Group 2 elements

Class 12 Chemistry, unit 5, book pages 79–93

Practise 48 MCQs Take the unit quiz Review 40 cards

Group 2 contains the alkaline earth metals beryllium to barium, all with two outer electrons ( $n s^{2}$ ) that they lose to form $M X^{2 +}$ ions. This unit follows how their size, ionisation energy and reactivity with oxygen, water and acids change down the group, and uses the size and polarising power of the $M X^{2 +}$ ion to explain the thermal stability of carbonates and nitrates and the opposite solubility trends of hydroxides and sulphates. It also covers beryllium's unusual behaviour, complex formation, extraction, and the everyday uses of Group 2 compounds in building, farming and medicine.

Learning outcomes (13)

Describe the properties and trends of Group 2 elements, including their electron configuration, reactivity, and common compounds such as oxides, hydroxides and carbonates.
Explain the chemical reactivity of Group 2 elements, including their reactions with oxygen, water and acids (Be, Mg, Ca).
Explain the reactivity of Group 2 elements in terms of their electron configuration and valence electrons.
Describe the industrial and everyday use of Group 2 compounds, including their role in medicine and agriculture.
Explain and use the term reactivity series and its application in predicting the outcome of chemical reactions.
Explain the extraction and purification process of Group 2 elements and their compounds.
Understand and use the term thermal decomposition and its application in the analysis of Group 2 compounds, especially carbonates and nitrates.
Explain the trend in solubility of Group 2 sulphates and hydroxides using the terms enthalpy of hydration and enthalpy of solution.
Compare the properties and reactivity of Group 2 elements with Group 1 in the periodic table.
Explain the term complex ion and its application in the formation of Group 2 compounds.
Explain the term basic oxide and its application in the formation of Group 2 compounds.
Describe qualitatively the trend in the thermal stability of nitrates and carbonates, including the effect of ionic radius on the polarisation of the large anion.
Describe qualitatively the variation in solubility and enthalpy change of solution, $Δ H_{so l}$ , of hydroxides and sulphates in terms of the relative magnitudes of the enthalpy change of hydration and the lattice energy.

5Group 2: the alkaline earth metals

Group 2 of the Periodic Table contains beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra). Radium is radioactive and rare, so this unit studies Be to Ba. You already know sodium and potassium from Group 1. Group 2 metals are similar but a little less reactive, and they always form ions with a $2 +$ charge.

They are used widely. Magnesium powder burns with a brilliant white light, so it is used in flares and fireworks. Calcium carbonate (limestone) is used for building and to make cement, and barium sulphate is used in hospitals for X-ray pictures of the gut.

Why the name alkaline earth metals

Their oxides and hydroxides dissolve in water (at least partly) to give alkaline solutions, and their oxides were known to the old chemists as earths, because they are found in the Earth's crust and do not change on heating. So they are called the alkaline earth metals.

Electronic configuration

Every Group 2 atom has two electrons in the outermost $s$ subshell. The general valence shell configuration is $n s^{2}$ , where $n$ is the period number.

Electronic configurations of the Group 2 elements
Element	Atomic number	Electronic configuration	Short form
Beryllium (Be)	4	$1 s^{2} 2 s^{2}$	$[He] 2 s^{2}$
Magnesium (Mg)	12	$1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2}$	$[Ne] 3 s^{2}$
Calcium (Ca)	20	$1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2}$	$[Ar] 4 s^{2}$
Strontium (Sr)	38	$\dots 4 s^{2} 3 d^{10} 4 p^{6} 5 s^{2}$	$[Kr] 5 s^{2}$
Barium (Ba)	56	$\dots 5 s^{2} 4 d^{10} 5 p^{6} 6 s^{2}$	$[Xe] 6 s^{2}$

The key idea of the whole unit

A Group 2 atom reacts by losing its two outer electrons to form an $M X^{2 +}$ ion with the stable configuration of a noble gas:

M M X^{2 +} + 2 e X^{-}

Oxidation of a Group 2 metal

Because they lose electrons, the metals are oxidised and act as reducing agents. The easier it is to remove the two electrons, the more reactive the metal. Almost every trend in this unit (reactivity, thermal stability, solubility) comes from the size and charge of the $M X^{2 +}$ ion.

5.1, 5.1.1Trends in ionisation energy and chemical reactivityOften examined

Ionisation energy (IE): The minimum energy required to remove one mole of electrons from one mole of isolated gaseous atoms (or ions). Unit: $kJ mo l^{- 1}$ .
First and second IE: $M (g) M X^{+} (g) + e X^{-}$ is the first IE; $M X^{+} (g) M X^{2 +} (g) + e X^{-}$ is the second IE.

First and second ionisation energies of Group 2

Be

Mg

Ca

Sr

Ba

300

600

900

1200

1500

1800

kJ mo l^{- 1}

second IE

first IE

Both ionisation energies fall steadily from Be to Ba (after the book's Fig 5.1). The second IE is always larger, because the second electron is removed from a positive ion.

Ionisation energies of Group 2 elements (data values read from the book's Fig 5.1)
Element	First IE / $kJ mo l^{- 1}$	Second IE / $kJ mo l^{- 1}$
Be	900	1757
Mg	738	1451
Ca	590	1145
Sr	550	1064
Ba	503	965

Decreasing order of first ionisation energy: $Be > Mg > Ca > Sr > Ba$ .

Why ionisation energy decreases down the group

Reasoning

Atomic size increases. Each element has one more shell of electrons than the one above it, so the outer $s$ electrons are further from the nucleus.
Shielding increases. More inner shells lie between the nucleus and the outer electrons, and they screen the outer electrons from the nuclear charge.
The larger nuclear charge is outweighed. The number of protons rises, but the extra distance and shielding matter more, so the attraction on the outer electrons becomes weaker.
Result: less energy is needed to remove the outer electrons, so both the first and second IE decrease from Be to Ba.

Reactivity increases down the group

Group 2 metals react by losing two electrons. Since this becomes easier down the group, reactivity increases from Be to Ba. The metals also become stronger reducing agents down the group.

With dilute hydrochloric acid: bubbles of hydrogen are given off faster and the reaction is more vigorous down the group. $Mg (s) + 2 HCl (aq) MgCl X_{2} (aq) + H X_{2} (g)$
With oxygen: reactivity increases down the group. Barium is so reactive that it must be stored under oil to keep it away from air.
With water: Be does not react, Mg reacts only with steam, and Ca, Sr and Ba react with cold water faster and faster.

5.1.2Trends in atomic radius

The atoms of Group 2 elements get larger going down the group. Increasing order of atomic radius: $Be < Mg < Ca < Sr < Ba$ .

Metallic radius of Group 2 atoms

Be

Mg

Ca

Sr

Ba

0.05

0.1

0.15

0.2

0.25

r / nm

0.122

0.160

0.197

0.215

0.217

The radius increases down the group: Be < Mg < Ca < Sr < Ba (book Fig 5.2 and Table 5.1; the book's axis label 2.25 should read 0.25).

Reason: going down the group, one new shell of electrons is added for each element. The extra shells and the increased shielding reduce the attraction between the nucleus and the valence electrons, so the outer electrons are held further away and the atom is bigger.

The ions follow the same order. Each $M X^{2 +}$ ion is much smaller than its atom because the whole outer shell has been lost: $Be X^{2 +}$ (about $0.031 nm$ ) $<$ $Mg X^{2 +}$ ( $0.072$ ) $<$ $Ca X^{2 +}$ ( $0.100$ ) $<$ $Sr X^{2 +}$ ( $0.118$ ) $<$ $Ba X^{2 +}$ ( $0.135 nm$ ).

5.1.3Trends in melting point and other physical properties

Group 2 elements are metals. In a metal, positive ions sit in a 'sea' of delocalised electrons, and the metallic bond is the attraction between the ions and this sea. Each Group 2 atom gives two electrons to the sea, so their metallic bonds are stronger than those of Group 1 metals, and their melting points are higher.

Trend: the melting point generally decreases going down the group. The outer electrons are further from the nucleus in larger atoms, so the attraction between the nuclei and the delocalised electrons is weaker, the metallic bond is weaker, and less energy is needed to melt the metal.

Melting points of Group 2 metals

Be

Mg

Ca

Sr

Ba

300

600

900

1200

1500

^{\circ} C

1280

650

838

768

714

The general trend is downward from Ca to Ba. Beryllium is far higher than the rest, and magnesium is lower than the trend because of its different crystal structure (book Fig 5.3).

Table 5.1: Physical properties of the alkaline earth metals (book values)
Element	Metallic radius / nm	Atomic number	Melting point / $^{\circ} C$	Density / $g c m^{- 3}$
beryllium (Be)	0.122	4	1280	1.85
magnesium (Mg)	0.160	12	650	1.74
calcium (Ca)	0.197	20	838	1.55
strontium (Sr)	0.215	38	768	2.66
barium (Ba)	0.217	56	714	3.56

Other physical properties: all are silvery-white, fairly soft metals (Be is harder), good conductors of heat and electricity. Density generally increases down the group (Ca is the lowest) because the mass of the atoms increases faster than their volume.

5.2 (a)Reaction with oxygen and the basic oxidesOften examined

Group 2 metals burn in air, and more rapidly in pure oxygen, to form white solid oxides $MO$ . For example, magnesium ribbon burns with a bright white flame once lit in a Bunsen flame:

2 Mg (s) + O X_{2} (g) 2 MgO (s)

General: 2M + O2 -> 2MO

Beryllium reacts only when it is powdered and strongly heated, because a thin protective layer of $BeO$ covers its surface. Reactivity with oxygen increases down the group: the larger atoms lose their two outer electrons more readily. Barium is so reactive that it is stored under oil.

Strontium and barium also form peroxides ( $MO X_{2}$ , containing the $O X_{2} X^{2 -}$ ion) when heated in excess oxygen, because their large cations can stabilise the large peroxide ion:

Ba (s) + O X_{2} (g) BaO X_{2} (s)

Basic oxides

Basic oxide: A metal oxide that reacts with an acid to give a salt and water (and, if soluble, dissolves in water to give an alkaline solution). Example: $CaO + 2 HCl CaCl X_{2} + H X_{2} O$ .
Amphoteric oxide: An oxide that reacts with both acids and bases. Example: $BeO$ .

All Group 2 oxides are basic except $BeO$ , which is amphoteric. The oxides of Ca, Sr and Ba react with water to form hydroxides, which dissolve to give alkaline solutions:

CaO (s) + H X_{2} O (l) BaO (s) + H X_{2} O (l) Ca (OH) X_{2} (s / aq) Ba (OH) X_{2} (aq)

$BeO$ and $MgO$ are insoluble in water (MgO reacts only very slightly). The solubility of the oxides (as hydroxides) increases down the group, and basic character increases from Be to Ba, because the metal-oxygen bond becomes more ionic as the cation becomes larger.

Beryllium oxide is amphoteric

$BeO$ reacts with an acid (so it behaves as a base) and also with a strong alkali (so it behaves as an acid):

BeO (s) + 2 HCl (aq) BeCl X_{2} (aq) + H X_{2} O (l)

BeO as a base

BeO (s) + 2 NaOH (aq) Na X_{2} BeO X_{2} (aq) + H X_{2} O (l)

BeO as an acid

$Na X_{2} BeO X_{2}$ is sodium beryllate. With more water the product is written as the complex $Na X_{2} [Be (OH) X_{4}]$ : $BeO + 2 NaOH + H X_{2} O Na X_{2} [Be (OH) X_{4}]$ .

Why BeO is covalent and amphoteric: the $Be X^{2 +}$ ion is extremely small and has a $2 +$ charge, so its polarising power is very high. It pulls the electron cloud of the oxide ion towards itself, so the electrons are shared between the two nuclei. The Be-O bond is therefore largely covalent, not purely ionic, and the oxide can react with bases as well as acids.

Flame test

Some Group 2 compounds give a characteristic colour to a flame. The colour comes from the $M X^{2 +}$ ions: heat excites their electrons to higher energy levels, and light of a particular colour is emitted when the electrons fall back. A nichrome wire is cleaned with concentrated hydrochloric acid, dipped into the salt, and held in a non-luminous Bunsen flame.

Flame test colours
Ion in compound	Flame colour
calcium, $Ca X^{2 +}$	brick-red
strontium, $Sr X^{2 +}$	scarlet (crimson) red
barium, $Ba X^{2 +}$	apple-green
beryllium and magnesium	no colour (their electrons are held too tightly to be excited by a Bunsen flame)

5.2 (b)Reaction with waterOften examined

The reactivity of Group 2 metals with water increases down the group: hydrogen gas is released more and more rapidly.

Reactions of Group 2 metals with water
Metal	Reaction with water	Equation
Be	no reaction, even with steam; a protective $BeO$ layer covers the metal	none
Mg	very slow with cold water (a few bubbles over several days); burns in steam with a bright white light	$Mg (s) + H X_{2} O (g) MgO (s) + H X_{2} (g)$
Ca	reacts steadily with cold water; bubbles of hydrogen, cloudy white suspension	$Ca (s) + 2 H X_{2} O (l) Ca (OH) X_{2} (aq) + H X_{2} (g)$
Sr	reacts more rapidly with cold water	$Sr (s) + 2 H X_{2} O (l) Sr (OH) X_{2} (aq) + H X_{2} (g)$
Ba	reacts vigorously with cold water	$Ba (s) + 2 H X_{2} O (l) Ba (OH) X_{2} (aq) + H X_{2} (g)$

Magnesium and cold water give magnesium hydroxide, but the reaction is extremely slow because the $Mg (OH) X_{2}$ is insoluble and coats the metal: $Mg (s) + 2 H X_{2} O (l) Mg (OH) X_{2} (s) + H X_{2} (g)$ . With steam the product is the oxide, not the hydroxide, because the hydroxide decomposes at that temperature.

Calcium gives a cloudy white suspension because calcium hydroxide is only slightly soluble. The part that dissolves makes the solution alkaline (pH about 12).

M (s) + 2 H X_{2} O (l) M (OH) X_{2} (aq) + H X_{2} (g) (M = Ca, Sr, Ba)

General equation

5.2 (c), (d)Reaction of the metals and their oxides with dilute acidsOften examined

Metals with dilute acids

Group 2 metals undergo a redox reaction with dilute acids to form a salt and hydrogen gas. The metal is oxidised to $M X^{2 +}$ and the $H X^{+}$ ions are reduced to $H X_{2}$ . The reactions become more vigorous and more exothermic down the group.

M (s) + 2 HCl (aq) MCl X_{2} (aq) + H X_{2} (g)

General equation

Mg (s) + 2 HCl (aq) Ca (s) + 2 HCl (aq) MgCl X_{2} (aq) + H X_{2} (g) CaCl X_{2} (aq) + H X_{2} (g)

Mg (s) + 2 H X^{+} (aq) Mg X^{2 +} (aq) + H X_{2} (g)

Ionic equation

Beryllium also dissolves in dilute acids, but only slowly, because of its oxide layer.

Group 2 oxides with acids

Group 2 oxides are basic, so they neutralise acids to form a salt and water. These reactions are used to make useful compounds:

MgO (s) + H X_{2} SO X_{4} (aq) MgSO X_{4} (aq) + H X_{2} O (l)

CaO (s) + 2 HCl (aq) CaCl X_{2} (aq) + H X_{2} O (l)

Magnesium sulphate (Epsom salt) is used for short-term relief of constipation, and as a soaking solution to relieve minor sprains, bruises, muscle aches and joint stiffness.
Calcium chloride is used as a solidifying agent in paint production and as a coagulant in the manufacture of rubber. It is also used as a drying agent in the laboratory.

Concept assessment 5.1, questions 1 to 3

(1) What is the general trend in the melting points going down Group 2? (2) Explain why the atoms in Group 2, as in any other group, get larger with increasing atomic number. (3) Write a balanced chemical equation, including state symbols, for the reaction of (i) strontium with oxygen, (ii) strontium oxide with hydrochloric acid.

(1) The melting point generally decreases down the group (Be $128 0^{\circ} C$ , Ba $71 4^{\circ} C$ ), because the larger atoms hold the delocalised electrons less strongly, so the metallic bond is weaker. Magnesium is an exception and is lower than calcium.

(2) Each element down the group has one more occupied shell of electrons. The outer electrons are further from the nucleus and more shielded by inner shells, so they are held less tightly and the atom is larger.

(3)(i) 2 Sr (s) + O X_{2} (g) 2 SrO (s)

(3)(ii) SrO (s) + 2 HCl (aq) SrCl X_{2} (aq) + H X_{2} O (l)

Concept assessment 5.1, question 4

Barium reacts vigorously when added to water. (i) Write a balanced chemical equation, including state symbols, for the reaction of barium with water. (ii) State two observations that could be made during this reaction. (iii) Suggest the approximate pH of the resulting solution. (iv) Will the reaction be more or less vigorous than the reaction of barium with water? Explain your answer.

(i) Ba (s) + 2 H X_{2} O (l) Ba (OH) X_{2} (aq) + H X_{2} (g)

(ii) Rapid fizzing (bubbles of a colourless gas, hydrogen); the metal dissolves and disappears; the water gets warm; a colourless solution forms.

(iii) Barium hydroxide is soluble and fully ionised, so the solution is strongly alkaline, pH about 13 to 14.

(iv) Part (iv) is misprinted in the book: it compares barium with itself. It is meant to compare another Group 2 metal (such as strontium or calcium) with barium. Strontium (or calcium) reacts less vigorously than barium, because it is higher in the group: its atoms are smaller, its ionisation energies are larger, and it loses its two outer electrons less easily.

Concept assessment 5.1, question 5

Describe what you would see when magnesium reacts with cold water and steam. Also write an equation for the reaction with steam.

Cold water: the reaction is very slow. Only a few tiny bubbles of hydrogen form on the surface of the magnesium, over several days. The insoluble $Mg (OH) X_{2}$ layer slows the reaction further.

Steam: heated magnesium burns in steam with a bright white light, leaving a white powder of magnesium oxide. Hydrogen is given off and can be burned at the end of the apparatus.

Mg (s) + H X_{2} O (g) MgO (s) + H X_{2} (g)

5.3Comparison of Group 1 and Group 2 elementsOften examined

The reactivity of a metal depends mainly on its ionisation energy: the smaller the ionisation energy, the greater the reactivity. Group 2 metals have higher ionisation energies than the Group 1 metals of the same period, so they are less reactive.

Why Group 2 ionisation energies are higher

Greater nuclear charge. Mg (12 protons) has one more proton than Na (11 protons), but its extra electron goes into the same shell, where it shields poorly.
Smaller atoms. The stronger pull of the nucleus draws the electrons closer, so a Group 2 atom is smaller than the Group 1 atom of the same period (Mg 0.160\,\mathrm{nm}, Na 0.186\,\mathrm{nm}).
**Paired ns^2 electrons.** A Group 2 atom must lose two electrons from a stable, filled s subshell; a Group 1 atom loses one electron from a half-filled ns^1 subshell.
Result: first IE of Mg is 738\,\mathrm{kJ\,mol^{-1}} but that of Na is only 496\,\mathrm{kJ\,mol^{-1}} (Be 900 vs Li 520). Group 1 metals lose electrons more easily, so they are stronger reducing agents and more reactive.

Reaction with water shows the difference. Alkali metals react vigorously with cold water; with sodium or potassium the reaction is so fast that the hydrogen may catch fire. Alkaline earth metals react slowly (Ca) or hardly at all (Mg, Be).

\ce{2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)}\qquad\text{fast}

\ce{Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)}\qquad\text{slow, steady}

Group 1 and Group 2 compared
Property	Group 1 (alkali metals)	Group 2 (alkaline earth metals)
Valence configuration	ns^1	ns^2
Ion formed, oxidation state	\ce{M+}, +1	\ce{M^2+}, +2
Atomic size (same period)	larger	smaller
First ionisation energy	lower	higher
Reactivity and reducing power	higher	lower
Metallic bond, melting point, hardness	weaker (one electron per atom), low m.p., soft	stronger (two electrons per atom), higher m.p., harder
Reaction with cold water	vigorous (all)	Be none, Mg very slow, Ca to Ba steady to fast
Hydroxides	strong alkalis, all very soluble	weaker bases, less soluble
Carbonates on heating	stable (except \ce{Li2CO3})	decompose to \ce{MO + CO2}
Nitrates on heating	give nitrite and \ce{O2} (except \ce{LiNO3})	give oxide, \ce{NO2} and \ce{O2}
Tendency to form complexes	very small	greater (especially Be and Mg)

SLOThe reactivity series and predicting reactions

Reactivity series: A list of metals arranged in order of decreasing reactivity, that is, in order of decreasing tendency to lose electrons and form positive ions.

Part of the series, with Group 2 metals highlighted in bold (carbon and hydrogen are included for comparison):

\ce{K > Na > \mathbf{Ca} > \mathbf{Mg} > Al > (C) > Zn > Fe > Pb > (H) > Cu > Ag > Au}

Down Group 2 reactivity increases, so in the full series barium and strontium lie above calcium, and beryllium lies below magnesium.

Using the series to predict the outcome of reactions

Displacement: a metal displaces any metal below it from a solution of its salt. \ce{Mg(s) + CuSO4(aq) -> MgSO4(aq) + Cu(s)} happens (blue colour fades, brown copper forms), but \ce{Cu + MgSO4} gives no reaction.
Reaction with acids: metals above hydrogen give \ce{H2} with dilute acids (Mg, Ca, Zn, Fe); metals below hydrogen (Cu, Ag) do not.
Reaction with water: only metals near the top (K, Na, Ca, Sr, Ba) react with cold water; Mg, Zn and Fe need steam.
Reducing oxides: a metal can remove oxygen from the oxide of a metal below it: \ce{2Mg + CO2 -> 2MgO + C}, and magnesium reduces \ce{BeF2} to Be during extraction.
Extraction method: metals above carbon (Group 1, Group 2, Al) cannot be made by heating the oxide with carbon; they are extracted by electrolysis of the molten chloride (or by reduction with a more reactive metal).

5.4Uses of Group 2 compoundsOften examined

The limestone cycle: the three calcium compounds and their main uses.

Limestone in industry

Limestone is mainly calcium carbonate, \ce{CaCO3}. It is cut into building blocks, which are joined with mortar (once made from lime and sand, now usually cement and sand). Marble, another form of \ce{CaCO3}, is used for expensive tiles and buildings.

Most calcium carbonate is used to make cement. The first stage is roasting limestone in a lime kiln. At high temperature calcium carbonate decomposes to calcium oxide, called lime or quicklime:

\ce{CaCO3(s) ->[heat] CaO(s) + CO2(g)}

The calcium oxide is then roasted with clay to make cement. Cement mixed with sand and small pieces of rock makes concrete, the most widely used building material in the world. Its tensile strength is improved by letting it set around iron rods (reinforced concrete).

Slaked lime in agriculture

Calcium hydroxide, \ce{Ca(OH)2} (slaked lime), is made by adding water to quicklime: \ce{CaO + H2O -> Ca(OH)2}. Farmers spread it on acidic soil. Being a base, it neutralises the acid and raises the pH of the soil so that crops grow better:

\ce{Ca(OH)2(s) + 2H+(aq) -> Ca^2+(aq) + 2H2O(l)}

Powdered limestone (\ce{CaCO3}) is also used for the same purpose; it acts more slowly but lasts longer.

Group 2 compounds in medicine

Barium sulphate (\ce{BaSO4}, 'barium meal') is given to patients before X-ray pictures of the stomach and intestines. Barium absorbs X-rays strongly, so the gut shows up clearly. Barium ions are poisonous, but \ce{BaSO4} is so insoluble that it is not absorbed into the blood from the gut.
Antacids: Group 2 hydroxides and carbonates are bases, so they neutralise excess hydrochloric acid in the stomach (acid indigestion). Tablets such as Gaviscon and Rennie contain \ce{CaCO3} and \ce{MgCO3}; Milk of Magnesia is a suspension of \ce{Mg(OH)2} in water.
Magnesium sulphate (Epsom salt) is a laxative and is used in soaking solutions for sprains and bruises.
Plaster of Paris, \ce{CaSO4.1/2H2O}, sets hard with water and is used to support broken bones.

\ce{CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + CO2(g) + H2O(l)}

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

Other uses

Magnesium powder in flares and fireworks; strontium and barium salts give red and green colours in fireworks.
Magnesium is alloyed with aluminium to make light, strong alloys for aircraft and car parts.
Beryllium is used in X-ray tube windows (it lets X-rays through) and in light, hard alloys with copper.

5.4, 5.4.1Thermal decomposition of Group 2 carbonates and nitratesOften examined

Thermal decomposition: The breakdown of a compound into two or more different substances by the application of heat.
Thermal stability: The resistance of a compound to decomposition on heating. A more stable compound needs a higher temperature to decompose.

Carbonates

Group 2 carbonates decompose on heating to form the metal oxide and carbon dioxide:

\ce{MCO3(s) ->[heat] MO(s) + CO2(g)}

General equation

\ce{MgCO3(s) ->[heat] MgO(s) + CO2(g)}\qquad \ce{CaCO3(s) ->[heat] CaO(s) + CO2(g)}

The temperature needed for decomposition increases going down the group.

Table 5.2: Decomposition of Group 2 carbonates
Group 2 carbonate	Decomposition temperature / ^\circ\mathrm{C}	Enthalpy change of reaction / \mathrm{kJ\,mol^{-1}}
Magnesium carbonate	540	+117
Calcium carbonate	900	+176
Strontium carbonate	1280	+238
Barium carbonate	1360	+268

Decomposition temperatures of Group 2 carbonates

A carbonate needs a higher temperature to decompose the further down the group its metal lies (data of book Table 5.2). Labels show the metal of each carbonate \ce{MCO3}.

The enthalpy change of decomposition becomes more positive down the group: more energy is needed to break the carbonate down, so the carbonate is more stable. (\ce{BeCO3} is so unstable that it decomposes even at room temperature unless kept under \ce{CO2}.)

Nitrates

Group 2 nitrates decompose on heating to form the metal oxide, nitrogen dioxide and oxygen. A brown gas, toxic nitrogen dioxide (\ce{NO2}), is seen, and the oxygen relights a glowing splint.

\ce{2M(NO3)2(s) ->[heat] 2MO(s) + 4NO2(g) + O2(g)}

General equation

\ce{2Ca(NO3)2(s) ->[heat] 2CaO(s) + 4NO2(g) + O2(g)}

Like the carbonates, the nitrates need a higher temperature to decompose as the group is descended.

Explaining the trend: polarisation of the anion

The small \ce{Mg^2+} ion pulls the electron cloud of the large \ce{CO3^2-} ion towards itself (polarisation). The larger \ce{Ca^2+} ion distorts it much less (after the book's Fig 5.4).

Derivation Why thermal stability increases down Group 2

Aim. Explain, using ion polarisation, why \ce{MgCO3} decomposes at 540^\circ\mathrm{C} but \ce{BaCO3} needs 1360^\circ\mathrm{C}.

The anion is large. The carbonate ion \ce{CO3^2-} (and the nitrate ion \ce{NO3-}) has a large ionic radius, so its electron cloud is easily distorted (polarised).
The cation's polarising power depends on its charge density. All Group 2 cations have the same charge, 2+, but their size increases down the group.
\ce{Mg^2+ < Ca^2+ < Sr^2+ < Ba^2+}\quad\text{(ionic radius)}
Small cations polarise more. The small \ce{Mg^2+} ion has the highest charge density. It attracts the electron cloud of the carbonate ion towards itself and distorts it, giving the bond considerable covalent character.
\text{polarising power: }\ce{Mg^2+ > Ca^2+ > Sr^2+ > Ba^2+}
Polarisation weakens a C-O bond. Electron density is pulled away from one of the carbon-oxygen bonds inside the carbonate ion, weakening it. Less heat is then needed to break it, releasing \ce{CO2} and leaving the oxide ion with the metal.
Down the group the effect weakens. \ce{Ba^2+} is large, so it hardly distorts the carbonate ion; the C-O bonds stay strong and a much higher temperature is needed.
The same argument applies to nitrates. A small cation polarises the \ce{NO3-} ion, weakens an N-O bond and helps release \ce{NO2} and \ce{O2}.

Result\text{Stability: }\ce{BaCO3 > SrCO3 > CaCO3 > MgCO3}\qquad \ce{Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2}

The smaller and more polarising the cation, the less stable its carbonate or nitrate. Thermal stability increases down the group, so \ce{MgCO3} and \ce{Mg(NO3)2} decompose most easily.

Concept assessment 5.2

(1) Which one of the three compounds listed will decompose at the lowest temperature? (i) Calcium carbonate, strontium carbonate, barium carbonate; (ii) barium nitrate, calcium nitrate, magnesium nitrate. (2) Write a balanced chemical equation, including state symbols, for the thermal decomposition of (i) strontium carbonate, (ii) barium nitrate.

(1)(i) Calcium carbonate. \ce{Ca^2+} is the smallest cation of the three, so it polarises the carbonate ion most.

(1)(ii) Magnesium nitrate. \ce{Mg^2+} is the smallest cation, with the greatest polarising power.

\text{(2)(i)}\quad\ce{SrCO3(s) ->[heat] SrO(s) + CO2(g)}

\text{(2)(ii)}\quad\ce{2Ba(NO3)2(s) ->[heat] 2BaO(s) + 4NO2(g) + O2(g)}

5.5 (a) to (d)Solubility of Group 2 hydroxidesOften examined

The solubility of the hydroxides increases down the group; barium hydroxide is quite soluble in water.

\ce{Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2}

Order of solubility

Table 5.3: Solubilities of Group 2 hydroxides
Group 2 hydroxide	Solubility at 298\,\mathrm{K} / mol per 100\,\mathrm{g} of water
\ce{Mg(OH)2}	2\times10^{-5}
\ce{Ca(OH)2}	1.5\times10^{-3}
\ce{Sr(OH)2}	3.4\times10^{-3}
\ce{Ba(OH)2}	1.5\times10^{-2}

So the solutions become more alkaline down the group. \ce{Mg(OH)2} is so slightly soluble that it is used as a mild antacid; a saturated solution of \ce{Ca(OH)2} is lime water (pH about 12); \ce{Ba(OH)2} gives a strongly alkaline solution (pH about 13 to 14).

Enthalpy of solution from lattice and hydration enthalpies

Lattice energy (lattice enthalpy): The energy released when one mole of an ionic solid is formed from its gaseous ions. The reverse, breaking the lattice into gaseous ions, is the lattice dissociation enthalpy (endothermic).
Enthalpy change of hydration, \Delta H_{hyd}: The enthalpy change when one mole of gaseous ions dissolves in water to form a very dilute solution: \ce{M^2+(g) -> M^2+(aq)}. Always exothermic.
Enthalpy change of solution, \Delta H_{sol}: The enthalpy change when one mole of a solute dissolves in water to form a very dilute solution.

Dissolving is imagined in two steps: break the lattice into gaseous ions (costs energy), then hydrate the ions (releases energy). Here hydration releases more than the lattice costs, so \Delta H_{sol} is negative.

\Delta H_{sol}=\Delta H_{latt\,diss}+\Delta H_{hyd}(\ce{M^2+})+2\,\Delta H_{hyd}(\ce{OH-})

Hess's law

The lattice dissociation term is positive (energy in); the hydration terms are negative (energy out).

As a rough guide, the more negative (exothermic) \Delta H_{sol} is, the more soluble the salt. A compound is likely to be soluble only if \Delta H_{sol} is negative or has a small positive value.

(b) Change in hydration enthalpy down the group

Hydration enthalpy is directly proportional to the charge of the ion and inversely proportional to its size. Smaller ions (with the same charge) attract water molecules more strongly, so they have more exothermic hydration enthalpies. Down the group it becomes less exothermic:

\Delta H_{hyd}:\ \ce{Mg^2+} (-1920) > \ce{Ca^2+} (-1650) > \ce{Sr^2+} (-1480) > \ce{Ba^2+} (-1360)\ \mathrm{kJ\,mol^{-1}}

Order of magnitude of the (exothermic) hydration enthalpy; approximate data values.

For the hydroxides this decrease is relatively small. It depends only on the cation, because the anion (the hydroxide ion) is the same in every case.

(c) Change in lattice energy down the group

Lattice energy is larger when the ions are small, and it is inversely proportional to the sum of the radii of cation and anion, r_++r_-. The hydroxide ion is small, so r_++r_- changes a lot as the cation grows. The lattice energy of the hydroxides therefore decreases by a relatively large amount down the group, in the order \ce{Mg(OH)2 > Ca(OH)2 > Sr(OH)2 > Ba(OH)2}.

(d) Conclusion: why hydroxides become more soluble

Reasoning

Down the group, lattice energy falls by a large amount (less energy needed to break the lattice).
Hydration enthalpy falls by a smaller amount (only a little less energy released on hydration).
So the energy cost falls faster than the energy gain, and \Delta H_{sol} becomes more exothermic (more negative) down the group.
\Delta H_{sol}=\underbrace{\Delta H_{latt\,diss}}_{\text{falls a lot}}+\underbrace{\Delta H_{hyd}}_{\text{falls a little}}\ \Rightarrow\ \text{more negative}
Therefore the solubility of Group 2 hydroxides increases from \ce{Mg(OH)2} to \ce{Ba(OH)2}.

Opposite solubility trends

\log_{10} of the solubilities in book Tables 5.3 (hydroxides, \mathrm{mol} per 100\,\mathrm{g} water) and 5.4 (sulphates, \mathrm{mol\,dm^{-3}}). Only the direction of each line matters: hydroxides become more soluble down the group, sulphates less soluble.

5.5 (e)Solubility of Group 2 sulphatesOften examined

The solubility of the sulphates decreases down the group; barium sulphate is insoluble in water. This is the opposite of the hydroxide trend.

\ce{MgSO4 > CaSO4 > SrSO4 > BaSO4}\qquad(\ce{BeSO4}\text{ is the most soluble of all})

Order of solubility

Table 5.4: Solubilities of Group 2 sulphates
Group 2 sulphate	Solubility at 298\,\mathrm{K} / \mathrm{mol\,dm^{-3}}
\ce{MgSO4}	1.83
\ce{CaSO4}	4.66\times10^{-2}
\ce{SrSO4}	7.11\times10^{-4}
\ce{BaSO4}	9.43\times10^{-6}

Explanation

Why sulphates become less soluble

Lattice energy changes very little. The sulphate ion is very large. Lattice energy depends on 1/(r_++r_-), and when r_- is large, a bigger cation hardly changes the sum. So the lattice dissociation enthalpy decreases only slightly down the group.
Hydration enthalpy falls considerably. The hydration enthalpy of the cation becomes much less exothermic as the cation gets larger (-1920 for \ce{Mg^2+} to -1360\,\mathrm{kJ\,mol^{-1}} for \ce{Ba^2+}).
Balance: the hydration enthalpy decreases more than the lattice dissociation enthalpy. The energy released on hydration is no longer enough to pay for breaking the lattice, so \Delta H_{sol} becomes more endothermic (less negative) down the group.
Result: solubility decreases from \ce{MgSO4} (soluble) to \ce{BaSO4} (insoluble).

The two opposite solubility trends side by side
	Hydroxides	Sulphates
Size of anion	small (\ce{OH-})	large (\ce{SO4^2-})
Fall in lattice energy down group	large	small
Fall in hydration enthalpy down group	smaller than lattice	larger than lattice
\Delta H_{sol} down group	more exothermic	more endothermic
Solubility down group	increases	decreases

Concept assessment 5.3

Explain why magnesium sulphate is more soluble than barium sulphate by referring to the relative values of the lattice energies and enthalpy changes of hydration.

The sulphate ion is large, so the lattice energies of \ce{MgSO4} and \ce{BaSO4} are similar: changing the cation from \ce{Mg^2+} to \ce{Ba^2+} hardly changes r_++r_-. However, the small \ce{Mg^2+} ion has a much more exothermic hydration enthalpy (about -1920\,\mathrm{kJ\,mol^{-1}}) than the large \ce{Ba^2+} ion (about -1360\,\mathrm{kJ\,mol^{-1}}).

For \ce{MgSO4} the hydration energy released is large enough to overcome the lattice energy, so \Delta H_{sol} is exothermic and it dissolves. For \ce{BaSO4} the hydration energy is much smaller than the lattice energy, \Delta H_{sol} is endothermic, and it is insoluble.

5.6Complexes of the alkaline earth metals

Complex ion: A central metal ion joined by coordinate (dative) bonds to a number of surrounding molecules or ions called ligands. Example: \ce{[Be(H2O)4]^2+}.
Ligand: A molecule or ion with a lone pair of electrons that it donates to the central metal ion (a Lewis base), such as \ce{H2O}, \ce{OH-}, \ce{F-}.
Coordination number: The number of coordinate bonds formed by the central ion: 4 in \ce{[Be(H2O)4]^2+}, 6 in \ce{[Mg(H2O)6]^2+}.

Group 2 elements almost always form ionic compounds containing the \ce{M^2+} ion. Because their ions have a higher charge (2+) and smaller radius than Group 1 ions, they have a much greater tendency to form complexes with Lewis bases than the alkali metals do. This tendency is greatest for the smallest cation, \ce{Be^2+}, and decreases rapidly as the ionic radius increases.

Beryllium complexes

The chemistry of \ce{Be^2+} is dominated by its behaviour as a Lewis acid (electron-pair acceptor). \ce{Be^2+} has lost both 2s electrons, so it has four empty orbitals (one 2s and three 2p). These hybridise (sp^3) and accept four lone pairs, giving an octet around beryllium and a tetrahedral complex.

\ce{Be^2+} has four empty orbitals (one 2s and three 2p). Four water molecules each donate a lone pair into them, giving the tetrahedral \ce{[Be(H2O)4]^2+} ion (after the book's figure on p. 89).

\ce{Be^2+} salts dissolve in water to form acidic solutions containing the tetrahedral \ce{[Be(H2O)4]^2+} ion. Because of its high charge-to-radius ratio, \ce{Be^2+} polarises the coordinated water molecules, weakening their O-H bonds, so they lose \ce{H+} easily:

\ce{[Be(H2O)4]^2+ -> [Be(H2O)3(OH)]+ + H+}

With a strong base, beryllium and its salts form the tetrahedral hydroxo complex \ce{[Be(OH)4]^2-}. This is why beryllium oxide and hydroxide are amphoteric:

\ce{Be(OH)2(s) + 2OH-(aq) -> [Be(OH)4]^2-(aq)}

Beryllium also forms very stable fluoride complexes such as \ce{[BeF4]^2-} and \ce{[BeF3]-}, and beryllium halides act as Lewis acids by forming adducts with Lewis bases (for example \ce{BeCl2 + 2Cl- -> [BeCl4]^2-}). All this is due to the small size and high charge density of \ce{Be^2+}.

Magnesium and calcium complexes

The heavier alkaline earth metals form few complexes, usually with a coordination number of 6 or more. Complex formation is most important for the smaller cations \ce{Mg^2+} and \ce{Ca^2+}. Aqueous solutions of \ce{Mg^2+} contain the octahedral \ce{[Mg(H2O)6]^2+} ion: magnesium can use one 3s, three 3p and two 3d orbitals of its outer shell (sp^3d^2) to accept six lone pairs. Chlorophyll, the green pigment of plants, is a magnesium complex.

Concept assessment 5.4, question 1

Why does beryllium form complex compounds?

\ce{Be^2+} is a very small ion with a 2+ charge, so it has a very high charge density and strongly attracts lone pairs of electrons. It also has four empty orbitals (2s and 2p) that can accept lone pairs. It therefore acts as a Lewis acid and forms coordinate bonds with Lewis bases, giving tetrahedral complexes such as \ce{[Be(H2O)4]^2+}, \ce{[Be(OH)4]^2-} and \ce{[BeF4]^2-}.

5.2, 5.6Unusual behaviour of berylliumOften examined

Beryllium, the first member of Group 2, differs from the other members in many ways. The reasons are the same in every case: its atom and \ce{Be^2+} ion are very small, the ion has a very high charge density and polarising power, its ionisation energy is high, and it has **no d orbitals** in its valence shell, so its maximum coordination number is 4.

How beryllium differs from the rest of Group 2
Property	Beryllium	Other Group 2 metals
Bonding in compounds	largely covalent (\ce{BeCl2} is covalent, low m.p.)	ionic
Oxide and hydroxide	amphoteric (\ce{BeO}, \ce{Be(OH)2})	basic
Reaction with water	none, even with steam	Mg with steam; Ca, Sr, Ba with cold water
Melting point, hardness	highest in group (1280^\circ\mathrm{C}), hard	lower, softer
Maximum coordination number	4 (no d orbitals)	6 (Mg can use 3d)
Complex formation	strong tendency	weak
Carbonate	very unstable	decomposes only on strong heating
Peroxide	not formed	Sr and Ba form peroxides
Flame colour	none	Ca, Sr, Ba give colours

Amphoteric beryllium oxide

\ce{BeO(s) + 2HCl(aq) -> BeCl2(aq) + H2O(l)}

Reacts with an acid

\ce{BeO(s) + 2NaOH(aq) -> Na2BeO2(aq) + H2O(l)}

Reacts with a base

5.7Extraction and purification of Group 2 elements

Group 2 metals are very reactive, so they are never found free in nature. They occur in minerals such as carbonates, sulphates, oxides and silicates. Because they lie above carbon in the reactivity series, they cannot be made by heating their oxides with carbon; electrolysis of the molten chloride, or reduction with a more reactive metal, is used.

1. Extraction

a. Ore preparation: the ore is mined and crushed to a fine powder.
b. Roasting: carbonates are unstable to heat and decompose to the oxide. For example, to extract magnesium from its carbonate it is first roasted: \ce{MgCO3(s) ->[heat] MgO(s) + CO2(g)}.
c. Reduction: the metal compound is reduced to the metal, usually by electrolysis of the molten compound. Calcium and strontium can also be obtained by reducing their halides with a more reactive metal such as sodium or magnesium.

2. Electrolysis and purification

Pure samples of most alkaline earth metals are obtained by electrolysis of their molten chlorides (or oxides). Electrolysis can also be used to purify the metal further. Beryllium was first obtained by reducing its chloride; radium chloride, which is radioactive, was obtained through a series of reactions and separations.

Examples

Extraction of beryllium and magnesium
Metal	Extraction	Purification
Beryllium	from beryl ore, \ce{Be3Al2(SiO3)6}, by acid digestion and solvent extraction, giving \ce{Be(OH)2}	\ce{Be(OH)2} is converted to \ce{BeF2}, which is reduced by magnesium: \ce{BeF2 + Mg ->[heat] Be + MgF2}
Magnesium	electrolysis of molten \ce{MgCl2} obtained from sea water or brines	fractional distillation or vacuum distillation

\text{cathode: }\ce{Mg^2+ + 2e- -> Mg(l)}\qquad\text{anode: }\ce{2Cl- -> Cl2(g) + 2e-}

Electrolysis of molten MgCl2

Concept assessment 5.4, question 2

Starting from calcium carbonate, how would you extract pure calcium metal?

Convert the carbonate to the chloride with hydrochloric acid, then evaporate the solution to obtain solid \ce{CaCl2}.
\ce{CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + CO2(g) + H2O(l)}
Melt the dry \ce{CaCl2} (a little \ce{CaF2} is added to lower its melting point) and electrolyse it. Calcium is formed at the cathode and chlorine at the anode.
\text{cathode: }\ce{Ca^2+ + 2e- -> Ca}\qquad\text{anode: }\ce{2Cl- -> Cl2 + 2e-}
Purify the calcium by distillation under vacuum.

(Alternatively, the carbonate can be roasted to \ce{CaO}, which is reduced with aluminium at high temperature under vacuum.)

Key points

Group 2 metals have the valence configuration ns^2 and form \ce{M^2+} ions; they are good conductors with fairly high melting points.
Down the group: atomic radius increases, ionisation energy decreases, reactivity with oxygen, water and acids increases.
The metals burn to white basic oxides; \ce{BeO} is amphoteric. Hydroxides become more soluble and more alkaline down the group.
Sulphates become less soluble down the group (\ce{BaSO4} is insoluble).
Carbonates and nitrates become more resistant to thermal decomposition down the group, because the cations become larger and less polarising.
Group 2 metals are less reactive than Group 1 metals because of their higher ionisation energies.
Limestone is used for building and cement; slaked lime neutralises acidic soil; \ce{BaSO4}, \ce{Mg(OH)2} and \ce{CaCO3} are used in medicine.
Pure metals are obtained by electrolysis of their molten chlorides or oxides.

Formula sheet

Formation of the Group 2 ion

\ce{M -> M^2+ + 2e-}

\ce{M}: Be, Mg, Ca, Sr or Ba; the metal is oxidised and acts as a reducing agent

Unit: none

Burning in oxygen

\ce{2M(s) + O2(g) -> 2MO(s)}

white basic oxide (\ce{BeO} amphoteric); Sr and Ba also form peroxides \ce{MO2}

Unit: none

Reaction with water (Ca, Sr, Ba)

\ce{M(s) + 2H2O(l) -> M(OH)2(aq) + H2(g)}

Mg reacts with steam instead: \ce{Mg + H2O(g) -> MgO + H2}; Be does not react

Unit: none

Reaction with dilute hydrochloric acid

\ce{M(s) + 2HCl(aq) -> MCl2(aq) + H2(g)}

ionic form \ce{M + 2H+ -> M^2+ + H2}

Unit: none

Basic oxide with acid

\ce{MO(s) + 2HCl(aq) -> MCl2(aq) + H2O(l)}

e.g. \ce{MgO + H2SO4 -> MgSO4 + H2O}

Unit: none

Amphoteric beryllium oxide

\ce{BeO + 2HCl -> BeCl2 + H2O}\qquad\ce{BeO + 2NaOH -> Na2BeO2 + H2O}

\ce{Na2BeO2}: sodium beryllate

Unit: none

Thermal decomposition of carbonates

\ce{MCO3(s) ->[heat] MO(s) + CO2(g)}

decomposition temperature rises down the group

Unit: none

Thermal decomposition of nitrates

\ce{2M(NO3)2(s) ->[heat] 2MO(s) + 4NO2(g) + O2(g)}

brown \ce{NO2} gas is seen

Unit: none

Enthalpy change of solution

\Delta H_{sol}=\Delta H_{latt\,diss}+\Delta H_{hyd}

\Delta H_{latt\,diss}: lattice dissociation enthalpy (positive); \Delta H_{hyd}: sum of hydration enthalpies of the ions (negative)

Unit: \mathrm{kJ\,mol^{-1}}

Lime kiln

\ce{CaCO3(s) ->[heat] CaO(s) + CO2(g)}

quicklime; with water gives slaked lime \ce{CaO + H2O -> Ca(OH)2}

Unit: none

Liming acidic soil

\ce{Ca(OH)2(s) + 2H+(aq) -> Ca^2+(aq) + 2H2O(l)}

raises the pH of the soil

Unit: none

Antacids

\begin{aligned}\ce{Mg(OH)2 + 2HCl &-> MgCl2 + 2H2O}\\\ce{CaCO3 + 2HCl &-> CaCl2 + CO2 + H2O}\end{aligned}

neutralise excess stomach acid

Unit: none

Hydrolysis of the hydrated beryllium ion

\ce{[Be(H2O)4]^2+ -> [Be(H2O)3(OH)]+ + H+}

makes \ce{Be^2+} solutions acidic

Unit: none

Electrolysis of molten magnesium chloride

\ce{MgCl2(l) -> Mg(l) + Cl2(g)}

cathode \ce{Mg^2+ + 2e- -> Mg}; anode \ce{2Cl- -> Cl2 + 2e-}

Unit: none

Key terms

Alkaline earth metals: The Group 2 elements Be, Mg, Ca, Sr, Ba and Ra, whose oxides ('earths') give alkaline solutions.
Ionisation energy: The minimum energy needed to remove one mole of electrons from one mole of gaseous atoms or ions; unit \mathrm{kJ\,mol^{-1}}.
Shielding effect: The reduction of the attraction of the nucleus on outer electrons caused by the inner electron shells.
Reducing agent: A substance that loses electrons and is itself oxidised; Group 2 metals are reducing agents.
Metallic bond: The attraction between positive metal ions and the sea of delocalised electrons around them.
Basic oxide: A metal oxide that reacts with acids to form a salt and water, such as \ce{MgO} and \ce{CaO}.
Amphoteric oxide: An oxide that reacts with both acids and bases, such as \ce{BeO}.
Peroxide: A compound containing the \ce{O2^2-} ion, such as \ce{BaO2}.
Flame test: A test in which a compound is heated in a non-luminous flame and identified by the colour its metal ions give.
Polarising power: The ability of a cation to distort the electron cloud of a nearby anion; high for small, highly charged ions such as \ce{Be^2+} and \ce{Mg^2+}.
Polarisation: The distortion of the electron cloud of an anion by a neighbouring cation, which gives the bond some covalent character.
Charge density: The charge of an ion divided by its size; small ions with a high charge have a high charge density.
Thermal decomposition: The breakdown of a compound into two or more different substances by heat.
Thermal stability: The resistance of a compound to decomposition on heating.
Lattice energy: The energy released when one mole of an ionic solid is formed from its gaseous ions; its reverse is the lattice dissociation enthalpy.
Enthalpy change of hydration: The enthalpy change when one mole of gaseous ions dissolves in water to give a very dilute solution; always exothermic.
Enthalpy change of solution: The enthalpy change when one mole of a solute dissolves in water to give a very dilute solution.
Complex ion: A central metal ion joined to ligands by coordinate bonds, such as \ce{[Be(H2O)4]^2+}.
Ligand: A molecule or ion that donates a lone pair of electrons to a central metal ion.
Lewis acid: An electron-pair acceptor, such as \ce{Be^2+} or \ce{BeCl2}.
Coordination number: The number of coordinate bonds formed by the central metal ion in a complex.
Reactivity series: A list of metals in order of decreasing reactivity, used to predict displacement and other reactions.
Diagonal relationship: The similarity between an element and the element diagonally below it to the right, such as Be and Al.
Quicklime: Calcium oxide, \ce{CaO}, made by heating limestone in a lime kiln.
Slaked lime: Calcium hydroxide, \ce{Ca(OH)2}, made by adding water to quicklime.
Barium meal: A suspension of \ce{BaSO4} swallowed before X-ray imaging of the gut.

Multiple choice questionsp. 92

i1 mark

The oxidation state shown by alkaline earth metals is

A+2Answer
B-2
C+1, +2
D-1, -2

Why

Every Group 2 atom has two valence electrons (ns^2) and loses both to form \ce{M^2+}, so the oxidation state is +2. Metals never show negative oxidation states in their compounds.

ii1 mark

Which one of the following is the most soluble in water?

A\ce{Mg(OH)2}
B\ce{Ca(OH)2}
C\ce{Sr(OH)2}
D\ce{Ba(OH)2}Answer

Why

Solubility of Group 2 hydroxides increases down the group (\ce{Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2}), because lattice energy falls faster than hydration enthalpy. Barium is lowest in the group.

iii1 mark

Which of the following alkaline earth metal hydroxides is amphoteric in character?

A\ce{Be(OH)2}Answer
B\ce{Sr(OH)2}
C\ce{Ca(OH)2}
D\ce{Ba(OH)2}

Why

\ce{Be(OH)2} reacts with acids (\ce{Be(OH)2 + 2HCl -> BeCl2 + 2H2O}) and with strong bases (\ce{Be(OH)2 + 2OH- -> [Be(OH)4]^2-}). The small, highly polarising \ce{Be^2+} ion causes this; the other hydroxides are purely basic.

iv1 mark

Of the metals Be, Mg, Ca and Sr of group 2A in the periodic table, the least ionic chloride would be formed by

ABeAnswer
BCa
CMg
DSr

Why

\ce{Be^2+} is the smallest ion with the highest charge density, so it polarises the chloride ion most and the bond has the most covalent character. \ce{BeCl2} is essentially covalent.

v1 mark

The order of solubility of sulphates of alkaline earth metals in water is

A\ce{Be > Mg > Ca > Sr > Ba}Answer
B\ce{Mg > Be > Ba > Ca > Sr}
C\ce{Be > Ca > Mg > Ba > Sr}
D\ce{Mg > Ca > Ba > Be > Sr}

Why

Sulphate solubility decreases steadily down the group: the large \ce{SO4^2-} ion keeps the lattice energy almost constant while the hydration enthalpy of the cation falls. \ce{BeSO4} is the most soluble and \ce{BaSO4} is insoluble.

vi1 mark

The solubilities of carbonates decrease down the magnesium group due to a decrease in

Ahydration energies of cationsAnswer
Binter-ionic attraction
Centropy of solution formation
Dlattice energies of solids

Why

The carbonate ion is large (like the sulphate ion), so the lattice energy changes very little down the group. The hydration energy of the cation falls considerably as the cation gets larger, so less energy is released on dissolving and solubility decreases.

Option D is tempting, but a decrease in lattice energy would make the carbonates more soluble, not less.

vii1 mark

In which of the following the hydration energy is higher than the lattice energy?

A\ce{MgSO4}Answer
B\ce{SrSO4}
C\ce{RaSO4}
D\ce{BaSO4}

Why

A salt dissolves when its hydration energy exceeds its lattice energy. \ce{MgSO4} is the only soluble sulphate listed: the small \ce{Mg^2+} ion has a large hydration energy. \ce{SrSO4}, \ce{BaSO4} and \ce{RaSO4} are insoluble.

viii1 mark

Which of the following alkaline earth metal sulphates has hydration enthalpy higher than the lattice enthalpy?

A\ce{CaSO4}
B\ce{BaSO4}
C\ce{BeSO4}Answer
D\ce{SrSO4}

Why

\ce{Be^2+} is the smallest ion, so its hydration enthalpy is the largest (about -2500\,\mathrm{kJ\,mol^{-1}}) and exceeds the lattice enthalpy of \ce{BeSO4}. That is why \ce{BeSO4} is very soluble while the others are sparingly soluble or insoluble.

ix1 mark

Which gas is released when \ce{CaCO3} reacts with dilute HCl?

A\ce{H2}
B\ce{O2}
C\ce{CO2}Answer
D\ce{Cl2}

Why

An acid reacts with a carbonate to give a salt, water and carbon dioxide: \ce{CaCO3 + 2HCl -> CaCl2 + CO2 + H2O}. Hydrogen is given only by a metal with an acid.

x1 mark

Brick red is characteristic flame colour of

ABeryllium
BMagnesium
CCalciumAnswer
DBarium

Why

Calcium compounds colour the flame brick-red. Barium gives apple-green, strontium scarlet red, and Be and Mg give no flame colour.

Short answer questionsp. 93

i3 marks

Describe and explain the trend observed in the thermal stabilities of the carbonates of Group II.

Solution

Trend: Group 2 carbonates decompose on heating to the oxide and carbon dioxide, \ce{MCO3 ->[heat] MO + CO2}. Thermal stability increases down the group: \ce{MgCO3} decomposes at about 540^\circ\mathrm{C} but \ce{BaCO3} needs about 1360^\circ\mathrm{C}.

\ce{BaCO3 > SrCO3 > CaCO3 > MgCO3}\quad\text{(stability)}

Explanation: the carbonate ion is large and easily polarised. All the cations are 2+, but they get larger down the group, so their charge density and polarising power decrease (\ce{Mg^2+ > Ca^2+ > Sr^2+ > Ba^2+}). The small \ce{Mg^2+} ion distorts the electron cloud of \ce{CO3^2-} strongly and weakens a C-O bond, so \ce{CO2} is released at a low temperature. The large \ce{Ba^2+} ion hardly distorts it, so a much higher temperature is needed.

ii3 marks

Describe the use of Group II elements or its compounds in agriculture.

Solution

Many soils become too acidic (from acid rain, fertilisers or decaying matter), and most crops grow poorly below about pH 6. Farmers add slaked lime, \ce{Ca(OH)2}, or powdered limestone, \ce{CaCO3}, to such soils. These are bases, so they neutralise the acid and raise the pH of the soil:

\ce{Ca(OH)2(s) + 2H+(aq) -> Ca^2+(aq) + 2H2O(l)}

\ce{CaCO3(s) + 2H+(aq) -> Ca^2+(aq) + CO2(g) + H2O(l)}

Liming also supplies calcium, a plant nutrient. Magnesium compounds such as magnesium sulphate (Epsom salt) are added as fertiliser to soils short of magnesium, which plants need to make chlorophyll.

iii3 marks

Write an equation to represent the thermal decomposition of calcium nitrate.

Solution

On heating, calcium nitrate decomposes to calcium oxide, nitrogen dioxide and oxygen:

\ce{2Ca(NO3)2(s) ->[heat] 2CaO(s) + 4NO2(g) + O2(g)}

Balancing check: left side 2 Ca, 4 N, 12 O; right side 2 Ca, 4 N, and 2+8+2=12 O. Observations: a white solid remains, a brown gas (\ce{NO2}, toxic) is given off, and the oxygen relights a glowing splint. Calcium nitrate decomposes at a higher temperature than magnesium nitrate but a lower one than barium nitrate, because thermal stability increases down the group.

iv3 marks

Why are the elements of Group 2 called alkaline earth metals?

Solution

They are called alkaline because their oxides and hydroxides dissolve in water (at least partly) to give alkaline solutions, for example \ce{CaO + H2O -> Ca(OH)2} and \ce{Ca(OH)2 -> Ca^2+ + 2OH-}, which turn red litmus blue.

They are called earth metals because their oxides were called 'earths' by early chemists: they are found in the Earth's crust (limestone, dolomite, gypsum) and are not changed by heat. Being metals that give such alkaline earths, Be, Mg, Ca, Sr, Ba and Ra are named the alkaline earth metals.

v3 marks

How do group 1 metals differ from Group 2 metals?

Solution

Property	Group 1	Group 2
Valence electrons, ion	ns^1, forms \ce{M+}	ns^2, forms \ce{M^2+}
Size and ionisation energy	larger atoms, lower IE	smaller atoms, higher IE
Reactivity	more reactive, stronger reducing agents	less reactive
Metallic bond	weaker: softer, lower m.p.	stronger: harder, higher m.p.
Hydroxides	very soluble, strong alkalis	less soluble, weaker bases
Carbonates on heating	stable (except \ce{Li2CO3})	decompose to \ce{MO + CO2}

Example: sodium reacts vigorously with cold water, \ce{2Na + 2H2O -> 2NaOH + H2}, while calcium reacts slowly and magnesium hardly at all. The main reason is the higher nuclear charge and smaller size of Group 2 atoms, which hold their valence electrons more strongly.

vi2026 Annual I, Old syllabus (Curriculum 2006), Q.2 (v) OR, old syllabus3 marks

Explain with the help of an equation the amphoteric nature of beryllium oxide.

Solution

An amphoteric oxide reacts with both acids and bases. Beryllium oxide shows both behaviours:

\ce{BeO(s) + 2HCl(aq) -> BeCl2(aq) + H2O(l)}

BeO acts as a base

\ce{BeO(s) + 2NaOH(aq) -> Na2BeO2(aq) + H2O(l)}

BeO acts as an acid

sodium beryllate; also written \ce{BeO + 2NaOH + H2O -> Na2[Be(OH)4]}

Reason: \ce{Be^2+} is very small with a high charge density, so it strongly polarises the oxide ion. The Be-O bond is largely covalent, so \ce{BeO} can accept hydroxide ions (forming \ce{[Be(OH)4]^2-}) as well as react with \ce{H+}. The other Group 2 oxides are ionic and purely basic.

vii3 marks

Explain qualitatively the variation in solubility of the hydroxides of the elements in Group II down the Group from magnesium to barium.

Solution

Trend: solubility increases down the group: \ce{Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2} (\ce{Mg(OH)2} is almost insoluble; \ce{Ba(OH)2} is quite soluble).

Explanation: whether a compound dissolves depends on \Delta H_{sol}=\Delta H_{latt\,diss}+\Delta H_{hyd}. Down the group the cation becomes larger.

Lattice energy depends on 1/(r_++r_-). The \ce{OH-} ion is small, so the sum changes a lot and lattice energy falls by a large amount.
Hydration enthalpy depends only on the cation (the anion is the same), and it falls by a smaller amount.

Since the energy needed to break the lattice falls faster than the energy released on hydration, \Delta H_{sol} becomes more exothermic down the group and the hydroxides become more soluble.

Long answer questionsp. 93

i8 marks

Explain the reactivity of group II elements with water and dilute HCl with chemical reactions.

Solution

Why reactivity increases down the group

Group 2 metals react by losing their two valence electrons: \ce{M -> M^2+ + 2e-}. Down the group the atomic radius and shielding increase, so the first and second ionisation energies decrease. The electrons are lost more easily, so the reactivity with water and acids increases from Be to Ba.

Reaction with water

Beryllium does not react with water or steam, because of a protective layer of \ce{BeO} on its surface.
Magnesium reacts extremely slowly with cold water (a few bubbles over days), because insoluble \ce{Mg(OH)2} coats it: \ce{Mg(s) + 2H2O(l) -> Mg(OH)2(s) + H2(g)}. Heated magnesium burns in steam with a bright white light, forming the oxide:

\ce{Mg(s) + H2O(g) -> MgO(s) + H2(g)}

Calcium reacts steadily with cold water, giving bubbles of hydrogen and a cloudy white suspension of slightly soluble calcium hydroxide. The solution is alkaline (pH about 12):

\ce{Ca(s) + 2H2O(l) -> Ca(OH)2(aq) + H2(g)}

Strontium and barium react faster and faster; barium reacts vigorously and gives a strongly alkaline solution:

\begin{aligned}\ce{Sr(s) + 2H2O(l) &-> Sr(OH)2(aq) + H2(g)}\\\ce{Ba(s) + 2H2O(l) &-> Ba(OH)2(aq) + H2(g)}\end{aligned}

General equation (Ca to Ba): \ce{M(s) + 2H2O(l) -> M(OH)2(aq) + H2(g)}. The resulting solutions become more alkaline down the group because the hydroxides become more soluble.

Reaction with dilute hydrochloric acid

All Group 2 metals react with dilute HCl to give the metal chloride and hydrogen. This is a redox reaction: the metal is oxidised to \ce{M^2+} and \ce{H+} is reduced to \ce{H2}.

\begin{aligned}\ce{M(s) + 2HCl(aq) &-> MCl2(aq) + H2(g)}\\\text{ionic: }\ce{M(s) + 2H+(aq) &-> M^2+(aq) + H2(g)}\end{aligned}

\begin{aligned}\ce{Mg(s) + 2HCl(aq) &-> MgCl2(aq) + H2(g)}\\\ce{Ca(s) + 2HCl(aq) &-> CaCl2(aq) + H2(g)}\end{aligned}

Beryllium reacts slowly (oxide layer). Magnesium reacts rapidly with fizzing; calcium, strontium and barium react more and more vigorously and exothermically. The rate of hydrogen evolution increases down the group.

Summary

Metal	Cold water	Dilute HCl
Be	no reaction	slow
Mg	very slow (reacts with steam)	rapid
Ca	steady	very rapid
Sr	fast	violent
Ba	vigorous	violent

ii2026 Annual I, Set B, Q.5 OR2026 Annual I, Set D, Q.3 OR8 marks

Analyse the trends in the thermal stability of nitrates and carbonates of alkaline earth metals.

Solution

Thermal decomposition

Thermal decomposition is the breakdown of a compound into simpler substances by heat. Group 2 carbonates and nitrates both decompose on heating to give the metal oxide.

Carbonates

\ce{MCO3(s) ->[heat] MO(s) + CO2(g)}\qquad\text{e.g. }\ce{CaCO3(s) ->[heat] CaO(s) + CO2(g)}

Carbonate	Decomposition temperature / ^\circ\mathrm{C}	\Delta H / \mathrm{kJ\,mol^{-1}}
\ce{MgCO3}	540	+117
\ce{CaCO3}	900	+176
\ce{SrCO3}	1280	+238
\ce{BaCO3}	1360	+268

The decomposition temperature and the enthalpy change both increase down the group, so thermal stability increases.

Nitrates

\begin{aligned}\ce{2M(NO3)2(s) &->[heat] 2MO(s) + 4NO2(g) + O2(g)}\\\text{e.g. }\ce{2Mg(NO3)2 &->[heat] 2MgO + 4NO2 + O2}\end{aligned}

A brown gas (\ce{NO2}) is seen and oxygen is released. Here too a higher temperature is needed further down the group.

Reason: polarisation of the anion by the cation

The small \ce{Mg^2+} ion pulls the electron cloud of the large \ce{CO3^2-} ion towards itself (polarisation). The larger \ce{Ca^2+} ion distorts it much less (after the book's Fig 5.4).

Derivation Explaining the stability trend

Aim. Relate the thermal stability to the size of the cation.

The carbonate and nitrate anions are large, so their electron clouds are easily distorted (polarised).
All Group 2 cations carry a 2+ charge, but their radius increases down the group, so their charge density and polarising power decrease.
\text{polarising power: }\ce{Mg^2+ > Ca^2+ > Sr^2+ > Ba^2+}
A small cation such as \ce{Mg^2+} pulls electron density from the anion towards itself. This weakens a C-O bond in \ce{CO3^2-} (or an N-O bond in \ce{NO3-}), so the anion breaks up easily into \ce{CO2} (or \ce{NO2} and \ce{O2}) and an oxide ion.
A large cation such as \ce{Ba^2+} hardly distorts the anion, so its bonds stay strong and much more heat is needed.

Result\text{Thermal stability increases down the group}

The less polarising the cation, the more stable its carbonate or nitrate.

Order of stability

\ce{BaCO3 > SrCO3 > CaCO3 > MgCO3}

\ce{Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2}

FBISE 2026 questions

Questions from the FBISE HSSC-II 2026 (1st Annual) papers that belong to this unit, answered the way an examiner expects.

MCQ 42026 Annual I, Set B, MCQ 41 mark

Which compound is least soluble in water?

A\ce{MgSO4}
B\ce{CaSO4}
C\ce{SrSO4}
D\ce{BaSO4}Answer

Why

Group 2 sulphates become less soluble down the group, because the hydration enthalpy of the cation falls more than the lattice energy (the large sulphate ion keeps the lattice energy nearly constant). Barium is lowest, so \ce{BaSO4} is the least soluble (about 10^{-5}\,\mathrm{mol\,dm^{-3}}).

Q.2 (iv)2026 Annual I, Set B, Q.2 (iv)1+1+1 = 3 marks

State the type of ions formed by Group II elements. Mention their charge, and compare their reactivity with that of Group I elements.

Solution

Type of ions (1 mark)

Group 2 elements are metals with the valence configuration ns^2. They lose both valence electrons and form positive ions (cations) with a noble gas configuration: \ce{M -> M^2+ + 2e-}, for example \ce{Mg -> Mg^2+ + 2e-}. Their compounds are ionic (except those of Be).

Charge (1 mark)

The charge on each ion is \mathbf{2+} (\ce{Be^2+}, \ce{Mg^2+}, \ce{Ca^2+}, \ce{Sr^2+}, \ce{Ba^2+}), so their oxidation state is +2.

Reactivity compared with Group I (1 mark)

Group 2 elements are less reactive than the Group 1 elements of the same period. Their atoms are smaller and have a greater nuclear charge, so their ionisation energies are higher (Mg 738, Na 496\,\mathrm{kJ\,mol^{-1}}), and they must lose two electrons instead of one. For example, sodium reacts vigorously with cold water, \ce{2Na + 2H2O -> 2NaOH + H2}, but magnesium reacts only very slowly.

Q.2 (v)2026 Annual I, Set B, Q.2 (v)1+1+1 = 3 marks

Compare with the help of chemical equations the reactivity of Li and Be with oxygen.

Solution

Lithium with oxygen (1 mark)

Lithium (Group 1) burns readily when heated in air or oxygen, with a crimson flame, forming the white normal oxide \ce{Li2O}:

\ce{4Li(s) + O2(g) -> 2Li2O(s)}

Beryllium with oxygen (1 mark)

Beryllium (Group 2) is much less reactive. The massive metal is protected by a thin layer of \ce{BeO}; it burns only when powdered and strongly heated, forming beryllium oxide:

\ce{2Be(s) + O2(g) ->[strong heat] 2BeO(s)}

Comparison and reason (1 mark)

Lithium is more reactive with oxygen than beryllium. Li loses one electron (first IE 520\,\mathrm{kJ\,mol^{-1}}) while Be must lose two electrons from a smaller atom (first IE 900, second IE 1757\,\mathrm{kJ\,mol^{-1}}). Both form only normal oxides (no peroxides), because their small ions cannot stabilise the large peroxide ion; \ce{Li2O} is basic while \ce{BeO} is amphoteric.

Note. Lithium belongs to Group 1, but the question compares it with beryllium, so it is answered here with the Group 2 reactions.

Q.5 OR2026 Annual I, Set B, Q.5 OR6 marks

Explain the trend in thermal stability of Group II carbonates and nitrates, giving reasons. Also, write the order of stability for each.

Solution

Trend (2 marks)

Group 2 carbonates and nitrates decompose on heating to the metal oxide. Thermal stability increases down the group: a higher temperature is needed to decompose the compound as the group is descended.

\begin{aligned}\ce{MCO3(s) &->[heat] MO(s) + CO2(g)}\\\ce{2M(NO3)2(s) &->[heat] 2MO(s) + 4NO2(g) + O2(g)}\end{aligned}

For example, \ce{MgCO3} decomposes at about 540^\circ\mathrm{C}, \ce{CaCO3} at 900^\circ\mathrm{C}, \ce{SrCO3} at 1280^\circ\mathrm{C} and \ce{BaCO3} at 1360^\circ\mathrm{C}. Nitrates give off brown \ce{NO2} gas and oxygen; \ce{Mg(NO3)2} decomposes most easily and \ce{Ba(NO3)2} least easily.

Reasons (3 marks)

The carbonate ion \ce{CO3^2-} and nitrate ion \ce{NO3-} are large anions whose electron clouds are easily polarised (distorted).
All Group 2 cations have a 2+ charge, but their size increases down the group (\ce{Mg^2+ < Ca^2+ < Sr^2+ < Ba^2+}), so their charge density and polarising power decrease.
A small cation like \ce{Mg^2+} pulls electron density towards itself and weakens a C-O (or N-O) bond in the anion, so the anion breaks up at a low temperature, releasing \ce{CO2} (or \ce{NO2} and \ce{O2}).
A large cation like \ce{Ba^2+} hardly distorts the anion, so its bonds remain strong and much more heat is needed. The enthalpy of decomposition also becomes more positive down the group (+117 for \ce{MgCO3} to +268\,\mathrm{kJ\,mol^{-1}} for \ce{BaCO3}).

The small \ce{Mg^2+} ion pulls the electron cloud of the large \ce{CO3^2-} ion towards itself (polarisation). The larger \ce{Ca^2+} ion distorts it much less (after the book's Fig 5.4).

Order of stability (1 mark)

\text{Carbonates: }\ce{BaCO3 > SrCO3 > CaCO3 > MgCO3}

\text{Nitrates: }\ce{Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2}

Note. The paper gives 6 marks for the whole question without a split; the headings show a sensible division.

MCQ 52026 Annual I, Set D, MCQ 51 mark

The reactivity of Group 2 elements ____ down the group:

ARemains constant
BFirst increases then decreases
CDecreases
DIncreasesAnswer

Why

Down the group the atoms get larger and the ionisation energies decrease, so the two outer electrons are lost more easily. Reactivity with oxygen, water and acids therefore increases from Be to Ba.

Q.2 (iv)2026 Annual I, Set D, Q.2 (iv)1+2 = 3 marks

Compare the ionization energies of Group 1 and Group 2 elements. Explain the reason for the difference.

Solution

Comparison (1 mark)

The first ionisation energies of Group 2 elements are higher than those of the Group 1 elements in the same period. Examples: Li 520 vs Be 900\,\mathrm{kJ\,mol^{-1}}; Na 496 vs Mg 738\,\mathrm{kJ\,mol^{-1}}; K 419 vs Ca 590\,\mathrm{kJ\,mol^{-1}}. (The second IE of Group 1 is, however, far higher than that of Group 2, because it removes an electron from a noble gas core.)

Reason (2 marks)

Greater nuclear charge: a Group 2 atom has one more proton than the Group 1 atom before it, while the added electron enters the same shell and shields poorly. The valence electrons feel a larger effective nuclear charge.
Smaller atomic size: the stronger attraction pulls the electrons closer, so Group 2 atoms are smaller (Mg 0.160\,\mathrm{nm}, Na 0.186\,\mathrm{nm}) and the outer electrons are held more tightly.
Stable configuration: Group 2 atoms have a completely filled ns^2 subshell, which is more stable than the half-filled ns^1 subshell of Group 1.

So more energy is needed to remove an electron from a Group 2 atom. This is why Group 2 metals are less reactive and weaker reducing agents than Group 1 metals.

Q.2 (v)2026 Annual I, Set D, Q.2 (v)1x3 = 3 marks

Write any three uses of group-II elements in industry, agriculture or medicines.

Solution

Industry (1 mark): limestone, \ce{CaCO3}, is used as a building material and is heated in a lime kiln to make quicklime, which is roasted with clay to make cement: \ce{CaCO3 ->[heat] CaO + CO2}. Magnesium is used in flares and in light alloys for aircraft.
Agriculture (1 mark): slaked lime, \ce{Ca(OH)2}, is spread on acidic soils to neutralise the acid and raise the pH: \ce{Ca(OH)2 + 2H+ -> Ca^2+ + 2H2O}.
Medicine (1 mark): barium sulphate, \ce{BaSO4}, is given as a 'barium meal' for X-ray pictures of the stomach and intestines; it absorbs X-rays and is too insoluble to be absorbed into the blood. Magnesium hydroxide (Milk of Magnesia) and calcium carbonate are antacids: \ce{Mg(OH)2 + 2HCl -> MgCl2 + 2H2O}.

Q.3 OR2026 Annual I, Set D, Q.3 OR2+3+2 = 7 marks

What is the trend in thermal stability of Group II carbonates and nitrates? Give reasons. Also, write the order of stability for each.

Solution

Trend (2 marks)

\begin{aligned}\ce{MCO3(s) &->[heat] MO(s) + CO2(g)}\\\ce{2M(NO3)2(s) &->[heat] 2MO(s) + 4NO2(g) + O2(g)}\end{aligned}

Reasons (3 marks)

The carbonate ion \ce{CO3^2-} and nitrate ion \ce{NO3-} are large anions whose electron clouds are easily polarised (distorted).
All Group 2 cations have a 2+ charge, but their size increases down the group (\ce{Mg^2+ < Ca^2+ < Sr^2+ < Ba^2+}), so their charge density and polarising power decrease.
A small cation like \ce{Mg^2+} pulls electron density towards itself and weakens a C-O (or N-O) bond in the anion, so the anion breaks up at a low temperature, releasing \ce{CO2} (or \ce{NO2} and \ce{O2}).
A large cation like \ce{Ba^2+} hardly distorts the anion, so its bonds remain strong and much more heat is needed. The enthalpy of decomposition also becomes more positive down the group (+117 for \ce{MgCO3} to +268\,\mathrm{kJ\,mol^{-1}} for \ce{BaCO3}).

The small \ce{Mg^2+} ion pulls the electron cloud of the large \ce{CO3^2-} ion towards itself (polarisation). The larger \ce{Ca^2+} ion distorts it much less (after the book's Fig 5.4).

Order of stability (2 marks)

\text{Carbonates: }\ce{BaCO3 > SrCO3 > CaCO3 > MgCO3}

\text{Nitrates: }\ce{Ba(NO3)2 > Sr(NO3)2 > Ca(NO3)2 > Mg(NO3)2}

MCQ 22026 Annual I, Old syllabus (Curriculum 2006), MCQ 2, old syllabus1 mark

Which of the following nitrates decomposes at the lowest temperature?

A\ce{Ba(NO3)2}
B\ce{Sr(NO3)2}
C\ce{Mg(NO3)2}Answer
D\ce{Ca(NO3)2}

Why

Thermal stability of Group 2 nitrates increases down the group. \ce{Mg^2+} is the smallest and most polarising of the four cations, so it weakens the N-O bonds of the nitrate ion most, and \ce{Mg(NO3)2} decomposes at the lowest temperature.

Q.2 (i)2026 Annual I, Old syllabus (Curriculum 2006), Q.2 (i), old syllabus1.5+1.5 = 3 marks

How does the solubility of \ce{MgSO4} compare with \ce{BaSO4}?

Solution

Comparison (1.5 marks)

\ce{MgSO4} is soluble in water (1.83\,\mathrm{mol\,dm^{-3}}), while \ce{BaSO4} is practically insoluble (9.43\times10^{-6}\,\mathrm{mol\,dm^{-3}}). Solubility of Group 2 sulphates decreases down the group: \ce{MgSO4 > CaSO4 > SrSO4 > BaSO4}.

Reason (1.5 marks)

The sulphate ion is very large, so the lattice energies of \ce{MgSO4} and \ce{BaSO4} are nearly the same (r_++r_- hardly changes). But the small \ce{Mg^2+} ion has a much larger hydration enthalpy (about -1920\,\mathrm{kJ\,mol^{-1}}) than the large \ce{Ba^2+} ion (about -1360\,\mathrm{kJ\,mol^{-1}}). For \ce{MgSO4} the hydration energy overcomes the lattice energy and \Delta H_{sol} is exothermic; for \ce{BaSO4} it does not, so \Delta H_{sol} is endothermic and it does not dissolve.

Q.2 (iii)2026 Annual I, Old syllabus (Curriculum 2006), Q.2 (iii), old syllabus1+1+1 = 3 marks

Compare the melting and boiling points of beryllium with magnesium. Explain the reason for the difference.

Solution

Melting points (1 mark)

Beryllium has a much higher melting point (about 1280^\circ\mathrm{C}) than magnesium (650^\circ\mathrm{C}).

Boiling points (1 mark)

Beryllium also has a much higher boiling point (about 2470^\circ\mathrm{C}) than magnesium (about 1090^\circ\mathrm{C}).

Reason (1 mark)

Both are held together by metallic bonds, each atom giving two electrons to the sea of delocalised electrons. The beryllium atom is much smaller (0.122\,\mathrm{nm} vs 0.160\,\mathrm{nm}), so its nucleus is closer to the delocalised electrons and attracts them more strongly. The metallic bond in Be is therefore much stronger, and more energy is needed to melt and boil it.

Q.2 (v) OR2026 Annual I, Old syllabus (Curriculum 2006), Q.2 (v) OR, old syllabus1.5+1.5 = 3 marks

Write two balanced chemical equations for the amphoteric behaviour of BeO.

Solution

An amphoteric oxide reacts with both an acid and a base. \ce{BeO} is amphoteric because the small, highly polarising \ce{Be^2+} ion gives the Be-O bond considerable covalent character.

With an acid: BeO acts as a base (1.5 marks)

\ce{BeO(s) + 2HCl(aq) -> BeCl2(aq) + H2O(l)}

With a base: BeO acts as an acid (1.5 marks)

\ce{BeO(s) + 2NaOH(aq) -> Na2BeO2(aq) + H2O(l)}

sodium beryllate; in water it is the complex \ce{Na2[Be(OH)4]}: \ce{BeO + 2NaOH + H2O -> Na2[Be(OH)4]}

Exam focus

Explain the trend in thermal stability of Group II carbonates and nitrates, with reasons and orders of stability. (long)
Asked in both 2026 new-syllabus sets (Set B Q.5 OR, Set D Q.3 OR) and in the book's long questions.
Go to answer
Explain the reactivity of Group II elements with water and dilute HCl, with equations. (long)
Book long question; tests equations and the ionisation energy argument together.
Go to answer
Explain the amphoteric nature of \ce{BeO} with equations. (short)
Asked in 2026 (old syllabus Q.2 (v) OR) and in the book's short questions.
Go to answer
Compare the ionisation energies and reactivity of Group 1 and Group 2 elements. (short)
Two 2026 questions (Set B Q.2 (iv), Set D Q.2 (iv)) test this comparison.
Go to answer
Explain the variation in solubility of Group 2 hydroxides using lattice and hydration enthalpies. (short)
A book short question and a named SLO; needs the lattice versus hydration argument.
Go to answer
Why is \ce{MgSO4} more soluble than \ce{BaSO4}? (short)
Asked in 2026 (old syllabus Q.2 (i)); the sulphate trend is the opposite of the hydroxide trend.
Go to answer
Three uses of Group II compounds in industry, agriculture and medicine. (short)
Asked in 2026 (Set D Q.2 (v)) and in the book (agriculture).
Go to answer
Compare the reactivity of Li and Be with oxygen, with equations. (short)
Asked in 2026 (Set B Q.2 (v)); combines Group 1 and Group 2 chemistry.
Go to answer
Write the equation for the thermal decomposition of calcium nitrate. (short)
Book short question; balancing the nitrate equation is a common slip.
Go to answer
MCQ trends: solubility of sulphates and hydroxides, nitrate decomposition, reactivity down the group, flame colours. (mcq)
Four 2026 MCQs came from these facts.
Go to answer