Periodicity of chemical properties of the elements in Period 3

1. Overview

Across Period 3 the elements change from reactive metals to non-metals, and their oxides and chlorides change with them: from ionic solids that give alkaline or neutral solutions to covalent substances that are acidic. This section covers the reactions of the elements with oxygen, chlorine and water, the oxidation numbers in their oxides and chlorides, and how these compounds behave with water, acids and sodium hydroxide. Almost all of it follows from one idea: how electronegative each element is compared with oxygen and chlorine. (The physical properties of the elements are in topic 9.1.)

Key Definitions

• Electronegativity: the power of an atom to attract electrons to itself.
• Oxidation number: the charge an atom would have if all its bonds were ionic, with the electrons of each bond given to the more electronegative atom (or shared equally between identical atoms); it is written with a sign, for example +5 for phosphorus in PCl₅.
• Valence (outer-shell) electrons: the electrons in the outer shell of an atom, the 3s and 3p electrons for a Period 3 element; these are the electrons available for bonding.
• Basic oxide: an oxide that reacts with an acid to form a salt and water but does not react with alkalis; in Period 3, Na₂O and MgO.
• Acidic oxide: an oxide that reacts with an alkali such as sodium hydroxide to form a salt and water; if it dissolves in water, the solution is acidic. In Period 3, SiO₂, P₄O₁₀, SO₂ and SO₃.
• Amphoteric: describes an oxide or hydroxide that reacts with both acids and bases, so it can act as a base and as an acid; in Period 3, Al₂O₃ and Al(OH)₃.
• Hydrolysis: a reaction in which water breaks bonds in a compound and new products form; covalent chlorides such as SiCl₄ are hydrolysed, forming HCl. Ionic NaCl only dissolves, so it is not hydrolysed.

Content

Electronegativity sets the pattern

Electronegativity increases across Period 3, because the nuclear charge increases and the atoms get smaller while the shielding stays almost the same (topic 3.1). Oxygen and chlorine are more electronegative than every element from sodium to sulfur, so in every oxide and chloride on this page the Period 3 element has a positive oxidation number.

Each element's Pauling electronegativity (Data Booklet) is in brackets, with its difference from oxygen (3.5) and from chlorine (3.0):

Element	Difference from O	Difference from Cl
Na (0.9)	2.6	2.1
Mg (1.3)	2.2	1.7
Al (1.5)	2.0	1.5
Si (1.9)	1.6	1.1
P (2.2)	1.3	0.8
S (2.6)	0.9	not needed

The larger the difference, the more ionic the bonding. The difference falls across the period, so the bonding changes from ionic to covalent: between Al and Si for the oxides, and between Mg and Al for the chlorides. Because oxygen is more electronegative than chlorine, Al₂O₃ (difference 2.0) is ionic while aluminium chloride (difference 1.5) is covalent.

Reactions of the elements with oxygen

Each element from sodium to sulfur burns in oxygen when heated (white phosphorus needs no heating). The element is oxidised: its oxidation number rises from 0.

Element	What you see
Sodium	bright yellow-orange flame; white solid
Magnesium	brilliant white flame; white powder
Aluminium	the powder burns with bright white sparks; white solid
Silicon	reacts only when heated strongly; white solid
Phosphorus	bright flame; dense white smoke
Sulfur	blue flame; colourless gas with a choking smell

4Na(s) + O₂(g) → 2Na₂O(s)

2Mg(s) + O₂(g) → 2MgO(s)

4Al(s) + 3O₂(g) → 2Al₂O₃(s)

Si(s) + O₂(g) → SiO₂(s)

P₄(s) + 5O₂(g) → P₄O₁₀(s)

S(s) + O₂(g) → SO₂(g)

• Aluminium is burned as a powder, because on a sheet a thin layer of Al₂O₃ protects the metal.
• Write P₄O₁₀ (molecular formula), not P₂O₅ (empirical formula). With too little oxygen, phosphorus forms P₄O₆; in plenty of oxygen, sodium also forms some peroxide. You need only P₄O₁₀ and Na₂O.
• Burning sulfur gives SO₂, not SO₃; SO₃ is made from SO₂ and oxygen over a catalyst (the Contact process, topic 7.1).

Reactions of the elements with chlorine

Each element reacts with chlorine on heating, and is again oxidised.

Element	What you see
Sodium	burns with a yellow-orange flame; white solid
Magnesium	burns with a white flame; white solid
Aluminium	white or pale yellow solid that turns to vapour easily
Silicon	colourless liquid
Phosphorus	with excess chlorine, an off-white or pale yellow solid (with too little, liquid PCl₃)

2Na(s) + Cl₂(g) → 2NaCl(s)

Mg(s) + Cl₂(g) → MgCl₂(s)

2Al(s) + 3Cl₂(g) → 2AlCl₃(s)

Si(s) + 2Cl₂(g) → SiCl₄(l)

P₄(s) + 10Cl₂(g) → 4PCl₅(s)

For phosphorus, 2P(s) + 5Cl₂(g) → 2PCl₅(s) is also correct. The chlorine must be dry, because the chlorides of aluminium, silicon and phosphorus react with water.

Sodium and magnesium with water

Sodium reacts vigorously with cold water. It floats, melts into a shiny ball, fizzes and moves about until it disappears, leaving a strongly alkaline solution (pH about 13 to 14).

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Magnesium reacts very slowly with cold water: a few bubbles collect on the metal over several days. Magnesium hydroxide is only slightly soluble, so the solution is weakly alkaline (pH about 9 to 10).

Mg(s) + 2H₂O(l) → Mg(OH)₂(s) + H₂(g)

Heated magnesium burns in steam with a bright white flame, and the product is the oxide:

Mg(s) + H₂O(g) → MgO(s) + H₂(g)

Oxidation numbers in the oxides and chlorides

In these compounds oxygen is −2 and chlorine is −1, and the oxidation numbers in a compound add up to zero. For P₄O₁₀: 4x + 10(−2) = 0, so x = +5.

The table gives the oxidation number of the Period 3 element, with its outer electrons:

Element	Oxide	Chloride
Na: 3s¹	Na₂O: +1	NaCl: +1
Mg: 3s²	MgO: +2	MgCl₂: +2
Al: 3s² 3p¹	Al₂O₃: +3	AlCl₃: +3
Si: 3s² 3p²	SiO₂: +4	SiCl₄: +4
P: 3s² 3p³	P₄O₁₀: +5	PCl₅: +5
S: 3s² 3p⁴	SO₃: +6; SO₂: +4	not needed

Explaining the trend:

• The highest oxidation number equals the number of outer-shell (valence) electrons, the 3s and 3p electrons, because all of them can be used in bonding: given to the more electronegative O or Cl in an ionic compound, or shared with it in a covalent one. Each step across the period adds one outer electron, so the highest value rises by one each time, up to +5 (PCl₅) for the chlorides and +6 (SO₃) for the oxides.
• Inner-shell electrons are not used. For magnesium, the third ionisation energy (7740 kJ mol⁻¹) is more than five times the second (1450 kJ mol⁻¹), because the third electron would come from the full second shell, closer to the nucleus and less shielded.
• In SO₂, sulfur uses only four of its six outer electrons in bonding; the other two stay as a lone pair, so it is +4.

Use oxidation numbers to check products. When an oxide or chloride reacts with water, an acid or NaOH(aq), no oxidation number changes: these are acid-base or hydrolysis reactions, not redox. So SO₂ (S +4) and water give sulfurous acid, H₂SO₃ (S +4), not H₂SO₄ (S +6). PCl₅ and P₄O₁₀ (P +5) give phosphoric(V) acid, H₃PO₄ (P +5), never an oxide of phosphorus.

The oxides with water

Oxide	With water	pH
Na₂O	reacts readily, giving out heat	13 to 14
MgO	only a little reacts	about 9 to 10
Al₂O₃	insoluble: no reaction	7 (water unchanged)
SiO₂	insoluble: no reaction	7 (water unchanged)
P₄O₁₀	reacts violently	about 1 to 2
SO₂	dissolves and reacts	about 1 to 3
SO₃	reacts violently	about 0 to 1

Na₂O(s) + H₂O(l) → 2NaOH(aq)

MgO(s) + H₂O(l) → Mg(OH)₂(s)

P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)

SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq)

SO₃(g) + H₂O(l) → H₂SO₄(aq)

These pH values are typical; the exact value depends on how much oxide dissolves. Learn the pattern: strongly alkaline (NaOH is very soluble), weakly alkaline (Mg(OH)₂ is only slightly soluble), neutral (insoluble oxides), then acidic. Sulfurous acid and phosphoric(V) acid are weak acids; sulfuric acid is strong.

Acid-base behaviour of the oxides and hydroxides

• Basic (Na₂O, MgO, NaOH, Mg(OH)₂): neutralise acids; they show no acidic behaviour with NaOH(aq).
• Amphoteric (Al₂O₃, Al(OH)₃): react with acids and with NaOH(aq).
• Acidic (P₄O₁₀, SO₂, SO₃): neutralise NaOH(aq); they do not react with acids.

Basic oxides and hydroxides give a salt and water with an acid:

Na₂O(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l)

MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l)

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

Aluminium oxide and hydroxide react with both. Al₂O₃ reacts slowly, so hot acid or hot, concentrated NaOH(aq) is used. The product with NaOH is sodium aluminate, NaAl(OH)₄, whose anion is [Al(OH)₄]⁻.

Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)

Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq)

Al(OH)₃(s) + 3HCl(aq) → AlCl₃(aq) + 3H₂O(l)

Al(OH)₃(s) + NaOH(aq) → NaAl(OH)₄(aq)

Acidic oxides give a salt and water with NaOH(aq):

P₄O₁₀(s) + 12NaOH(aq) → 4Na₃PO₄(aq) + 6H₂O(l)

SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l)

SO₃(g) + 2NaOH(aq) → Na₂SO₄(aq) + H₂O(l)

SiO₂ is acidic too, but it needs hot, concentrated NaOH(aq) to react at a useful rate (forming sodium silicate, Na₂SiO₃); its equation is not required.

Worked example 1: aluminium hydroxide and excess sodium hydroxide

Question: 0.400 mol dm⁻³ NaOH(aq) is added from a burette to 20.0 cm³ of 0.150 mol dm⁻³ AlCl₃(aq). (a) Describe what you see. (b) What volume of NaOH(aq) has been added when the amount of precipitate is greatest? (c) What volume has been added when the precipitate has just all dissolved?

Step 1: amount of aluminium ions. $$\begin{aligned} n(\text{Al}^{3+}) &= 0.150 \times \frac{20.0}{1000} \\ &= 3.00 \times 10^{-3}\ \text{mol} \end{aligned}$$

Step 2: forming the precipitate. Al³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) needs 3 mol of OH⁻ per mole of Al³⁺: $$\begin{aligned} n(\text{OH}^-) &= 3 \times 3.00 \times 10^{-3} \\ &= 9.00 \times 10^{-3}\ \text{mol} \end{aligned}$$ $$V = \frac{9.00 \times 10^{-3}}{0.400} = 0.0225\ \text{dm}^3$$

Step 3: dissolving it. Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) uses one more mole of OH⁻ per mole of Al, so 4 mol in total: $$\begin{aligned} n(\text{OH}^-) &= 4 \times 3.00 \times 10^{-3} \\ &= 1.20 \times 10^{-2}\ \text{mol} \end{aligned}$$ $$V = \frac{1.20 \times 10^{-2}}{0.400} = 0.0300\ \text{dm}^3$$

Answer: (a) a white precipitate forms, then dissolves in excess NaOH(aq) to leave a colourless solution. (b) 22.5 cm³. (c) 30.0 cm³. With 20.0 cm³ of 0.150 mol dm⁻³ MgCl₂(aq) instead, a white precipitate of Mg(OH)₂ would be complete at 15.0 cm³ and would not dissolve in excess NaOH(aq), because Mg(OH)₂ is basic, not amphoteric.

The chlorides with water

Chloride	With water	pH
NaCl	dissolves; no reaction	7
MgCl₂	dissolves; very slight reaction	about 6.5
AlCl₃	reacts, giving out heat; misty fumes of HCl if there is little water	about 3
SiCl₄	reacts vigorously: misty fumes of HCl and a white solid	about 1 to 2
PCl₅	reacts vigorously: misty fumes of HCl and a colourless solution	about 1 to 2

Ionic chlorides. Water molecules surround the separated ions:

NaCl(s) → Na⁺(aq) + Cl⁻(aq)

MgCl₂(s) → Mg²⁺(aq) + 2Cl⁻(aq)

NaCl does not react: no NaOH or HCl forms. The Mg²⁺ ion is smaller than Na⁺ and has a 2+ charge, so it pulls slightly on the electrons of nearby water molecules and a few of them release H⁺, making the solution very slightly acidic.

Aluminium chloride. In water, each Al³⁺ ion bonds to six water molecules:

AlCl₃(s) + 6H₂O(l) → [Al(H₂O)₆]³⁺(aq) + 3Cl⁻(aq)

Al³⁺ is small and has a 3+ charge, so it pulls electron density strongly from the O-H bonds of these water molecules, and some release H⁺:

[Al(H₂O)₆]³⁺(aq) ⇌ [Al(H₂O)₅(OH)]²⁺(aq) + H⁺(aq)

With only a little water the reaction goes further and HCl gas escapes; a simplified equation is AlCl₃(s) + 3H₂O(l) → Al(OH)₃(s) + 3HCl(g).

Silicon(IV) chloride and phosphorus(V) chloride. These are covalent. Chlorine is more electronegative than Si or P, so each Si-Cl or P-Cl bond is polar and the central atom is δ+. A lone pair on the oxygen of a water molecule is attracted to it, HCl is lost, and this repeats until no chlorine is left:

SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq)

PCl₅(s) + 4H₂O(l) → H₃PO₄(aq) + 5HCl(aq)

With only a little water, PCl₅ reacts only as far as POCl₃: PCl₅(s) + H₂O(l) → POCl₃(l) + 2HCl(g).

Worked example 2: hydrolysis of silicon(IV) chloride

Question: 1.19 g of SiCl₄ is added to an excess of water. (a) Give two observations and the equation. (b) Calculate the mass of white solid, after it is filtered off and heated to constant mass. (c) Calculate the volume of 0.500 mol dm⁻³ NaOH(aq) that neutralises the filtrate, assuming no HCl escapes.

Step 1: amount of SiCl₄. $M_{\text{r}}$ = 28.1 + (4 × 35.5) = 170.1 $$n(\text{SiCl}_4) = \frac{1.19}{170.1} = 6.996 \times 10^{-3}\ \text{mol}$$

Step 2: mass of SiO₂. One mole of SiCl₄ gives one mole of SiO₂, and $M_{\text{r}}$(SiO₂) = 28.1 + (2 × 16.0) = 60.1: $$m = 6.996 \times 10^{-3} \times 60.1 = 0.420\ \text{g}$$

Step 3: volume of NaOH(aq). One mole of SiCl₄ gives four moles of HCl, which react 1 : 1 with NaOH: $$\begin{aligned} n(\text{HCl}) &= 4 \times 6.996 \times 10^{-3} \\ &= 0.02798\ \text{mol} \end{aligned}$$ $$V = \frac{0.02798}{0.500} = 0.0560\ \text{dm}^3$$

Answer: (a) misty fumes and a white solid; the mixture also warms up. SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq). (b) 0.420 g of SiO₂. (c) 56.0 cm³. If some HCl escaped as fumes, less NaOH(aq) would be needed.

Explaining the trends: bonding and electronegativity

Oxide	Structure	Nature
Na₂O	giant ionic	basic
MgO	giant ionic	basic
Al₂O₃	giant ionic	amphoteric
SiO₂	giant covalent	acidic
P₄O₁₀	simple molecular	acidic
SO₂, SO₃	simple molecular	acidic

Chloride	Structure	pH in water
NaCl	giant ionic	7 (dissolves)
MgCl₂	giant ionic	about 6.5 (dissolves)
AlCl₃	simple molecular	about 3 (hydrolysed)
SiCl₄	simple molecular	1 to 2 (hydrolysed)
PCl₅	simple molecular	1 to 2 (hydrolysed)

• Ionic oxides are basic. Na and Mg have low electronegativities, so they give their outer electrons to oxygen, forming oxide ions, O²⁻. An oxide ion accepts protons: with water it forms hydroxide ions, O²⁻(s) + H₂O(l) → 2OH⁻(aq), and with an acid it forms water. NaOH and Mg(OH)₂ are basic for the same reason: their OH⁻ ions accept H⁺.
• Covalent oxides are acidic. P and S are close enough to oxygen in electronegativity to share electrons with it. With water they form acids with O-H groups bonded to the P or S atom, which pulls electron density away from the O-H bonds, so H⁺ is released. Sulfur at +6 (SO₃) gives a strong acid; at +4 (SO₂), a weak one.
• Aluminium is in between. Al₂O₃ is ionic, but Al³⁺ is small and highly charged. Its oxide ions accept H⁺ from acids (basic behaviour), and its Al³⁺ ions bond to OH⁻ ions from NaOH(aq) to form [Al(OH)₄]⁻ (acidic behaviour), so it is amphoteric. Al(OH)₃ behaves in the same way.
• Insoluble oxides. The attraction between the highly charged Al³⁺ and O²⁻ ions in Al₂O₃ is very strong, and SiO₂ is held together by strong Si-O bonds throughout its giant structure. Water cannot break either apart.
• Chlorides. As the electronegativity difference falls, the chlorides change from ionic, which dissolve with little or no reaction, to covalent, which are hydrolysed and release H⁺. So the pH falls across the period.

Describe AlCl₃ and PCl₅ as covalent and simple molecular: both turn to vapour easily and are hydrolysed like the other covalent chlorides. Heated aluminium chloride forms Al₂Cl₆ molecules, in which two AlCl₃ units are joined by two dative bonds (topic 3.4). The solids have more complicated structures, which you are not asked about.

Suggesting the bonding from properties

Evidence	Suggests
High melting point	giant structure, ionic or covalent
Low melting or boiling point, or turns to vapour easily	simple molecular (covalent)
Conducts electricity when molten, but not as a solid	ionic
Does not conduct when molten or liquid	covalent
Chloride gives a neutral or almost neutral solution	ionic chloride
Chloride fumes in moist air and gives a strongly acidic solution	covalent chloride
Oxide gives an alkaline solution or reacts only with acids	ionic oxide
Oxide gives an acidic solution or reacts with NaOH(aq)	covalent oxide (or amphoteric Al₂O₃, if it also reacts with acids)

Do not use the conductivity of a solution: covalent chlorides react with water to form H⁺ and Cl⁻ ions, so their solutions conduct too.

Worked example 3: suggesting bonding from observations

Question: E, G and L are each an oxide or a chloride of a Period 3 element. Suggest the bonding and structure of each, and identify it.

• E: white solid with a very high melting point that conducts electricity when molten; shaken with water, it gives a solution of pH about 9 to 10; it reacts with dilute hydrochloric acid but not with NaOH(aq).
• G: pale yellow solid that turns to vapour on gentle heating; reacts vigorously with water, giving misty fumes and a colourless solution of pH about 1. This solution gives a white precipitate with aqueous silver nitrate, but no precipitate when NaOH(aq) is added drop by drop.
• L: white solid with a high melting point that conducts when molten; dissolves in water without fuming to give a solution of pH 7; no precipitate with NaOH(aq).

E. A very high melting point and conducting when molten mean giant ionic. Reacting with acid but not with NaOH(aq) means a basic oxide. Na₂O would give pH 13 to 14, and Al₂O₃ would also react with NaOH(aq), so E is MgO.

G. Turning to vapour easily means simple molecules, so covalent bonding. The white precipitate with silver nitrate shows chloride ions, so G is a chloride, not an acidic oxide such as SO₃. Fumes and a strongly acidic solution mean a covalent chloride hydrolysed to HCl. Aluminium chloride would give a white precipitate of Al(OH)₃ with NaOH(aq); SiCl₄ is a liquid and gives a white solid with water. So G is PCl₅.

L. A high melting point and conducting when molten mean giant ionic. A neutral solution fits an ionic chloride (an ionic oxide would be alkaline), and no precipitate with NaOH(aq) rules out MgCl₂. So L is NaCl.

Answer: E is MgO (giant ionic, basic); G is PCl₅ (covalent, simple molecular); L is NaCl (giant ionic).

Key Equations

Oxidation numbers (in the compounds on this page O = −2, Cl = −1 and H = +1; the sum equals the overall charge, which is zero for a compound): $$\sum \text{oxidation numbers} = \text{overall charge}$$

Amount from mass ($m$ = mass in g; $M$ = molar mass in g mol⁻¹; $n$ in mol): $$n = \frac{m}{M}$$

Amount in a solution ($c$ = concentration in mol dm⁻³; $V$ = volume in cm³): $$n = c \times \frac{V}{1000}$$

Common Mistakes to Avoid

• Mistake: saying NaCl gives an acidic solution, or writing NaCl + H₂O → NaOH + HCl. Fix: NaCl only dissolves, NaCl(s) → Na⁺(aq) + Cl⁻(aq), so the pH is 7. The chlorides that give clearly acidic solutions are AlCl₃ (pH about 3), SiCl₄ and PCl₅ (pH about 1 to 2); do not leave out AlCl₃.
• Mistake: changing an oxidation number when an oxide or chloride reacts with water. Fix: these reactions are not redox. SO₂ gives H₂SO₃ (S stays +4), not H₂SO₄; PCl₅ gives H₃PO₄ and HCl (P stays +5), not an oxide of phosphorus.
• Mistake: calling SiCl₄ giant covalent, or giving a wrong formula for the white solid it forms with water. Fix: SiCl₄ is a simple molecular liquid. The giant covalent compound is SiO₂, and that is the white solid: SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq).
• Mistake: writing the steam reaction when asked about magnesium and cold water, or calling the solution acidic. Fix: cold water gives Mg(OH)₂ and H₂ very slowly, and the solution is weakly alkaline (pH about 9 to 10); steam gives MgO and H₂.
• Mistake: explaining the rise in oxidation number by "more protons", or by "more electrons" without saying which electrons. Fix: each element has one more outer-shell (3s and 3p) electron, and all of them can be used in bonding to the more electronegative O or Cl. Phosphorus, 3s² 3p³, reaches +5 in PCl₅.
• Mistake: giving Al(OH)₃ or AlCl₃ as the product of Al₂O₃ with NaOH(aq), or saying Mg(OH)₂ dissolves in excess NaOH(aq). Fix: Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq). Of the Period 3 oxides and hydroxides here, only aluminium's are amphoteric; Mg(OH)₂ stays as a white precipitate.
• Mistake: deciding the bonding in a chloride from the conductivity of its solution. Fix: use the pure compound, molten or liquid. SiCl₄ gives a conducting solution because it reacts to form H⁺ and Cl⁻ ions, yet SiCl₄ itself is covalent.

Exam Tips

• Describe what you would see: flame colour, fizzing, misty fumes, a white solid or precipitate, a colourless solution. "Misty fumes" is an observation; "HCl is formed" is not. Write "white precipitate", not "white solution", and "colourless", not "clear".
• Write equations balanced and with state symbols, using P₄O₁₀, Cl₂ and NaAl(OH)₄. In names, give the oxidation number in Roman numerals, for example phosphorus(V) chloride for PCl₅.
• State an oxidation number with its sign (+4 for S in SO₂), and a pH as a number or narrow range (pH 1 to 2 for SiCl₄ in water), not just "acidic".
• Explain a trend as a chain: electronegativity difference, then type of bonding, then behaviour. For example, Na (0.9) and O (3.5) differ by 2.6, so Na₂O is ionic; its O²⁻ ions accept H⁺, so it is basic.
• Suggest the bonding by naming both the bonding and the structure (giant ionic, giant covalent or simple molecular) and quoting the evidence you used.