Group 2

The Group 2 metals: two outer s electrons, and the physical trends down the group

Group 2 contains beryllium, magnesium, calcium, strontium, barium and radium; the AS course runs from magnesium to barium. Every Group 2 atom has two electrons in its outer s subshell, so the group sits in the s-block and its outer configuration is written ns2 - magnesium is [Ne] 3s2 and calcium is [Ar] 4s2. Both of those electrons are lost to form a 2+ ion: M -> M2+ + 2e-, and the ion left behind has the noble gas configuration of the element before it, so Mg2+ is [Ne] and Ba2+ is [Xe]. A third electron is never removed in a chemical reaction, because it would have to come out of a complete inner shell: the third ionisation energy of every Group 2 element is far higher than the second, so M2+ is the only ion you meet in compounds.

Because the ions have no partly filled d subshell and no unpaired electrons to absorb visible light, Group 2 compounds are white or colourless, and the compounds in this topic are all ionic. Magnesium oxide, calcium carbonate and barium sulfate are white solids, and solutions of magnesium chloride and barium hydroxide are colourless - a useful clue when you are identifying an unknown salt. Only beryllium, the first member of the group, forms largely covalent compounds, and you do not need it for any of the reactions here.

Going down the group the nuclear charge increases, but each element also has an extra filled shell of electrons. The extra shells add shielding and make the atom bigger, so the outer electrons are further from the nucleus and the pull on them falls, even though the charge on the nucleus is larger. That one idea explains the physical trends: atomic radius increases, the radius of the 2+ ion increases with it, and the first ionisation energy and the electronegativity both fall.

Density rises overall down the group, because the mass of each atom increases much faster than the volume it occupies, but calcium is an exception: it is less dense than magnesium because its atoms pack in a different crystal structure. Melting points are high - evidence of strong metallic bonding between the M2+ ions and the delocalised electrons - and they fall overall down the group as the ions get bigger and the metallic bond weakens, although calcium does not sit neatly in the trend because it melts higher than magnesium.

Reactions of the elements with oxygen, water and steam

With oxygen every Group 2 metal burns to give the oxide: 2Mg(s) + O2(g) -> 2MgO(s), and 2Ca(s) + O2(g) -> 2CaO(s). Magnesium is the classic demonstration, burning with a brilliant white flame and leaving a white powder of magnesium oxide; calcium, strontium and barium burn in the same way. The oxides are basic oxides. Barium heated in a large excess of oxygen also forms barium peroxide, BaO2, a useful reminder that the group is not perfectly uniform.

With cold water the metals give the hydroxide and hydrogen: M(s) + 2H2O(l) -> M(OH)2(aq) + H2(g). Magnesium reacts extremely slowly with cold water, and the reaction soon appears to stop, because the magnesium hydroxide that forms is only sparingly soluble and sticks to the metal as a thin coating that keeps the water away from the surface. Calcium reacts steadily. The metal sinks and bubbles of hydrogen stream off it, and a white cloud of sparingly soluble calcium hydroxide forms. Strontium reacts vigorously and barium violently, and the resulting solutions are strongly alkaline, so universal indicator turns blue or purple.

With steam magnesium reacts quite differently, and this is a favourite exam trap: Mg(s) + H2O(g) -> MgO(s) + H2(g). The reaction is fast enough for the metal to glow brightly, and the solid left behind is magnesium oxide, not magnesium hydroxide. The temperature of the steam is high enough to decompose the hydroxide as fast as it is formed: Mg(OH)2(s) -> MgO(s) + H2O(g). The hydrogen can be collected and tested with a lighted splint, which gives the usual squeaky pop.

Taken together, the reactions with water show that reactivity increases from magnesium to barium. The reason is the fall in the first and second ionisation energies down the group: the atoms are bigger, their outer electrons are more shielded, and both electrons are lost more easily, so the metal is oxidised from M to M2+ more readily.

Reactions of the elements with dilute hydrochloric and sulfuric acids

With dilute hydrochloric acid the metals dissolve with effervescence to give the chloride and hydrogen: M(s) + 2HCl(aq) -> MCl2(aq) + H2(g). Magnesium reacts briskly, calcium faster still, and barium reacts violently. The solutions that form are colourless, because the chlorides of all four metals are soluble, and the gas is hydrogen, which burns with a squeaky pop. The vigour increases down the group, for exactly the same reason as the reaction with water: the ionisation energies fall.

With dilute sulfuric acid the equation is the same shape, M(s) + H2SO4(aq) -> MSO4(aq) + H2(g), but the pattern of reactivity is reversed. Magnesium reacts steadily, because magnesium sulfate is very soluble, so the metal surface is continually washed clean and the acid keeps reaching it. Calcium reacts at first and then slows down and stops: calcium sulfate is only slightly soluble, so a layer of it forms on the metal and blocks the acid from reaching the surface below. Strontium sulfate is even less soluble and barium sulfate is essentially insoluble, so strontium reacts very little and barium barely reacts at all with dilute sulfuric acid once the first thin layer of barium sulfate has formed.

This is an important lesson in how to read a trend. The intrinsic reactivity of the metal still increases down the group, but the rate you observe in the test tube depends on whether the salt formed stays in solution. Where the salt is insoluble, the reaction stops no matter how reactive the metal is. That is why chemists use dilute hydrochloric acid rather than dilute sulfuric acid when they want a Group 2 metal or carbonate to react to completion.

If you are asked to write these equations, check two things: the 2+ charge on the metal ion (so two acid molecules are needed per metal atom), and the state symbols, which must show the hydrogen as a gas and the acid as aqueous.

Oxides and hydroxides: basicity, water and dilute acids

The oxides are basic. They react with water to give the hydroxide and with dilute acids to give a salt and water. Basicity and speed of reaction with water both increase down the group. Magnesium oxide reacts only slowly with water and gives a weakly alkaline suspension, because MgO is only slightly soluble. Calcium oxide, by contrast, reacts vigorously with water and the mixture gets very hot - the reaction is so exothermic that it is called slaking, and the calcium hydroxide produced is slaked lime: CaO(s) + H2O(l) -> Ca(OH)2(s). Strontium and barium oxides behave like calcium oxide. Barium hydroxide solution is a strong alkali.

With dilute hydrochloric acid, an oxide gives the chloride and water: MgO(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l). With dilute sulfuric acid the same pattern applies, but watch the solubility of the salt: BaO(s) + H2SO4(aq) -> BaSO4(s) + H2O(l) produces a white precipitate of insoluble barium sulfate, whereas MgO gives the soluble magnesium sulfate.

The hydroxides neutralise acids to give a salt and water: Mg(OH)2(s) + 2HCl(aq) -> MgCl2(aq) + 2H2O(l), Ca(OH)2(aq) + H2SO4(aq) -> CaSO4(s) + 2H2O(l) and Ba(OH)2(aq) + 2HCl(aq) -> BaCl2(aq) + 2H2O(l). Because they contain two hydroxide ions per formula unit, one mole of a hydroxide neutralises two moles of a monobasic acid such as hydrochloric acid.

Solubility of the hydroxides increases down the group - Mg(OH)2 is sparingly soluble, Ca(OH)2 is slightly soluble, and Sr(OH)2 and Ba(OH)2 are increasingly soluble. You can see the consequence in the classroom: magnesium hydroxide is used as a suspension (milk of magnesia) because so little of it dissolves, calcium hydroxide is used as a saturated solution (limewater), and barium hydroxide can be used as a strong alkali in a burette, because enough of it dissolves to give a genuinely alkaline solution.

• MgO: only slightly soluble, reacts slowly with water, weakly basic.
• CaO: reacts vigorously and very exothermically with water (slaking) to give slaked lime.
• SrO and BaO: react vigorously with water, and Ba(OH)2 solution is a strong alkali.
• Oxide + dilute hydrochloric acid -> chloride + water.
• Hydroxide + dilute acid -> salt + water, one mole of M(OH)2 neutralising two moles of HCl.

Carbonates: insoluble in water, but they react with dilute acids

The Group 2 carbonates are insoluble in water. You cannot make an alkaline solution from magnesium carbonate or calcium carbonate the way you can from calcium oxide, because almost none of the solid dissolves. They do, however, react readily with dilute acids, and this is the reaction behind the standard laboratory test for carbonates.

With dilute hydrochloric acid every Group 2 carbonate fizzes and dissolves: MCO3(s) + 2HCl(aq) -> MCl2(aq) + H2O(l) + CO2(g), for example CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + H2O(l) + CO2(g). The gas is carbon dioxide, and it is identified by bubbling it through limewater, a saturated solution of calcium hydroxide, which it turns milky: Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l). If you keep bubbling, the milkiness disappears again as the carbonate is converted to the soluble hydrogencarbonate, CaCO3(s) + H2O(l) + CO2(g) -> Ca(HCO3)2(aq).

With dilute sulfuric acid the reaction starts but does not run to completion for calcium, strontium and barium. Calcium carbonate fizzes for a moment and then stops, because the sparingly soluble calcium sulfate formed on the surface seals the solid off from the acid. Strontium carbonate and barium carbonate behave in the same way, and more decisively, because their sulfates are even less soluble. Only magnesium carbonate reacts steadily with dilute sulfuric acid, because magnesium sulfate is soluble. That is why dilute hydrochloric acid, not sulfuric acid, is used to test for a carbonate or to dissolve a carbonate in analysis.

The chemistry of carbonates and carbon dioxide is not confined to the laboratory: the same two reactions - dissolving limestone with carbonic acid and precipitating calcium carbonate from limewater - are what shape limestone caves.

Thermal decomposition of the carbonates and nitrates

On strong heating, a Group 2 carbonate decomposes to the oxide and carbon dioxide: MCO3(s) -> MO(s) + CO2(g), for example CaCO3(s) -> CaO(s) + CO2(g). A Group 2 nitrate decomposes to the oxide, brown nitrogen dioxide and oxygen: 2M(NO3)2(s) -> 2MO(s) + 4NO2(g) + O2(g), for example 2Mg(NO3)2(s) -> 2MgO(s) + 4NO2(g) + O2(g). The brown gas turns damp blue litmus paper red because it forms an acid with water, and the oxygen relights a glowing splint. Note that every Group 2 nitrate gives the oxide, unlike sodium and potassium nitrates, which give a nitrite and oxygen.

Thermal stability increases down the group: the temperature needed to decompose the carbonate or the nitrate rises from magnesium to barium. Magnesium carbonate breaks down on quite gentle heating, calcium carbonate needs the full heat of a roaring Bunsen flame in a lime kiln, and the strontium and barium salts need much higher temperatures still.

The explanation is the polarising power of the cation. A cation pulls electron density away from the large anion next to it, distorting it. The higher the charge density of the cation - the more charge packed into a small volume - the stronger that distortion. A strongly polarised carbonate ion has one of its C-O bonds weakened and lengthened, so less energy is needed to break the ion apart and release the carbon dioxide; in a nitrate it is an N-O bond that is weakened. Going down the group the ionic radius increases, so the charge density and the polarising power fall, the anion is distorted less, and a higher temperature is required.

Notice what is not happening. The bond that breaks is inside the anion, not the bond between the metal ion and the anion, so an argument about the strength of the metal-oxygen bond will not score. The industrial version of the reaction is the lime kiln, where limestone is decomposed at about 1000 °C to make quicklime.

Solubility trends: hydroxides become more soluble, sulfates less

The hydroxides become more soluble down the group: Mg(OH)2 is sparingly soluble, Ca(OH)2 is slightly soluble - a saturated solution at room temperature holds only about 1.7 g dm-3, and that solution is limewater - Sr(OH)2 is more soluble, and Ba(OH)2 is the most soluble of the four, giving a strongly alkaline solution. The sulfates become less soluble down the group: MgSO4 is very soluble, CaSO4 is only slightly soluble, and SrSO4 and BaSO4 are essentially insoluble.

Both trends come from the same balance. An ionic solid dissolves when the energy released by hydrating its ions as they separate (the hydration enthalpy) is at least as large as the energy needed to pull the lattice apart (the lattice enthalpy). Going down the group the cation gets bigger, so its hydration enthalpy becomes less exothermic, and the lattice enthalpy also falls, because the ions are further apart.

Which way the balance tips depends on how fast each quantity falls. For the hydroxides, the lattice enthalpy falls more quickly than the hydration enthalpy of the cation, so dissolving becomes more favourable and solubility increases. For the sulfates, the situation is different because the sulfate ion is large: the hydration enthalpy of the cation falls steeply as the cation gets bigger (the small Mg2+ ion has a high charge density and is strongly hydrated, the large Ba2+ ion much less so), while the lattice enthalpy of the sulfate changes comparatively little. Dissolving therefore becomes less favourable and solubility decreases.

These two trends have practical consequences you should be able to quote. Limewater works as a test for carbon dioxide because calcium hydroxide is soluble enough to give a usable solution but its carbonate is insoluble, so the precipitate forms. Barium sulfate can be swallowed as a barium meal before an X-ray because it is essentially insoluble: no Ba2+ ions are released into the body, and barium ions are poisonous. Soluble barium salts such as barium chloride must therefore be handled with care, and barium chloride solution is used in the standard test for sulfate ions: acidify the unknown solution with dilute hydrochloric acid and add barium chloride solution, when a white precipitate of BaSO4 forms that does not dissolve in the acid: Ba2+(aq) + SO4^2-(aq) -> BaSO4(s).

Using the trends to make predictions

An exam question rarely asks you to recall one fact on its own. It gives you a Group 2 element or compound - sometimes one you have not been asked about before, such as strontium or radium - and asks you to predict what it will do. To do that, sort the question into one of the trends below and apply the trend in the right direction.

For the elements: reactivity with oxygen and water increases down the group, because the ionisation energies fall. For the oxides and hydroxides: solubility and basicity increase down the group, so a barium compound gives the most alkaline solution. For the carbonates and nitrates: thermal stability increases down the group, so a magnesium salt decomposes most easily and a barium salt needs the strongest heating. For the sulfates: solubility decreases down the group, so barium sulfate is the least soluble and magnesium sulfate the most soluble.

Then check the reaction against solubility. Any time a Group 2 compound reacts with dilute sulfuric acid, ask whether the sulfate that would form is soluble: if it is not, a coating will form and the reaction will slow or stop. Any time a test involves a white precipitate, ask whether that precipitate is soluble in dilute acid, because that decides whether it can be used to confirm the presence of the ion.

Finally, always give the reason in terms of particles. Say which property changes (charge density, ionisation energy, lattice enthalpy, hydration enthalpy) and which way it changes down the group. A prediction with the correct direction but no reason usually scores only the first mark.

• Elements: reactivity with oxygen and water increases down the group.
• Oxides and hydroxides: solubility and basicity increase down the group.
• Carbonates and nitrates: thermal stability increases down the group.
• Sulfates: solubility decreases down the group.
• Before predicting a reaction with dilute sulfuric acid, decide whether the sulfate formed is soluble.