P and S:

Points that the observations hinge on:

• Burning produces a solid, smoke or fumes, never a "precipitate". Precipitates only form in solutions.
• For sulfur, "a gas" or "a colourless gas" is not enough: say it is pungent or choking.
• Sodium is stored under oil and white phosphorus under water, to stop them reacting with oxygen in the air.
• Aluminium seems unreactive in air because a thin layer of Al₂O₃ forms on its surface; the oxide is insoluble and impermeable, so it protects the metal underneath.

P₄O₁₀ is one molecule containing 4 phosphorus atoms and 10 oxygen atoms. That is why its formula is not simplified to P₂O₅.

SO₃ forms when sulfur dioxide reacts further with oxygen. This step is reversible and slow without a catalyst:

2SO₂(g) + O₂(g) ⇌ 2SO₃(g)

In the highest oxides, the oxidation state of the Period 3 element rises by one each step (it equals the group number). Use this to check a formula:

Structure and bonding of the oxides

The type of bonding follows the electronegativity difference between the element and oxygen:

• Large difference (metal + oxygen) → ionic bonding.
• Smaller difference (non-metal + oxygen) → covalent bonding.

Use "molecular" or "simple molecular" for P₄O₁₀, SO₂ and SO₃. "Simple covalent" on its own describes the bonding, not the structure, and does not earn the structure mark.

You can show the metal oxides are ionic by melting them: the molten oxide conducts electricity. Dissolving them in water is not a valid test (MgO is only sparingly soluble, Al₂O₃ is insoluble, and many covalent substances also form ions in water).

Trend in the melting points of the highest oxides

The melting point is high for the ionic and giant covalent oxides (Na₂O to SiO₂) and low for the molecular oxides (P₄O₁₀, SO₃). For SO₂ it is lower still, about 200 K.

Explain each part of the trend by naming the structure, the force, and how strong it is:

1. Na₂O and MgO: giant ionic lattices. There is strong electrostatic attraction between oppositely charged ions, so a lot of energy is needed to overcome it.
2. MgO higher than Na₂O. The Mg²⁺ ion has a greater charge (and is smaller) than the Na⁺ ion, so it attracts the O^2- ions more strongly.
3. Al₂O₃ lower than MgO, even though Al³⁺ has a higher charge. The bonding in Al₂O₃ has some covalent character, so it is not a purely ionic lattice.
4. SiO₂: macromolecular. Melting means breaking many strong covalent bonds between atoms, which needs a lot of energy.
5. P₄O₁₀, SO₃ and SO₂: molecular. Only the weak van der Waals (and dipole-dipole) forces between molecules are overcome; the covalent bonds inside the molecules do not break. Little energy is needed.
6. Between the molecular oxides: P₄O₁₀ > SO₃ > SO₂, because a larger molecule (greater M_r, more electrons) has stronger van der Waals forces between molecules.

Still don't get it? · what actually breaks when an oxide melts

Imagine three Lego builds. The first is one huge castle where every brick is clicked firmly to its neighbours: to make it fall apart you have to unclick thousands of bricks. The second is a pile of strong magnets stuck together: still hard to pull apart. The third is a bag of small finished Lego cars that are only loosely stuck to each other with a bit of tape: you can shake them apart easily, and every car stays in one piece.

Now the chemistry, one step at a time:

1. Melting means letting the particles move past each other. You only have to overcome whatever holds the particles to each other.
2. In Na₂O and MgO the particles are ions (the magnets). Holding them together is the strong attraction between oppositely charged ions.
3. In SiO₂ there are no separate molecules at all. The whole crystal is one network of Si and O atoms joined by covalent bonds (the castle). To melt it you must break covalent bonds, lots of them.
4. In P₄O₁₀, SO₃ and SO₂ the particles are small molecules (the cars). The covalent bonds are inside each molecule and they do not break on melting. Only the weak forces between molecules (the tape) are overcome.

Exam version: always say the structure first (giant ionic lattice, macromolecular, or molecular), then name the force that is overcome (attraction between oppositely charged ions, covalent bonds between atoms, or van der Waals forces between molecules), then compare how much energy it takes. Never mix the words: "covalent bonds between molecules of SiO₂" and "intermolecular forces in SiO₂" both lose the mark.

Reactions of the oxides with water, and the pH formed

The pH values are approximate; they depend on how much oxide is added.

The trend follows the bonding: ionic oxides give alkaline solutions, covalent oxides give acidic solutions, and the insoluble oxides leave the water neutral.

Why the ionic oxides are basic. Na₂O and MgO contain oxide ions, O^2-. The oxide ion accepts a proton from water, forming hydroxide ions:

O^2-(s) + H₂O(l) → 2OH^-(aq)

The OH^- ions make the solution alkaline. MgO gives a lower pH than Na₂O because Mg(OH)₂ is only sparingly soluble, so fewer OH^- ions are released into solution.

Why Al₂O₃ and SiO₂ do not react. Water cannot supply enough energy to break the strong ionic lattice of Al₂O₃ or the many strong covalent bonds of macromolecular SiO₂, so both are insoluble and the pH stays at 7.

Why the molecular oxides are acidic. P₄O₁₀, SO₂ and SO₃ are covalent molecules that react with water to form acids, which release H⁺ ions. When a question asks for the ions formed, write the products as ions:

• P₄O₁₀(s) + 6H₂O(l) → 12H⁺(aq) + 4PO₄^3-(aq)
• SO₂(g) + H₂O(l) ⇌ H⁺(aq) + HSO₃^-(aq)
• SO₃(g) + H₂O(l) → 2H⁺(aq) + SO₄^2-(aq)

SO₂ gives a weak acid: the reaction is an equilibrium that lies mainly to the left, so the acid is only partially dissociated. That is why its pH is higher than that of SO₃.

Still don't get it? · why some oxides make alkalis and others make acids

Picture a room full of people (the water molecules), each holding a small ball. Some visitors walk in with empty, grabbing hands: they take balls away from the people. Other visitors walk in carrying spare balls and hand them out. Whether the room ends up "alkaline" or "acidic" depends on which kind of visitor arrived.

Now the chemistry, one step at a time:

1. The "ball" is a proton, H⁺.
2. An ionic oxide such as Na₂O is made of ions, and one of them is the oxide ion, O^2-. It has a 2- charge, so it strongly attracts a proton. It is the visitor with grabbing hands.
3. Each O^2- takes a proton from a water molecule. The oxide ion becomes OH^-, and the water molecule that lost its proton is also left as OH^-. So O^2- + H₂O → 2OH^-. More OH^- means alkaline.
4. A covalent oxide such as SO₃ has no oxide ions. Instead it joins with water to make an acid (H₂SO₄), whose O-H groups release protons into the water. It is the visitor carrying spare balls. More H⁺ means acidic.
5. Al₂O₃ and SiO₂ never really enter the room: they are insoluble, so the water stays neutral.

Exam version: "Na₂O is ionic and contains O^2- ions. The oxide ions accept protons from water to form OH^- ions, so the solution is alkaline." Saying only "sodium oxide forms sodium hydroxide" misses the marks for the oxide ion and the proton transfer.

Structures of the acids and anions formed

When P₄O₁₀, SO₂ and SO₃ react with water they form these acids, which lose protons to give these anions:

Features to get right when you draw them:

• Each acidic H is on an oxygen (O-H), never bonded directly to P or S.
• H₂SO₃ has one S=O and two S-O-H groups, with a lone pair on the sulfur. H₂SO₄ has two S=O and two S-O-H groups.
• In each anion, every O-H of the acid has become O^-, so the charge equals the number of protons lost.
• Partly deprotonated ions also exist and can be credited, for example HSO₃^-, HSO₄^-, H₂PO₄^- and HPO₄^2-.

Reactions of the oxides with acids and bases

Most of these are acid + base → salt + water. Work out the salt from the charges on its ions (Na⁺, Mg²⁺, Al³⁺, PO₄^3-, SO₃^2-, SO₄^2-, SiO₃^2-), rather than memorising every equation.

Basic oxides (Na₂O, MgO) react with acids:

• Na₂O(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l)
• Na₂O(s) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O(l)
• MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)
• 3MgO(s) + 2H₃PO₄(aq) → Mg₃(PO₄)₂(s) + 3H₂O(l)
• Ionic form: MgO(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂O(l)

Aluminium oxide is amphoteric: it reacts with both acids and bases.

Al₂O₃ acting as a base (with an acid):

• Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)
• Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l)
• Ionic form: Al₂O₃(s) + 6H⁺(aq) → 2Al³⁺(aq) + 3H₂O(l)

Al₂O₃ acting as an acid (with a base). This one does not follow the salt + water pattern, so learn it:

• Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq)
• Ionic form: Al₂O₃(s) + 2OH^-(aq) + 3H₂O(l) → 2[Al(OH)₄]^-(aq)

Acidic oxides react with bases:

• SiO₂(s) + 2NaOH(aq) → Na₂SiO₃(aq) + H₂O(l) (needs hot, concentrated NaOH)
• SiO₂ + CaO → CaSiO₃ (with a basic oxide, at high temperature)
• P₄O₁₀(s) + 12NaOH(aq) → 4Na₃PO₄(aq) + 6H₂O(l)
• SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l)
• SO₃(g) + 2NaOH(aq) → Na₂SO₄(aq) + H₂O(l)
• Ionic forms: P₄O₁₀(s) + 12OH^-(aq) → 4PO₄^3-(aq) + 6H₂O(l); SO₂(g) + 2OH^-(aq) → SO₃^2-(aq) + H₂O(l)

SiO₂ is classed as an acidic oxide (it reacts with bases), even though it is insoluble in water. It is not amphoteric: it does not react with acids.

An acidic oxide and a basic oxide also react directly, with no water formed:

• 6Na₂O(s) + P₄O₁₀(s) → 4Na₃PO₄(s)
• 6MgO(s) + P₄O₁₀(s) → 2Mg₃(PO₄)₂(s)

Practical skills: burning the elements and testing the oxides

The spec suggests burning the elements in oxygen and testing the pH of the oxides.

1. Heat a small sample on a deflagrating spoon and lower it into a gas jar of oxygen. Burn phosphorus and sulfur in a fume cupboard, since their oxides are harmful.
2. Record the flame colour and the product (see the oxygen table above).
3. Add a little water to the jar and shake, so the oxide reacts with or dissolves in the water.
4. Test the solution with universal indicator or a pH meter. Na₂O gives a blue or purple colour (alkaline); P₄O₁₀ gives red, and SO₂ gives red to orange (acidic).

The water step is essential: a dry oxide has no pH.

Worked examples

Example 1: writing an equation for an oxide with a given acid or base

Write an equation for the reaction of an excess of magnesium oxide with phosphoric(V) acid.

1. With excess MgO the acid is fully neutralised, so the anion is PO₄^3- (not HPO₄^2-).
2. Balance the charges in the salt: three Mg²⁺ (total 6+) with two PO₄^3- (total 6-), so the salt is Mg₃(PO₄)₂.
3. A basic oxide plus an acid gives a salt plus water (never hydrogen):

3MgO(s) + 2H₃PO₄(aq) → Mg₃(PO₄)₂(s) + 3H₂O(l)

4. Check: Mg 3 = 3; P 2 = 2; O 3 + 8 = 11 and 8 + 3 = 11; H 6 = 6.

Write an ionic equation for aluminium oxide reacting with an excess of hydroxide ions.

1. Keep Al₂O₃ as the whole formula (it is the reagent), and use OH^-, not NaOH.
2. The product is the aluminate ion [Al(OH)₄]^-; two Al atoms make two of them.
3. Balance O and H with water:

Al₂O₃(s) + 2OH^-(aq) + 3H₂O(l) → 2[Al(OH)₄]^-(aq)

4. Check: Al 2 = 2; O 3 + 2 + 3 = 8 and 2 × 4 = 8; H 2 + 6 = 8 and 2 × 4 = 8; charge 2- = 2-.

Write an equation for sulfur trioxide reacting with potassium hydroxide. The products are a salt and water, not "K₂O + H₂SO₄":

SO₃(g) + 2KOH(aq) → K₂SO₄(aq) + H₂O(l)

Example 2: using the phosphorus(V) oxide mole ratio

1.42 g of P₄O₁₀ is added to water. Calculate the mass of sodium hydroxide needed to neutralise the acid formed completely. (A_r: Na 23.0, O 16.0, H 1.0, P 31.0)

Step 1: the overall ratio. One P₄O₁₀ gives 4H₃PO₄, and each H₃PO₄ needs 3NaOH, so 1 mol P₄O₁₀ needs 12 mol NaOH:

$\text{P}_4\text{O}_{10} + 12\,\text{NaOH} \rightarrow 4\,\text{Na}_3\text{PO}_4 + 6\,\text{H}_2\text{O}$

Step 2: moles of P₄O₁₀:

$M_r(\text{P}_4\text{O}_{10}) = (4 \times 31.0) + (10 \times 16.0) = 284.0$

$n(\text{P}_4\text{O}_{10}) = \frac{1.42\ \text{g}}{284.0\ \text{g mol}^{-1}} = 5.00 \times 10^{-3}\ \text{mol}$

Step 3: moles of NaOH:

$n(\text{NaOH}) = 12 \times 5.00 \times 10^{-3}\ \text{mol} = 6.00 \times 10^{-2}\ \text{mol}$

Step 4: mass of NaOH:

$M_r(\text{NaOH}) = 23.0 + 16.0 + 1.0 = 40.0$

$m(\text{NaOH}) = 6.00 \times 10^{-2}\ \text{mol} \times 40.0\ \text{g mol}^{-1} = 2.40\ \text{g}$

Answer: 2.40 g (3 significant figures, matching the 1.42 g given).

Where the marks are lost: forgetting either the factor of 4 (one P₄O₁₀ gives four H₃PO₄) or the factor of 3 (each H₃PO₄ needs three NaOH). Using a 1 : 3 ratio gives 0.600 g and a 1 : 4 ratio gives 0.800 g, both wrong.

Example 3: model answer, melting points of silicon dioxide and phosphorus(V) oxide

Explain why silicon dioxide has a much higher melting point than phosphorus(V) oxide. Four points, in this order:

1. SiO₂ is macromolecular (giant covalent).
2. P₄O₁₀ is molecular (simple molecules).
3. To melt SiO₂, many strong covalent bonds between atoms must be broken; to melt P₄O₁₀, only weak van der Waals forces between molecules are overcome.
4. Covalent bonds are much stronger, so much more energy is needed to melt SiO₂.

Where the mark is lost: any mention of molecules or intermolecular forces in SiO₂, or calling P₄O₁₀ macromolecular.

Example 4: model answer, why sodium oxide gives an alkaline solution

Explain why sodium oxide forms an alkaline solution in water. Three points:

1. Na₂O is ionic.
2. It contains oxide ions, O^2-, which accept protons from water.
3. This forms hydroxide ions: O^2- + H₂O → 2OH^-.

Where the mark is lost: writing only "it forms NaOH" without the oxide ion and the proton transfer.

Common exam mistakes

• Writing P₂O₅ or P₂ in equations. The oxide is P₄O₁₀ and white phosphorus is P₄. The reason is that one molecule contains 4 P and 10 O atoms, not that P₄O₁₀ is "more stable".
• Not balancing the sodium equation: 2Na + O₂ → Na₂O is wrong; it is 4Na + O₂ → 2Na₂O.
• Giving a "red" flame for sodium (it is yellow or orange), or describing the white solid formed on burning as a "precipitate". Precipitates form only in solution.
• For sulfur burning, writing only "a colourless gas" or "smells of bad eggs". The marks are for a blue flame and a choking or pungent gas.
• Writing Mg(OH)₂ as the product of magnesium with steam (it is MgO), or saying you would see bubbles of gas: the hydrogen is invisible and there is no solution for effervescence.
• Calling SO₂, SO₃ or P₄O₁₀ "simple covalent" when asked for the structure. Say molecular (simple molecular).
• Thinking P₄O₁₀ is macromolecular or ionic. It is molecular, which is why its melting point is low.
• Contradicting yourself on SiO₂: "macromolecular with strong covalent bonds between the molecules", or "intermolecular forces in SiO₂". There are no molecules in SiO₂; the covalent bonds are between atoms.
• Leaving out "between molecules" when describing van der Waals forces, or calling them "bonds". "van der Waals forces" alone is not enough.
• Explaining the melting point of the elements (metallic bonding in Na and Mg) when the question asks about the oxides. Any mention of metallic bonding or molecules in Na₂O or MgO scores zero.
• Explaining why Na₂O is basic without mentioning the O^2- ion accepting a proton to form OH^-.
• Writing H₂SO₃ when the question says ions are formed from SO₂ and water. Write H⁺ + HSO₃^-.
• Using SO₃ where the question is about SO₂ (or the reverse), and confusing the acids: SO₂ gives H₂SO₃, not sulfuric(VI) acid.
• Drawing H₂SO₃ with H atoms bonded to sulfur, or with two S=O bonds. It has one S=O and two S-O-H.
• In an ionic equation for Al₂O₃, using HCl or NaOH. Use H⁺ and OH^-: Al₂O₃ + 6H⁺ → 2Al³⁺ + 3H₂O. Keep the whole oxide formula as the reactant (MgO + 2H⁺ → Mg²⁺ + H₂O), not just O^2- + 2H⁺ → H₂O.
• Giving the products of SO₃ + KOH as "K₂O + H₂SO₄". An acidic oxide plus a base gives a salt and water (K₂SO₄ + H₂O).
• Balancing an acid-base equation by producing hydrogen gas. Oxide + acid gives salt + water.
• Getting the formula of magnesium phosphate wrong (it is Mg₃(PO₄)₂), or ignoring that MgO is in excess and giving MgHPO₄.
• Suggesting you test whether an oxide is ionic by dissolving it and testing the solution's conductivity. Melt it and show the molten oxide conducts.
• Calling SiO₂ amphoteric. It is acidic: it reacts with bases but not with acids. Al₂O₃ is the amphoteric one.

Key definitions

• Basic oxide: an oxide that reacts with acids to form a salt and water. The Period 3 basic oxides are ionic and contain O^2- ions, which accept protons to form OH^-.
• Acidic oxide: an oxide that reacts with bases to form a salt (and water), and that forms an acidic solution if it reacts with water. The Period 3 acidic oxides are covalent.
• Amphoteric: reacts with both acids and bases (shows both acidic and basic properties). Al₂O₃ is amphoteric.
• Giant ionic lattice: a regular arrangement of ions held by strong electrostatic attraction between oppositely charged ions.
• Macromolecular (giant covalent): a giant structure of atoms joined by many strong covalent bonds (between atoms), with no separate molecules.
• Molecular (simple molecular): separate molecules with covalent bonds within the molecules and weak van der Waals (and dipole-dipole) forces between molecules.
• Brønsted-Lowry base: a proton acceptor (the role of the oxide ion with water).

Specification

• I can describe and write equations for the reactions of sodium with cold water and magnesium with cold water and with steam, including what I would see.
• I can write equations, with observations, for the reactions of Na, Mg, Al, Si, P and S with oxygen, forming Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₂ and SO₃.
• I can state the structure and bonding of each of these oxides.
• I can explain the trend in melting point of the highest oxides Na₂O to SO₃ in terms of their structure and bonding.
• I can write equations for the reactions of Na₂O, MgO, P₄O₁₀, SO₂ and SO₃ with water, and state that Al₂O₃ and SiO₂ do not react, giving the approximate pH of each solution.
• I can explain the trends in the reactions of the oxides with water in terms of the type of bonding in each oxide.
• I can draw the structures of the acids and anions formed when P₄O₁₀, SO₂ and SO₃ react with water.
• I can write equations for the reactions of the oxides Na₂O to SO₃ with given acids and bases, including the amphoteric reactions of Al₂O₃.
• I can describe how to burn elements in oxygen and test the pH of the oxides formed.

Related notes

• Periodicity
• Bonding
• Group 2, the alkaline earth metals
• Amount of substance
• Acids and bases
• Reactions of ions in aqueous solution (A-level)