Period 3 Oxides, Hydroxides and Chlorides

AS · 14 min

Put the oxides and chlorides of the Period 3 elements into water and you see the whole period in miniature: on the left, ionic compounds that give alkaline or neutral solutions; in the middle, aluminium compounds that can behave as acids or bases; on the right, covalent compounds that react with water to give strongly acidic solutions. This note covers every reaction the syllabus lists, the likely pH of each solution, amphoteric behaviour, and, most importantly, how to explain all of it using bonding and electronegativity. It is one of the most heavily examined areas of AS inorganic chemistry, in both structured Paper 2 questions and Paper 1 multiple choice.

The organising idea: electronegativity and bonding

Electronegativity increases across Period 3 (Na 0.9 to Cl 3.0). Oxygen (3.5) and chlorine (3.0) are both highly electronegative.

Left of the period (Na, Mg): the electronegativity difference between the element and O or Cl is large, so the oxides and chlorides are ionic.
Aluminium: the difference is intermediate. $\ce{Al2O3}$ is ionic with some covalent character; aluminium chloride is essentially covalent.
Right of the period (Si, P, S): the difference is small, so the oxides and chlorides are covalent. $\ce{SiO2}$ is giant covalent; the others are simple molecular.

Ionic oxides contain the oxide ion, $\ce{O^2-}$ , which is a strong base. Covalent oxides of non-metals react with water to form acids. So the oxides go from basic to amphoteric to acidic across the period.

oxide	$\ce{Na2O}$	$\ce{MgO}$	$\ce{Al2O3}$	$\ce{SiO2}$	$\ce{P4O10}$	$\ce{SO2}$ / $\ce{SO3}$
bonding	ionic	ionic	ionic with covalent character	covalent	covalent	covalent
structure	giant ionic	giant ionic	giant ionic	giant covalent	simple molecular	simple molecular
melting point (approximate)	about 1400 K	about 3100 K	about 2300 K	about 2000 K	about 600 K (sublimes)	$\ce{SO2}$ 200 K, $\ce{SO3}$ 290 K
nature	basic	basic	amphoteric	acidic	acidic	acidic

Oxides with water

oxide	reaction with water	likely pH
$\ce{Na2O}$	$\ce{Na2O(s) + H2O(l) -> 2NaOH(aq)}$ ; dissolves exothermically	13–14
$\ce{MgO}$	$\ce{MgO(s) + H2O(l) -> Mg(OH)2(s)}$ ; only slightly soluble	about 9
$\ce{Al2O3}$	no reaction; insoluble	7
$\ce{SiO2}$	no reaction; insoluble	7
$\ce{P4O10}$	$\ce{P4O10(s) + 6H2O(l) -> 4H3PO4(aq)}$ ; violent reaction	about 2 (0–2)
$\ce{SO2}$	$\ce{SO2(g) + H2O(l) <=> H2SO3(aq)}$ ; sulfurous acid, a weak acid	about 2–3
$\ce{SO3}$	$\ce{SO3(g) + H2O(l) -> H2SO4(aq)}$ ; violent, misty fumes	about 0–1

Why each happens:

$\ce{Na2O}$ and $\ce{MgO}$ contain $\ce{O^2-}$ ions, which accept protons from water: $\ce{O^2-(aq) + H2O(l) -> 2OH-(aq)}$ . $\ce{NaOH}$ is very soluble so the pH is very high; $\ce{Mg(OH)2}$ is only sparingly soluble so the $\ce{[OH-]}$ is low and the pH is only about 9.
$\ce{Al2O3}$ has a very strong lattice (small, highly charged $\ce{Al^3+}$ and $\ce{O^2-}$ ), so it is insoluble and does not react; the water stays at pH 7.
$\ce{SiO2}$ is a giant covalent structure; water cannot break its strong Si–O bonds, so there is no reaction.
$\ce{P4O10}$ , $\ce{SO2}$ and $\ce{SO3}$ are covalent molecules with electron-deficient central atoms bonded to electronegative oxygen. Water molecules attack them, forming acids that release $\ce{H+}$ ions.

Watch out

" $\ce{Al2O3}$ and $\ce{SiO2}$ are neutral" is wrong. They have no effect on the pH of water because they are insoluble, but chemically $\ce{Al2O3}$ is amphoteric and $\ce{SiO2}$ is acidic. Say "insoluble, so pH 7" and give their nature separately.

Acid–base behaviour of the oxides and hydroxides

A basic oxide reacts with acids to form a salt and water. An acidic oxide reacts with bases (alkalis) to form a salt and water. An amphoteric oxide or hydroxide reacts with both acids and bases.

Definition

An amphoteric substance is one that can react both as an acid and as a base.

Basic: sodium and magnesium

\ce{Na2O(s) + 2HCl(aq) -> 2NaCl(aq) + H2O(l)}

\ce{MgO(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l)}

The hydroxides behave the same way:

\ce{NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l)}

\ce{Mg(OH)2(s) + 2HCl(aq) -> MgCl2(aq) + 2H2O(l)}

Neither $\ce{Na2O}$ nor $\ce{MgO}$ reacts with sodium hydroxide.

Amphoteric: aluminium

Aluminium oxide and aluminium hydroxide react with acids and with sodium hydroxide.

With acid (acting as a base):

\ce{Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)}

\ce{Al(OH)3(s) + 3HCl(aq) -> AlCl3(aq) + 3H2O(l)}

With sodium hydroxide (acting as an acid), forming the soluble aluminate ion, $\ce{[Al(OH)4]-}$ :

\ce{Al2O3(s) + 2NaOH(aq) + 3H2O(l) -> 2NaAl(OH)4(aq)}

\ce{Al(OH)3(s) + NaOH(aq) -> NaAl(OH)4(aq)}

or, as an ionic equation, $\ce{Al(OH)3(s) + OH-(aq) -> [Al(OH)4]-(aq)}$ .

Tip

Some books write the product with hot, concentrated sodium hydroxide as sodium aluminate, $\ce{NaAlO2}$ : $\ce{Al2O3 + 2NaOH -> 2NaAlO2 + H2O}$ . Both forms are accepted, as long as the equation balances. $\ce{NaAl(OH)4}$ is the more common form in mark schemes.

This amphoteric behaviour is the basis of the qualitative-analysis test for $\ce{Al^3+}$ : adding $\ce{NaOH(aq)}$ gives a white precipitate of $\ce{Al(OH)3}$ which dissolves in excess, while $\ce{Mg^2+}$ gives a white precipitate that does not.

Acidic: silicon, phosphorus and sulfur

The acidic oxides react with sodium hydroxide to form salts and water.

\ce{SiO2(s) + 2NaOH(aq) -> Na2SiO3(aq) + H2O(l)}

(only with hot, concentrated $\ce{NaOH}$ ; silicon dioxide is a very weakly acidic oxide)

\ce{P4O10(s) + 12NaOH(aq) -> 4Na3PO4(aq) + 6H2O(l)}

\ce{SO2(g) + 2NaOH(aq) -> Na2SO3(aq) + H2O(l)}

\ce{SO3(g) + 2NaOH(aq) -> Na2SO4(aq) + H2O(l)}

None of these oxides reacts with hydrochloric acid.

Key result

oxide	reacts with $\ce{HCl}$ ?	reacts with $\ce{NaOH}$ ?	nature
$\ce{Na2O}$ , $\ce{MgO}$	yes	no	basic
$\ce{Al2O3}$	yes	yes	amphoteric
$\ce{SiO2}$ , $\ce{P4O10}$ , $\ce{SO2}$ , $\ce{SO3}$	no	yes	acidic

Chlorides with water

The chlorides split into two groups: ionic chlorides that simply dissolve, and covalent chlorides that are hydrolysed (react with water, breaking bonds), producing hydrogen chloride and an acidic solution.

chloride	bonding and structure	what happens in water	likely pH
$\ce{NaCl}$	ionic, giant lattice	dissolves; ions hydrated: $\ce{NaCl(s) -> Na+(aq) + Cl-(aq)}$	7
$\ce{MgCl2}$	ionic, giant lattice	dissolves; very slight hydrolysis	about 6.5
$\ce{AlCl3}$	covalent ( $\ce{Al2Cl6}$ )	hydrolysed; steamy fumes of $\ce{HCl}$ with a little water	about 3
$\ce{SiCl4}$	covalent, simple molecular	hydrolysed completely; white solid and steamy fumes	about 1–2
$\ce{PCl5}$	covalent, simple molecular	hydrolysed completely; steamy fumes	about 1–2

Sodium chloride and magnesium chloride

Sodium chloride dissolves to give a neutral solution. Magnesium chloride also dissolves, but the small, doubly charged $\ce{Mg^2+}$ ion attracts the lone pairs of surrounding water molecules strongly enough to weaken a few O–H bonds and release a small number of $\ce{H+}$ ions. The solution is very slightly acidic, about pH 6.5.

Aluminium chloride

Aluminium chloride reacts exothermically with water. In a large amount of water the $\ce{Al^3+}$ ion becomes hydrated, $\ce{[Al(H2O)6]^3+}$ . Because $\ce{Al^3+}$ is small and highly charged (high charge density), it pulls electron density from the O–H bonds of the water molecules around it so strongly that the complex loses $\ce{H+}$ :

\ce{AlCl3(s) + 6H2O(l) -> [Al(H2O)6]^3+(aq) + 3Cl-(aq)}

\ce{[Al(H2O)6]^3+(aq) <=> [Al(H2O)5(OH)]^2+(aq) + H+(aq)}

The solution has a pH of about 3. With only a little water (or in moist air), hydrolysis goes further and steamy fumes of hydrogen chloride are seen:

\ce{AlCl3(s) + 3H2O(l) -> Al(OH)3(s) + 3HCl(g)}

Silicon tetrachloride

\ce{SiCl4(l) + 2H2O(l) -> SiO2(s) + 4HCl(aq)}

A vigorous reaction: steamy (white, misty) fumes of $\ce{HCl}$ and a white solid of hydrated silicon dioxide. The hydrochloric acid formed gives a pH of about 1–2.

Phosphorus pentachloride

\ce{PCl5(s) + 4H2O(l) -> H3PO4(aq) + 5HCl(aq)}

A vigorous reaction with steamy fumes of $\ce{HCl}$ ; two acids form, so the pH is about 1–2. With only a little water, phosphorus oxychloride forms first: $\ce{PCl5 + H2O -> POCl3 + 2HCl}$ .

Tip

Extension, not required: carbon tetrachloride, $\ce{CCl4}$ , does not react with water, even though $\ce{SiCl4}$ does. Silicon is larger and has empty orbitals in its third shell that can accept a lone pair from a water molecule, so water can attack it. Carbon in $\ce{CCl4}$ is small, crowded by four chlorine atoms, and has no low-energy empty orbital. This illustrates that hydrolysis needs a route for water to attack, not just polar bonds.

Explaining the trends

The syllabus asks you to explain all of these trends in terms of bonding and electronegativity. One chain of reasoning covers everything.

Key result

Across Period 3 the electronegativity of the element increases, so the difference in electronegativity between the element and oxygen (or chlorine) decreases.
So bonding changes from ionic ( $\ce{Na2O}$ , $\ce{MgO}$ , $\ce{NaCl}$ , $\ce{MgCl2}$ ) through intermediate ( $\ce{Al2O3}$ , aluminium chloride) to covalent ( $\ce{SiO2}$ , $\ce{P4O10}$ , $\ce{SO2}$ , $\ce{SO3}$ , $\ce{SiCl4}$ , $\ce{PCl5}$ ).
Ionic oxides contain $\ce{O^2-}$ , which reacts with water or acids as a base: basic oxides.
Covalent oxides react with water to give acids, or react with alkalis: acidic oxides.
Ionic chlorides dissolve to give hydrated ions: neutral (or nearly neutral) solutions.
Covalent chlorides are hydrolysed to give $\ce{HCl}$ : acidic solutions.
The oxidation number rises because more outer electrons are available for bonding (previous note).

Suggesting bonding from observations

Exam questions often describe an unknown oxide or chloride and ask you to deduce its bonding and structure. Use the evidence like this:

observation	suggests
high melting point; conducts when molten or in solution, not when solid	giant ionic
very high melting point; does not conduct even when molten; insoluble	giant covalent
low melting or boiling point; does not conduct when molten	simple molecular (covalent)
fumes in moist air; reacts with water to give an acidic solution and $\ce{HCl}$	covalent chloride
dissolves in water to give a neutral solution	ionic chloride
oxide reacts with acids but not alkalis	ionic, basic
oxide or hydroxide reacts with both acids and alkalis	amphoteric (Al)

Worked examples

Oxides with water

Write equations for the reactions of sodium oxide and of sulfur trioxide with water, and give the approximate pH of each solution.

Solution

$\ce{Na2O(s) + H2O(l) -> 2NaOH(aq)}$ , pH 13–14 (strongly alkaline).

$\ce{SO3(g) + H2O(l) -> H2SO4(aq)}$ , pH 0–1 (strongly acidic).

Sodium oxide is ionic; its $\ce{O^2-}$ ions accept protons from water to form $\ce{OH-}$ . Sulfur trioxide is covalent; it reacts with water to form sulfuric acid, which releases $\ce{H+}$ .

Showing amphoteric behaviour

Aluminium hydroxide is amphoteric. Write two equations to show this, and describe what is seen when sodium hydroxide solution is added dropwise until in excess to a solution containing $\ce{Al^3+}$ ions.

Solution

As a base: $\ce{Al(OH)3(s) + 3HCl(aq) -> AlCl3(aq) + 3H2O(l)}$

As an acid: $\ce{Al(OH)3(s) + NaOH(aq) -> NaAl(OH)4(aq)}$

Observation: a white precipitate forms first ( $\ce{Al^3+(aq) + 3OH-(aq) -> Al(OH)3(s)}$ ), which then dissolves in excess sodium hydroxide to give a colourless solution.

Contrasting two chlorides

Describe and explain the difference between what happens when sodium chloride and when silicon tetrachloride are added to water. Include equations and pH values.

Solution

Sodium chloride is ionic. It dissolves; the ions separate and become hydrated, and there is no reaction with water. $\ce{NaCl(s) -> Na+(aq) + Cl-(aq)}$ . The solution is neutral, pH 7.

Silicon tetrachloride is covalent (simple molecular). It is hydrolysed: water attacks the silicon atom, the Si–Cl bonds break and hydrogen chloride forms. Steamy fumes and a white solid are seen. $\ce{SiCl4(l) + 2H2O(l) -> SiO2(s) + 4HCl(aq)}$ . The solution is strongly acidic, pH 1–2.

The difference arises because the electronegativity difference between Na and Cl is large (ionic bonding) but between Si and Cl is small (covalent bonding).

Exam-style: titrating the products of hydrolysis

$0.425\ \text{g}$ of silicon tetrachloride is added to water and the mixture is made up to $250\ \text{cm}^3$ . A $25.0\ \text{cm}^3$ portion is titrated with $0.0500\ \text{mol dm}^{-3}$ sodium hydroxide. Calculate the expected titre. ( $M(\ce{SiCl4}) = 170.1\ \text{g mol}^{-1}$ ; assume silicon dioxide does not react with the alkali.)

Solution

$n(\ce{SiCl4}) = \dfrac{0.425}{170.1} = 2.499 \times 10^{-3}\ \text{mol}$

$\ce{SiCl4 + 2H2O -> SiO2 + 4HCl}$ , so $n(\ce{HCl}) = 4 \times 2.499 \times 10^{-3} = 9.994 \times 10^{-3}\ \text{mol}$ in $250\ \text{cm}^3$ .

In $25.0\ \text{cm}^3$ : $9.994 \times 10^{-4}\ \text{mol}$ .

$\ce{HCl + NaOH -> NaCl + H2O}$ ( $1 : 1$ ), so $n(\ce{NaOH}) = 9.994 \times 10^{-4}\ \text{mol}$ .

V = \frac{9.994 \times 10^{-4}}{0.0500} = 0.01999\ \text{dm}^3 = 20.0\ \text{cm}^3

Exam-hard: deducing an element from its compounds

Element E is in Period 3. Its oxide has a very high melting point and is insoluble in water, but it dissolves in both dilute hydrochloric acid and hot aqueous sodium hydroxide. Its chloride sublimes at about 450 K, and a solution of the chloride in water has a pH of about 3. Identify E, explain each observation, and write equations for the reactions of the oxide.

Solution

E is aluminium.

Very high melting point oxide: $\ce{Al2O3}$ has a giant ionic lattice (with some covalent character) and strong attraction between small, highly charged $\ce{Al^3+}$ and $\ce{O^2-}$ ions.
Insoluble in water: the lattice is too strong to be broken by hydration, so the pH of water is unchanged.
Dissolves in both acid and alkali: the oxide is amphoteric. $\ce{Al2O3(s) + 6HCl(aq) -> 2AlCl3(aq) + 3H2O(l)}$ $\ce{Al2O3(s) + 2NaOH(aq) + 3H2O(l) -> 2NaAl(OH)4(aq)}$
Chloride sublimes at a low temperature: aluminium chloride is covalent ( $\ce{Al2Cl6}$ molecules held by weak intermolecular forces), not a giant ionic lattice.
pH about 3: the small, highly charged $\ce{Al^3+}$ ion in $\ce{[Al(H2O)6]^3+}$ polarises the water molecules bonded to it, releasing $\ce{H+}$ : $\ce{[Al(H2O)6]^3+ <=> [Al(H2O)5(OH)]^2+ + H+}$ .

Watch out

Write $\ce{P4O10}$ , not $\ce{P2O5}$ , and $\ce{H3PO4}$ as the product with water.
$\ce{SO2}$ with water gives sulfurous acid, $\ce{H2SO3}$ , not sulfuric acid. Only $\ce{SO3}$ gives $\ce{H2SO4}$ .
Do not say that $\ce{SiCl4}$ "dissolves". It reacts (is hydrolysed). Use "hydrolysis" or "reacts with water".
For $\ce{MgO}$ with water, the pH is about 9, not 13–14. Magnesium hydroxide is only sparingly soluble.
"Steamy fumes" means $\ce{HCl}$ gas meeting moist air; do not call them "white smoke" (that is a solid) or "steam".

Exam tip

pH values: examiners accept a sensible range. Learn 13–14 ( $\ce{Na2O}$ , $\ce{NaOH}$ ), 9 ( $\ce{MgO}$ ), 7 ( $\ce{Al2O3}$ , $\ce{SiO2}$ , $\ce{NaCl}$ ), 6.5 ( $\ce{MgCl2}$ ), 3 ( $\ce{AlCl3}$ ), 1–2 ( $\ce{SiCl4}$ , $\ce{PCl5}$ , $\ce{P4O10}$ ), 0–1 ( $\ce{SO3}$ ).
To show amphoteric behaviour you need two equations, one with an acid and one with a base, and you must use $\ce{NaOH}$ (the syllabus specifies sodium hydroxide).
"Explain in terms of bonding" answers must name the bonding type (ionic or covalent) and link it to electronegativity difference. A pH value alone is not an explanation.
Many Paper 1 questions give a table of melting points, conductivities and pH values and ask which oxide or chloride fits. Use the "suggesting bonding" table above.

Practical skills

The pH of the solution formed by an oxide or chloride is found by adding a small amount of the solid to about $5\ \text{cm}^3$ of distilled water in a test-tube and testing with universal indicator paper or solution (or a pH meter). Covalent chlorides must be handled in a fume cupboard because they release $\ce{HCl}$ . To test whether an oxide or hydroxide is amphoteric, add a little of the solid to two test-tubes, one with dilute $\ce{HCl(aq)}$ and one with $\ce{NaOH(aq)}$ , and warm gently: dissolving in both shows amphoteric behaviour. Record "no visible change" or "white solid remains" where nothing happens; Paper 3 expects every observation recorded.

Summary

Across Period 3, oxides change from basic ( $\ce{Na2O}$ , $\ce{MgO}$ ) to amphoteric ( $\ce{Al2O3}$ ) to acidic ( $\ce{SiO2}$ , $\ce{P4O10}$ , $\ce{SO2}$ , $\ce{SO3}$ ).
With water: $\ce{Na2O}$ pH 13–14, $\ce{MgO}$ pH 9, $\ce{Al2O3}$ and $\ce{SiO2}$ insoluble (pH 7), $\ce{P4O10}$ pH 2, $\ce{SO2}$ pH 2–3 ( $\ce{H2SO3}$ ), $\ce{SO3}$ pH 0–1 ( $\ce{H2SO4}$ ).
$\ce{Al2O3}$ and $\ce{Al(OH)3}$ react with acids and with $\ce{NaOH}$ (giving $\ce{NaAl(OH)4}$ ).
Chlorides: $\ce{NaCl}$ dissolves (pH 7), $\ce{MgCl2}$ dissolves (pH 6.5), aluminium chloride hydrolyses (pH 3), $\ce{SiCl4}$ and $\ce{PCl5}$ hydrolyse completely (pH 1–2) with steamy fumes of $\ce{HCl}$ .
Everything follows from electronegativity: large difference gives ionic compounds (basic oxides, neutral chlorides); small difference gives covalent compounds (acidic oxides, hydrolysed chlorides).

Practice

Question

Write equations for the reactions of magnesium oxide and of phosphorus(V) oxide with water, and state the approximate pH of each solution.
Classify each oxide as basic, amphoteric or acidic: $\ce{Na2O}$ , $\ce{Al2O3}$ , $\ce{SiO2}$ , $\ce{SO2}$ .
Write an equation for the reaction of sulfur dioxide with excess sodium hydroxide.
Explain why aluminium oxide does not change the pH of water even though it is amphoteric.
Write an equation for the reaction of phosphorus(V) chloride with water and describe what is seen.
$2.00\ \text{g}$ of sodium oxide is dissolved in water and made up to $500\ \text{cm}^3$ . Calculate the concentration of sodium hydroxide in the solution. ( $A_r$ : Na 23.0, O 16.0)
Calculate the volume of $2.00\ \text{mol dm}^{-3}$ sodium hydroxide needed to dissolve $5.10\ \text{g}$ of aluminium oxide, forming $\ce{NaAl(OH)4}$ . ( $A_r$ : Al 27.0, O 16.0)
Explain, in terms of bonding, why magnesium chloride gives a nearly neutral solution in water but aluminium chloride gives an acidic one.
$1.00\ \text{g}$ of phosphorus(V) chloride is completely hydrolysed. All the acidic hydrogen atoms in the products are then neutralised by $1.00\ \text{mol dm}^{-3}$ sodium hydroxide (phosphoric acid reacts with three moles of $\ce{NaOH}$ ). Calculate the volume of sodium hydroxide needed. ( $M(\ce{PCl5}) = 208.5\ \text{g mol}^{-1}$ )
A white solid oxide Y has a melting point above 2000 K, does not conduct electricity when molten, is insoluble in water and in dilute acids, but dissolves slowly in hot concentrated sodium hydroxide. The chloride of the same element is a colourless liquid that fumes in moist air. Identify the element, deduce the structure and bonding of the oxide and of the chloride, and write equations for the reaction of the oxide with sodium hydroxide and of the chloride with water.

Answers

$\ce{MgO(s) + H2O(l) -> Mg(OH)2(s)}$ , pH about 9. $\ce{P4O10(s) + 6H2O(l) -> 4H3PO4(aq)}$ , pH about 2 (0–2).
$\ce{Na2O}$ basic; $\ce{Al2O3}$ amphoteric; $\ce{SiO2}$ acidic; $\ce{SO2}$ acidic.
$\ce{SO2(g) + 2NaOH(aq) -> Na2SO3(aq) + H2O(l)}$ .
Aluminium oxide is insoluble in water: its lattice of small, highly charged ions is very strong. Because it does not dissolve or react, it releases neither $\ce{H+}$ nor $\ce{OH-}$ and the pH stays at 7. Its amphoteric nature shows only with acids and alkalis.
$\ce{PCl5(s) + 4H2O(l) -> H3PO4(aq) + 5HCl(aq)}$ . A vigorous reaction, the solid disappears, steamy (misty white) fumes of hydrogen chloride are given off and the solution is strongly acidic (pH 1–2).
$M(\ce{Na2O}) = 62.0$ ; $n = 2.00 / 62.0 = 0.03226\ \text{mol}$ . $n(\ce{NaOH}) = 2 \times 0.03226 = 0.06452\ \text{mol}$ . $c = 0.06452 / 0.500 = 0.129\ \text{mol dm}^{-3}$ .
$M(\ce{Al2O3}) = 102.0$ ; $n = 5.10 / 102.0 = 0.0500\ \text{mol}$ . $\ce{Al2O3 + 2NaOH + 3H2O -> 2NaAl(OH)4}$ , so $n(\ce{NaOH}) = 0.100\ \text{mol}$ . $V = 0.100 / 2.00 = 0.0500\ \text{dm}^3 = 50.0\ \text{cm}^3$ .
Magnesium chloride is ionic: it dissolves to give hydrated $\ce{Mg^2+}$ and $\ce{Cl-}$ ions, and $\ce{Mg^2+}$ polarises water only slightly, so the pH is about 6.5. Aluminium chloride is covalent, and the $\ce{Al^3+}$ ion has a much higher charge density (higher charge, smaller radius). In $\ce{[Al(H2O)6]^3+}$ it draws electron density from the O–H bonds of the bound water molecules, so $\ce{H+}$ is released: $\ce{[Al(H2O)6]^3+ <=> [Al(H2O)5(OH)]^2+ + H+}$ , giving pH about 3.
$n(\ce{PCl5}) = 1.00 / 208.5 = 4.796 \times 10^{-3}\ \text{mol}$ . Each mole gives 5 mol $\ce{HCl}$ (needs 5 mol $\ce{NaOH}$ ) and 1 mol $\ce{H3PO4}$ (needs 3 mol $\ce{NaOH}$ ): 8 mol $\ce{NaOH}$ in total. $n(\ce{NaOH}) = 8 \times 4.796 \times 10^{-3} = 0.03837\ \text{mol}$ . $V = 0.03837 / 1.00 = 0.0384\ \text{dm}^3 = 38.4\ \text{cm}^3$ .
The element is silicon. The oxide, $\ce{SiO2}$ , is giant covalent (giant molecular): very high melting point because many strong Si–O covalent bonds must be broken, no mobile ions or electrons so no conduction when molten, and insoluble. It is a (weakly) acidic oxide, so it reacts with hot concentrated alkali but not with acids: $\ce{SiO2(s) + 2NaOH(aq) -> Na2SiO3(aq) + H2O(l)}$ . The chloride, $\ce{SiCl4}$ , is simple molecular covalent: weak id–id forces between molecules make it a liquid; it is hydrolysed by water vapour, releasing $\ce{HCl}$ fumes: $\ce{SiCl4(l) + 2H2O(l) -> SiO2(s) + 4HCl(aq)}$ .

The organising idea: electronegativity and bonding

Oxides with water

Acid–base behaviour of the oxides and hydroxides

Basic: sodium and magnesium

Amphoteric: aluminium

Acidic: silicon, phosphorus and sulfur

Chlorides with water

Sodium chloride and magnesium chloride

Aluminium chloride

Silicon tetrachloride

Phosphorus pentachloride

Explaining the trends

Suggesting bonding from observations

Worked examples

Practice

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