Group 2, the alkaline earth metals

In a nutshell

The alkaline earth metals are Group 2: the elements you need are magnesium to barium (Mg-Ba). They are reactive metals that all lose their two outer electrons to form 2+ ions.

This subtopic is really two stories: the physical trends down the group (atomic radius, first ionisation energy, melting point) and the patterns in reactions and solubility (reactions with water, the opposite solubility trends of the hydroxides and sulfates, and the uses and the sulfate test that come out of them).

Assumed knowledge: Atomic structure, Bonding, Periodicity.

Core content

Atomic radius down the group

Atomic radius increases down the group (Mg to Ba).

Going down, each element has one more shell of electrons, so the outer electrons are further from the nucleus and the atom is larger.

The extra nuclear charge does not win here because the added inner shells also add more shielding, so the effect of the extra protons on the outer shell is largely cancelled.

First ionisation energy down the group

First ionisation energy decreases down the group.

The first ionisation energy is the energy to remove one mole of electrons from one mole of gaseous atoms:

$Mg(g) \rightarrow Mg^{+}(g) + e^{-}$

Three linked reasons, and the exam wants them compared down the group:

• The outer electron is in a shell that is further from the nucleus.
• There are more inner shells, so the outer electron is more shielded from the nuclear charge.
• Both of these outweigh the increase in nuclear charge, so the outer electron is less strongly attracted and easier to remove.

Still don't get it? · why ionisation energy falls even though there are more protons

Think of the nucleus as a magnet and the outer electron as a paperclip you are trying to pull off.

Going down the group you add protons, so you might expect the magnet to get stronger and the paperclip harder to pull off. But two other things change at the same time.

First, each step down adds a whole new layer, so the paperclip sits further away. A magnet's pull drops off fast with distance, so distance alone makes it easier to remove.

Second, the inner layers of electrons sit between the nucleus and the outer electron and push it away (they "shield" it). More inner layers means more of this pushing-out, so the outer electron barely feels the extra protons.

Put together: the outer electron is further away and better shielded, and those two effects beat the extra nuclear charge. So the attraction on the outer electron gets weaker down the group, and the ionisation energy decreases. In the exam, say all three parts and make them comparative: further from the nucleus, more shielding, and these outweigh the greater nuclear charge.

Melting point down the group

The Group 2 metals are giant metallic lattices: a regular array of positive ions in a sea of delocalised electrons, held together by the attraction between them (metallic bonding).

The general trend is that melting point decreases from Mg to Ba. As the ions get larger down the group:

• the delocalised electrons are further from the positive ions,
• so the attraction between the positive ions and the delocalised electrons is weaker,
• meaning weaker metallic bonding and a lower melting point.

Magnesium is the exception: its melting point is lower than calcium's, because the metals do not all pack in the same crystal structure, so the simple size argument only runs cleanly from Ca to Ba.

Two phrases are worth more than the trend itself, because they are exactly where the mark is lost:

• The number of delocalised electrons is the same for every Group 2 metal (each atom releases two). So never explain the trend by saying there are "more" or "fewer" delocalised electrons. The change is the size of the ion, not the number of electrons.
• These are metallic structures. Any mention of molecules, intermolecular forces or covalent bonds scores zero on the whole explanation.

Reactions of Mg-Ba with water

Reactivity increases down the group, because the first ionisation energy decreases, so the atoms lose their two outer electrons more easily.

Magnesium with steam gives magnesium oxide and hydrogen. Magnesium burns with a bright white flame and a white solid (MgO) is left:

$Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$

Magnesium with warm water reacts very slowly to give magnesium hydroxide (a different product) and hydrogen, with no flame. The hydroxide is only sparingly soluble, so it forms as a solid coating, which is part of why the reaction is slow:

$Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(s) + H_2(g)$

Calcium, strontium and barium with cold water react with increasing vigour down the group, forming the hydroxide and hydrogen:

$Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$

Down the group you see more fizzing, the metal dissolving faster and the mixture getting hotter.

Magnesium in the extraction of titanium

Titanium is extracted from titanium(IV) chloride, TiCl₄, by reaction with a more reactive metal, magnesium:

$TiCl_4 + 2Mg \rightarrow Ti + 2MgCl_2$

• Magnesium is the reducing agent (it reduces the titanium, going from Ti in TiCl₄ to the metal).
• The reaction is carried out in an inert argon atmosphere, so that the hot, reactive magnesium and titanium do not react with oxygen (or nitrogen) in the air.

Solubility of the hydroxides

The hydroxides become more soluble down the group.

• Mg(OH)₂ is sparingly soluble (only slightly soluble).
• Solubility rises to Ba(OH)₂, which is soluble.

More of the hydroxide dissolves down the group, so more OH⁻ ions are released and the solution becomes more alkaline (higher pH) down the group. A sparingly soluble hydroxide forms a white precipitate, for example:

$Mg^{2+}(aq) + 2OH^{-}(aq) \rightarrow Mg(OH)_2(s)$

Solubility of the sulfates

The sulfates become less soluble down the group (the opposite trend to the hydroxides).

• MgSO₄ is soluble.
• BaSO₄ is insoluble.

So the insoluble end is the bottom of the group: adding sulfate ions to a solution of Ba²⁺ gives a white precipitate of barium sulfate:

$Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$

Uses that come from the solubility trends

The uses are just the trends applied, so learn them together:

• Mg(OH)₂ in medicine ("milk of magnesia") is an antacid: it neutralises excess hydrochloric acid in the stomach. $Mg(OH)_2 + 2HCl \rightarrow MgCl_2 + 2H_2O$ It is safe because it is only sparingly soluble, so the suspension is only weakly alkaline and not corrosive.
• Ca(OH)₂ in agriculture is spread on fields to neutralise acidic soils, raising the pH for better crop growth.
• CaO or CaCO₃ in flue-gas desulfurisation removes the acidic gas SO₂ from power-station waste gases. The basic solid neutralises the acidic gas: $CaO + SO_2 \rightarrow CaSO_3$ $CaCO_3 + SO_2 \rightarrow CaSO_3 + CO_2$
• BaSO₄ in medicine is the "barium meal" swallowed before an X-ray of the gut: barium is good at absorbing X-rays, so the gut shows up. Barium ions are toxic, but BaSO₄ is safe because it is insoluble, so it is not absorbed into the blood.

Test for sulfate ions

To test an unknown solution for sulfate ions (SO₄²⁻):

1. Add dilute hydrochloric acid to the solution first.
2. Then add acidified barium chloride solution, BaCl₂(aq).
3. A white precipitate (of barium sulfate) confirms a sulfate is present:

$Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$

Why the barium chloride is acidified: the acid reacts with and removes carbonate (and sulfite) ions, which would otherwise form their own white precipitate (BaCO₃ or BaSO₃) and give a false positive result.

Why the acid must not be sulfuric acid: sulfuric acid contains sulfate ions, so it would give a white precipitate itself. Use dilute hydrochloric acid (nitric acid also works).

Still don't get it? · why acidified barium chloride, and why not sulfuric acid

Imagine you are testing a crowd for one specific person wearing a red hat. The trouble is, a couple of other people are also wearing red hats and could fool you. So before you look, you quietly ask those lookalikes to leave.

The "red hat" here is a white precipitate with barium. Sulfate ions give one, but so do carbonate and sulfite ions. If any of those are hiding in the solution, you would see a white precipitate and wrongly call it a sulfate: a false positive.

Adding dilute acid first gets rid of the lookalikes: the acid reacts with the carbonate and sulfite ions and removes them, so only a genuine sulfate can give the white precipitate when you add the barium chloride.

Now the twist: you must not use sulfuric acid to do this, because sulfuric acid is itself full of sulfate ions, so it would make its own barium sulfate precipitate and ruin the test. Use dilute hydrochloric acid (or nitric). In the exam: acidify with dilute HCl to remove carbonate ions that would give a false positive, then add BaCl₂(aq); a white precipitate of BaSO₄ shows sulfate ions are present.

Worked examples

These two are the answers examiners most often see thrown away. The lesson in each is the number and order of the points, and the exact wording.

Model answer: "Explain why the first ionisation energy of barium is lower than that of magnesium." (3 marks)

Each mark is one comparative point, in a chain:

1. In barium the outer electron is in a shell further from the nucleus than in magnesium.
2. Barium has more inner shells, so more shielding of the outer electron from the nuclear charge.
3. These outweigh the increased nuclear charge, so the outer electron is less strongly attracted and needs less energy to remove.

Where the mark is lost: giving one reason without the other, or forgetting to make it comparative (you must compare barium with magnesium, not just describe barium). Never mention molecules.

Model answer: "Describe how you would test a solution to show it contains sulfate ions, and explain why the reagent is acidified." (4 marks)

1. Add dilute hydrochloric acid, then add barium chloride solution (BaCl₂(aq)).
2. A white precipitate forms if sulfate ions are present.
3. Ionic equation: $Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$
4. The acid is added to remove carbonate ions (and sulfite ions), which would otherwise give a white precipitate and a false positive.

Where the mark is lost: naming the reagent as "barium" or "Ba²⁺ ions" instead of a real reagent like barium chloride (no named reagent means the observation mark is lost too); and using sulfuric acid to acidify, which contains sulfate ions and gives a false positive.

Common exam mistakes

• Explaining the ionisation energy trend with just one reason. The marks need the outer electron further from the nucleus, more shielding, and that these outweigh the greater nuclear charge, stated comparatively.
• Explaining the melting point trend by saying there are "more" or "fewer" delocalised electrons down the group. Every Group 2 atom releases two; the number is the same. The change is the size of the ion and the distance to the delocalised electrons.
• Mentioning molecules, intermolecular forces or covalent bonds in a metallic-bonding or atomic-property answer. This is a contradiction error and scores zero.
• Confusing the two water reactions of magnesium: steam gives the oxide (MgO) with a bright white flame; warm water gives the hydroxide (Mg(OH)₂) very slowly. The products are different.
• Reversing the two solubility trends, or giving both the same direction. Hydroxides increase, sulfates decrease down the group. Also, do not describe the solubility of the metals; the question is about the compounds.
• Calling Mg(OH)₂ "insoluble". The required word is sparingly soluble (slightly soluble); that slight solubility is why it works as an antacid.
• Wrong formulae for Group 2 compounds, the single most common slip: it is MgCl₂ (not MgCl), Mg(OH)₂ (not MgOH), BaCl₂ (not BaCl), Ba²⁺ (not Ba⁺), and TiCl₄. A wrong formula then makes the balanced equation impossible.
• In the sulfate test, naming the reagent as "barium" or "Ba²⁺" rather than a specific reagent such as barium chloride; using sulfuric acid to acidify (it contains sulfate ions); or thinking the acid is the test rather than a step to remove interfering ions.
• Confusing the uses of the two medicinal hydroxides and compounds: Mg(OH)₂ is the antacid, Ca(OH)₂ is for soil, and BaSO₄ is the barium meal.
• Missing or wrong state symbols in the equations (especially the (s) on a precipitate and (aq) on a dissolved reactant); examiners penalise these.

Key definitions

• First ionisation energy - the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
• Shielding - the repulsion of outer electrons by inner-shell electrons, which reduces the effective attraction of the nucleus for the outer electrons.
• Metallic bonding - the electrostatic attraction between positive metal ions and the delocalised electrons.
• Sparingly soluble - only slightly soluble; a small amount dissolves and the rest stays as a solid.
• Reducing agent - an electron donor (the species that reduces another and is itself oxidised).
• Positive test for sulfate ions - a white precipitate (of barium sulfate) forms when acidified barium chloride solution is added.

Specification

• I can explain the trend in atomic radius down Group 2 (more shells, so larger atoms).
• I can explain the trend in first ionisation energy down Group 2 (further from the nucleus, more shielding, outweighing the greater nuclear charge), comparatively.
• I can explain the trend in melting point in terms of structure and bonding (giant metallic lattice; larger ions, so weaker attraction to the delocalised electrons), and know Mg is an exception.
• I can write equations for the reactions of Mg-Ba with water, and distinguish Mg + steam (MgO) from Mg + warm water (Mg(OH)₂).
• I can describe the use of magnesium to extract titanium from TiCl₄, with Mg as the reducing agent.
• I can state the trend in solubility of the hydroxides (increases; Mg(OH)₂ sparingly soluble).
• I can state the trend in solubility of the sulfates (decreases; BaSO₄ insoluble).
• I can give the uses of Mg(OH)₂ (medicine) and Ca(OH)₂ (agriculture), and of CaO/CaCO₃ to remove SO₂ from flue gases.
• I can describe and explain the acidified BaCl₂ test for sulfate ions (why acidified, and why not sulfuric acid), and give the use of BaSO₄ in medicine.

Related notes

• Periodicity
• Group 7 (the halogens)
• Atomic structure
• Bonding
• Properties of Period 3 elements and their oxides (A-level)