What this chapter covers10.1
Group 2 is the topic where trends stop being an abstraction and start doing work. Four metals, one column of the Periodic Table, and every difference between them — how fast they react with water, how hot you must make their carbonates before they crack, whether their sulfate dissolves — traces back to a single change: the ion gets bigger while its charge stays at 2+. Once you can say that sentence and follow it through, most of this chapter can be reconstructed rather than remembered.
The syllabus is unusually specific about which elements it means. The sub-topic is headed the Group 2 metals, magnesium to barium, and everything below keeps to those four. Beryllium appears only where it is useful to see something that does not fit, and it is labelled as such wherever it does.
What topic 10 asks you to do
10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds
10.1.1 describe, and write equations for, the reactions of the
elements with oxygen, water and dilute hydrochloric and sulfuric acids
10.1.2 describe, and write equations for, the reactions of the oxides, hydroxides and
carbonates with water and dilute hydrochloric and sulfuric acids
10.1.3 describe, and write equations for, the thermal decomposition of the nitrates and
carbonates, to include the trend in thermal stabilities
10.1.4 describe, and make predictions from, the trends in physical and chemical properties
of the elements involved in the reactions in 10.1.1 and the compounds involved in 10.1.2, 10.1.3
and 10.1.5
10.1.5 state the variation in the solubilities of the hydroxides and sulfates
Read those verbs carefully, because they are not the same from one statement to the next. 10.1.1, 10.1.2 and 10.1.3 say describe and write equations — they want the facts and the chemistry, not a mechanism. 10.1.5 says only state the variation: two sentences about solubility, with no explanation expected. 10.1.4 is the one that asks you to think, and it is where most of the marks in this topic actually live.
Where the explanations belong
Two explanations that feel as though they belong here are examined elsewhere. The reasoning behind the ionisation-energy trend is built in topic 1.3, and this page uses it rather than proves it. The full charge-density treatment of thermal stability — polarising power set out properly — is an A Level statement examined later in the course; 10.1.3 asks for the trend and the equations. The argument is given here anyway, because it is short and it makes the trend memorable, but do not feel short-changed if an AS question only wants "stability increases down the group".
The four metals10.1.4
Before any trend means anything, here is what changes from one element to the next — and, just as importantly, what does not.
| Element | Proton number | Electronic configuration | Outer electrons | Ion formed | Flame colour |
|---|---|---|---|---|---|
| Be beryllium | 4 | 1s22s2 | 2 | Be2+ (atypical) | none |
| Mg magnesium | 12 | 1s22s22p63s2 | 2 | Mg2+ | brilliant white |
| Ca calcium | 20 | [Ar]4s2 | 2 | Ca2+ | brick red |
| Sr strontium | 38 | [Kr]5s2 | 2 | Sr2+ | crimson red |
| Ba barium | 56 | [Xe]6s2 | 2 | Ba2+ | apple green |
The one sentence this whole chapter rests on
Every one of these elements has two outer s electrons and forms a 2+ ion by losing both. So the charge on the ion never changes. The only thing that changes down the group is size — and every trend on this page, physical or chemical, is a consequence of size changing while charge stands still. When an exam question asks you to explain a group 2 trend, that is the argument it is looking for.
The model below plots any of the measured properties of these elements. It also measures the shape of what it has plotted, which matters more than it sounds: one of these trends is not smooth, and a page that simply asserted "decreases down the group" would be wrong about it.
Why beryllium is kept at arm's length
Be2+ has a radius of roughly 45 pm and carries two units of charge, which gives it an extraordinary pull on any nearby anion — enough to drag the electron cloud so far towards itself that the bonding is no longer properly ionic. Beryllium chloride is a covalent, chain-forming solid; beryllium hydroxide is amphoteric, not simply basic. None of the statements in 10.1 would survive being applied to it, which is exactly why the syllabus wrote "magnesium to barium".
Atomic and ionic radius10.1.4
Atomic radius rises from 160 pm at magnesium to 217 pm at barium, and ionic radius from 72 pm to 135 pm. Both rise at every step, which makes this the easiest trend in the topic to state and the one most often stated for the wrong reason.
The explanation, in the form that earns the marks
Going down group 2, each element has one more occupied electron shell than the one above. The outer electrons are therefore further from the nucleus and are shielded from it by more complete inner shells. Both effects weaken the attraction between the nucleus and the outer electrons, so the atom is larger — even though the nuclear charge has increased.
Trap: "the nuclear charge increases, so the atom gets smaller"
Wrong, and it is the commonest wrong answer in the topic. The nuclear charge does increase — from +12 to +56 — and across a period that argument would be right. Down a group it is beaten, because the extra protons arrive alongside a full extra shell of inner electrons that cancel almost all of them. Never use a period argument on a group trend.
The ion is always much smaller than the atom
Magnesium goes from 160 pm to 72 pm on becoming Mg2+; barium from 217 pm to 135 pm. The reason is the same in every case, and it is worth having ready as a sentence:
Losing both outer electrons removes an entire occupied shell, so the outermost electrons of the ion are one shell closer in. At the same time the unchanged nuclear charge now has two fewer electrons to hold, so it holds each of the remaining ones more tightly. Both effects shrink it.
Why the ionic radii matter more than the atomic ones
Almost every compound in this chapter contains M2+, not M. So when you come to explain thermal stability, or the solubility of a sulfate, the number that does the work is the ionic radius — and because the charge is fixed at 2+ across all four, ionic radius alone fixes the charge density. Atomic radius is what you quote when the question is about the metal itself: its melting point, its reactivity, its ionisation energy.
Ionisation energy10.1.4
First ionisation energy falls steadily from 738 kJ mol−1 at magnesium to 503 kJ mol−1 at barium. The second ionisation energy falls in the same way, from 1451 to 965 kJ mol−1. This is the trend that explains chemical reactivity, so it is worth more than its two lines in the syllabus suggest.
First ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Mg(g) → Mg+(g) + e−
The explanation is the same one as for atomic radius, which is the point — a group trend has one argument behind it and several symptoms. The electron being removed is further from the nucleus and shielded by more complete inner shells, so less energy is needed to take it away.
The number that makes the argument watertight is the effective nuclear charge. The model counts it for each element rather than asserting it.
Trap: "barium has 56 protons, so its electrons are held more tightly"
Wrong. Barium does have 56 protons — and 54 inner-shell electrons cancelling them. The outer pair feels a net pull of about +2, exactly as magnesium's does. Nuclear charge quoted without shielding predicts the opposite of what actually happens, which is why an answer built on it scores nothing even though the fact inside it is true.
Putting the four in order by ionisation energy is a useful check, because the order is the reverse of the reactivity order and it is easy to state one while meaning the other.
Two ionisation energies, not one
Forming M2+ costs the sum of the first two ionisation energies: 2189 kJ mol−1 for magnesium, 1468 for barium. Both numbers look enormous, and on their own they suggest these ions should never form. What pays for them is the lattice energy of the solid, or the hydration energy of the ion in solution — and a 2+ ion attracts far more strongly than a 1+ one would. That is why group 2 elements form 2+ ions rather than stopping at 1+ where it would be cheaper.
Melting points10.1.4
This is the one trend in the topic that is not clean, and being honest about it is worth more than smoothing it over.
| Metal | Melting point / K | Melting point / °C | Ionic radius / pm | Crystal structure |
|---|---|---|---|---|
| Mg | 923 | 650 | 72 | hexagonal close-packed |
| Ca | 1115 | 842 | 100 | face-centred cubic |
| Sr | 1050 | 777 | 118 | face-centred cubic |
| Ba | 1000 | 727 | 135 | body-centred cubic |
From calcium to barium the melting point falls, and the explanation is the standard one for metallic bonding:
Why metallic bonding weakens down the group
Each atom releases the same two electrons into the delocalised sea, so the charge density of the electron sea does not change. What does change is the size of the cation: as it grows, the distance between the nucleus and the delocalised electrons increases, so the electrostatic attraction between the cations and the sea of delocalised electrons is weaker and less energy is needed to melt the metal.
Magnesium is out of line — say so
Magnesium melts 192 K below calcium, which the argument above cannot explain: magnesium has the smallest cation of the four and ought to melt highest. The reason is that its crystal packs differently from calcium's and strontium's, and melting point depends on how the lattice is arranged as well as on how strongly it is held. The safe exam statement is "the melting points generally decrease down the group, with magnesium as an exception" — or simply describe the fall from calcium to barium, which is uncontroversial.
Trap: explaining a metal's melting point with ionic charge
"The charge on the ion decreases down the group, so the bonding is weaker" — wrong, because the charge is 2+ on every one of them. The charge argument belongs to a comparison across a period, where Na+, Mg2+ and Al3+ really do differ. Down group 2 the only variable is size.
Checking the physical trends10.1.4
Six claims, each with a gap in it. Every blank below is marked against the numbers in the data table rather than against a stored answer, so getting one right means you have agreed with the measurements and not merely with a form of words.
Reaction with oxygen10.1.1
All four metals burn in oxygen to give the oxide, and all four do it more readily down the group. The equation is the same shape every time, because the metal always goes to +2 and the oxygen always to −2.
2Mg(s) + O2(g) → 2MgO(s)
2Ca(s) + O2(g) → 2CaO(s)
2Ba(s) + O2(g) → 2BaO(s)
The redox, in oxidation numbers
The metal goes from 0 to +2 — it loses two electrons and is oxidised. Oxygen goes from 0 to −2 — it gains two electrons and is reduced. The metal is the reducing agent; the oxygen is the oxidising agent. Every reaction in 10.1.1 is a redox reaction of exactly this kind, with the metal always the one being oxidised.
Strontium and barium also make peroxides
In excess oxygen, strontium and barium go one stage further and form SrO2 and BaO2. These contain the peroxide ion, O22− — one negative charge on each oxygen atom, not two on a single one. The bigger cations at the bottom of the group are better able to stabilise the bigger anion. A question asking for "the reaction with oxygen" wants the simple oxide unless it says excess.
The model below assembles any of the reactions in 10.1.1 from the formulae and checks that it balances before it prints it.
Flame colours10.1.1
Group 2 compounds colour a flame, and the colour identifies the metal. This is the one place in the topic where a physical observation is enough to name an element.
| Ion | Flame colour | What to write |
|---|---|---|
| Mg2+ | brilliant white | white — often described as giving no colour to the flame |
| Ca2+ | brick red | brick red, or orange-red |
| Sr2+ | crimson red | crimson, or scarlet |
| Ba2+ | apple green | green, or pale green |
Why a flame test works
The heat of the flame promotes an electron to a higher energy level. When that electron drops back down, the energy it releases is emitted as light of a definite frequency, fixed by the size of the gap between the two levels. Because the energy levels are characteristic of the element, so is the colour.
Trap: "the metal burns with a coloured flame"
Wrong. A flame test uses a compound — a chloride on a clean wire — and nothing is burning. The colour is emission from excited electrons falling back, not combustion. Say excited, higher energy level and emits, and the mark is safe.
Reaction with water10.1.1
Every one of these metals reduces hydrogen out of water, and the speed at which it does so is the clearest demonstration in the topic that reactivity increases down the group.
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
Sr(s) + 2H2O(l) → Sr(OH)2(aq) + H2(g)
Ba(s) + 2H2O(l) → Ba(OH)2(aq) + H2(g)
| Metal | With cold water | pH of the solution |
|---|---|---|
| Mg | extremely slow — a few bubbles after days | 9–10 |
| Ca | steady effervescence; the mixture warms and goes cloudy | 12–13 |
| Sr | faster still | 13 |
| Ba | vigorous | 13–14 |
Magnesium is the exception, twice over
Magnesium barely reacts with cold water, and there are two reasons rather than one. It is the least reactive of the four — and the magnesium hydroxide it produces is almost insoluble, so it forms a coating that keeps the water away from the metal underneath. With steam, though, magnesium burns readily, and the product is the oxide rather than the hydroxide:
Mg(s) + H2O(g) → MgO(s) + H2(g)
Cold water gives the hydroxide; steam gives the oxide
Read the question for the word steam. It changes both the product and the equation's coefficients — 2H2O for the hydroxide route, one H2O for the steam route. Writing Ca + H2O → CaO + H2 for calcium in cold water is a common and avoidable error.
Why the solution is alkaline, and why the pH differs
The alkalinity comes from the hydroxide produced, so how alkaline the solution ends up depends on how much of that hydroxide will dissolve. Magnesium hydroxide barely dissolves, so the pH stops near 10. Barium hydroxide dissolves freely, so the pH goes past 13. That link — from solubility to pH — is the one this topic keeps returning to, and it is set out properly in the solubility sections.
Trap: writing the hydroxide as MOH
Group 1 gives NaOH. Group 2 gives Ca(OH)2, with brackets and a subscript 2, because the ion is 2+ and needs two hydroxide ions to balance it. "CaOH" is not a real substance, and an equation containing it cannot balance.
Reaction with dilute acids10.1.1
With dilute hydrochloric acid, every one of these metals behaves the way a reactive metal should: effervescence, hydrogen given off, the metal dissolving to leave a colourless solution of the chloride.
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
Ca(s) + 2HCl(aq) → CaCl2(aq) + H2(g)
With dilute sulfuric acid the equation looks just as tidy —
Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g)
Ca(s) + H2SO4(aq) → CaSO4(s) + H2(g)
— and for calcium, strontium and barium it does not describe what you would see, because the sulfate produced is insoluble.
The trap this whole section exists for
Put calcium in dilute sulfuric acid and it fizzes for a moment and then stops. The calcium sulfate formed is only sparingly soluble, so it builds up as a layer on the surface of the metal and keeps the acid away from it. The reaction has not finished — it has been sealed off. Put the same calcium in hydrochloric acid and it dissolves completely, because calcium chloride is soluble.
The effect gets worse down the group as the sulfates get less soluble: barium in dilute sulfuric acid does almost nothing at all. Magnesium is the exception — magnesium sulfate is soluble, so magnesium reacts freely with both acids.
Keep to the two acids named
10.1.1 says dilute hydrochloric and sulfuric acids, and that is deliberate. Nitric acid is an oxidising acid and gives nitrogen oxides rather than hydrogen; concentrated sulfuric acid behaves differently again. Neither is being asked about here, so do not volunteer them.
Why reactivity increases down the group10.1.4
Three separate observations — burning in oxygen, reacting with water, dissolving in acid — all give the same order: Mg < Ca < Sr < Ba. One explanation covers all three, and it is the explanation the mark scheme wants.
The reactivity argument, in full
In every one of these reactions the metal is oxidised: it loses its two outer electrons to form M2+. So "more reactive" means "gives up those two electrons more readily".
Going down the group, the outer electrons are in a shell further from the nucleus and are shielded by more complete inner shells, so the ionisation energies fall. Less energy is needed to form the 2+ ion, so the reaction happens more readily.
| Metal | 1st I.E. | 2nd I.E. | Sum / kJ mol−1 | E°(M2+/M) / V |
|---|---|---|---|---|
| Mg | 738 | 1451 | 2189 | −2.37 |
| Ca | 590 | 1145 | 1735 | −2.87 |
| Sr | 549 | 1064 | 1613 | −2.89 |
| Ba | 503 | 965 | 1468 | −2.91 |
The electrode potential is the same statement in a different currency. A more negative E°(M2+/M) means the metal parts with its electrons more readily, and the values become more negative at every step down the group — exactly matching the ionisation-energy column beside them.
Two columns, one story — but they are not the same measurement
Ionisation energy is about gaseous atoms in isolation. Electrode potential is about a metal in contact with a solution, so it also includes the energy released when the ion is hydrated, and the energy needed to break up the metal lattice. That the two agree so neatly is worth noticing rather than assuming: the ionisation energies fall fast enough down the group to outweigh the fact that a larger ion is hydrated less strongly.
Trap: "barium is more reactive because it has more electrons to lose"
Wrong. All four lose exactly two. Nothing about the number of electrons changes down the group — only how tightly those two are held. An answer that does not mention distance and shielding has not explained anything.
Reactivity and ionisation energy run opposite ways
The most reactive metal has the lowest ionisation energy. It is easy, under exam pressure, to put the four "in order of reactivity" and hand in the ionisation-energy order by mistake. Write down which quantity you are ordering before you order it.
The oxides and hydroxides with water10.1.2
Every group 2 oxide reacts with water to give the hydroxide. The equation is the same for all four; what differs is how vigorous the reaction is and how alkaline the result.
MgO(s) + H2O(l) → Mg(OH)2(s)
CaO(s) + H2O(l) → Ca(OH)2(s)
BaO(s) + H2O(l) → Ba(OH)2(aq)
| Oxide | What happens with water | pH of the solution |
|---|---|---|
| MgO | very little visible reaction; the hydroxide is nearly insoluble | ≈ 9–10 |
| CaO | reacts vigorously and exothermically — the solid crumbles and steams | ≈ 12–13 |
| SrO | as calcium, but faster; more of the hydroxide dissolves | ≈ 13 |
| BaO | reacts readily; the hydroxide is the most soluble of the four | ≈ 13.5 |
Two names worth knowing
Quicklime is calcium oxide, CaO. Slaked lime is calcium hydroxide, Ca(OH)2, and the reaction of quicklime with water is called slaking — it is violently exothermic. Limewater is a saturated solution of calcium hydroxide, and it is only mildly alkaline because so little of it will dissolve.
Strong base, weak solution
All four hydroxides are strong bases: every formula unit that dissolves is fully dissociated into M2+ and 2OH−. Magnesium hydroxide is not a weak base — it is a strong base that is almost insoluble. The pH of a saturated solution is set by how much dissolves, not by how far it ionises, and confusing the two is the single most common error in this part of the topic.
The hydroxides themselves do not react with water at all. They can only dissolve in it, and the extent to which they do is the trend in 10.1.5.
The oxides and hydroxides with acids10.1.2
Oxides and hydroxides are both bases, so both neutralise an acid to give a salt and water.
MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
Ca(OH)2(s) + 2HCl(aq) → CaCl2(aq) + 2H2O(l)
MgO(s) + H2SO4(aq) → MgSO4(aq) + H2O(l)
Ba(OH)2(aq) + H2SO4(aq) → BaSO4(s) + 2H2O(l)
The sulfate problem again
Exactly the same thing happens here as with the metals themselves. With hydrochloric acid every reaction goes to completion, because all four chlorides are soluble. With sulfuric acid the reactions of the calcium, strontium and barium compounds slow down and stop, because the sulfate produced is insoluble and coats the solid. Only magnesium reacts freely with both.
If a question asks you to prepare a group 2 salt, this decides your method: soluble sulfates can be made by neutralisation, insoluble ones have to be made by precipitation.
The model builds any of the reactions in 10.1.2 and works out whether it will actually run to completion, from the solubility of the salt it would produce.
The carbonates with water and acids10.1.2
The carbonates are the odd ones out, in one specific way: they are all essentially insoluble in water, and none of them reacts with it. That is why calcium carbonate exists as chalk, limestone and marble rather than having dissolved away.
With an acid, though, they react readily, and the reaction has a signature that no other compound in this chapter shares:
MgCO3(s) + 2HCl(aq) → MgCl2(aq) + H2O(l) + CO2(g)
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
MgCO3(s) + H2SO4(aq) → MgSO4(aq) + H2O(l) + CO2(g)
Why acid dissolves what water cannot
Water fails because the energy released by hydrating the ions does not pay for pulling the lattice apart. Acid does not have to pay that bill. It removes the carbonate ion from the equation altogether — CO32− + 2H+ → H2O + CO2 — and the carbon dioxide leaves the mixture as a gas. With one of the products escaping, the solid keeps dissolving until it is gone.
Carbon dioxide, not hydrogen
A metal with an acid gives hydrogen. A carbonate with an acid gives carbon dioxide. Both fizz, and telling them apart is the point of the gas test: carbon dioxide turns limewater milky, hydrogen gives a squeaky pop with a lit splint.
Trap: carbonates plus sulfuric acid
Only magnesium carbonate reacts properly with dilute sulfuric acid. Calcium, strontium and barium carbonates fizz briefly and then stop, coated in their own insoluble sulfate — the same effect as everywhere else in this chapter. Use hydrochloric acid if you want the reaction to finish.
Heating the carbonates10.1.3
Heat a group 2 carbonate strongly enough and it falls apart into the oxide and carbon dioxide. The equation is the same for every one of them:
MgCO3(s) → MgO(s) + CO2(g)
CaCO3(s) → CaO(s) + CO2(g)
BaCO3(s) → BaO(s) + CO2(g)
What is not the same is how hot you have to make it, and that is the trend 10.1.3 asks for.
| Carbonate | Ionic radius of M2+ / pm | Decomposes near / °C | Thermal stability |
|---|---|---|---|
| MgCO3 | 72 | 540 | least stable |
| CaCO3 | 100 | 900 | |
| SrCO3 | 118 | 1290 | |
| BaCO3 | 135 | 1360 | most stable |
These are the temperatures normally quoted at this level. A carbonate does not have a sharp decomposition point the way a solid has a melting point, so treat them as indicative — it is the order and the size of the gaps that matter.
Thermal stability increases down group 2
The carbonates become more thermally stable down the group, so a higher temperature is needed to decompose them. Magnesium carbonate decomposes below 600 °C; barium carbonate needs more than twice that.
Nothing is oxidised or reduced
Check the oxidation numbers before calling this redox: the metal stays at +2, carbon stays at +4, oxygen stays at −2. The carbonate ion is simply splitting into an oxide ion and a carbon dioxide molecule. Compare it with the nitrates in the next section, where something genuinely is reduced.
Heating the nitrates10.1.3
Group 2 nitrates decompose to the oxide, and they give off two gases rather than one. This equation has to be learnt as a shape, because getting the coefficients right is a mark in itself:
2Mg(NO3)2(s) → 2MgO(s) + 4NO2(g) + O2(g)
2Ca(NO3)2(s) → 2CaO(s) + 4NO2(g) + O2(g)
2Ba(NO3)2(s) → 2BaO(s) + 4NO2(g) + O2(g)
How to see that it balances
Take two formula units, because one would leave you half an oxygen molecule. Two units contain 4 nitrogen and 12 oxygen. The two oxides take 2 of those oxygens; the 4 nitrogens leave as 4 NO2, taking 8 more. That leaves 2 oxygen atoms — one O2 molecule. Every group 2 nitrate follows this pattern exactly, so once you can rebuild it you never have to recall it.
What you would observe
The white solid melts, then gives off a brown gas — nitrogen dioxide — and a gas that relights a glowing splint — oxygen. A white solid is left behind: the oxide. The same trend applies as for the carbonates: the nitrates become more stable down the group and need stronger heating.
Group 1 nitrates stop halfway
Sodium and potassium nitrates decompose only as far as the nitrite — 2NaNO3 → 2NaNO2 + O2 — releasing oxygen but no brown gas. Their 1+ ions are far too weakly polarising to break the nitrate ion all the way down. Lithium, whose ion is small and 1+, behaves like a group 2 metal and goes to the oxide. That comparison is the best evidence there is that charge density, and not something about the anion, is what is doing the work.
Why stability increases down the group10.1.3
Two compounds, one explanation, and it rests on the sentence this chapter opened with: the charge stays at 2+ while the ion gets bigger.
The polarising-power argument
A small cation with a high charge has a high charge density, and therefore a strong polarising power — it pulls the electron cloud of a neighbouring anion towards itself and distorts it. Distorting the carbonate ion weakens the C—O bonds on the side facing the cation, until the ion breaks into O2− and CO2.
Down group 2 the cation gets larger while the charge stays at 2+, so the charge density and the polarising power fall. The carbonate ion is distorted less, so more heat is needed before it will decompose. Thermal stability therefore increases down the group.
Because volume goes as the cube of the radius, the fall in charge density down the group is far steeper than the radii themselves suggest. The model computes it from the ionic radii and then asks whether it really does track the decomposition temperatures.
And here is the same idea as a picture — the distortion drawn is scaled to the charge density the model has just calculated for whichever ion you pick.
The same argument explains three other things
Polarising power is one of the most portable ideas in inorganic chemistry. It explains why lithium behaves unlike the rest of group 1, why beryllium compounds are covalent enough to fall outside this topic altogether, and why a small highly charged cation such as Al3+ makes its solutions acidic. Whenever you meet a group 1 or group 2 anomaly, ask what the charge density of the cation is before looking for anything more exotic.
How much of this an AS question wants
Statement 10.1.3 asks you to describe the trend and write the equations. The polarising-power explanation is examined in full later in the course, at A Level. It is given here because it makes the trend impossible to forget and costs two sentences — but if an AS question gives one mark for the trend, give the trend.
Solubility of the hydroxides10.1.5
The hydroxides become more soluble down the group. That is the whole of what 10.1.5 asks for on this half of the topic — the statement says state, not explain.
| Hydroxide | Solubility / mol dm−3 | Description | pH of saturated solution |
|---|---|---|---|
| Mg(OH)2 | 1.2 × 10−4 | very sparingly soluble | ≈ 10.4 |
| Ca(OH)2 | 1.6 × 10−2 | moderately soluble | ≈ 12.5 |
| Sr(OH)2 | 3.4 × 10−2 | moderately soluble | ≈ 12.8 |
| Ba(OH)2 | 1.5 × 10−1 | soluble | ≈ 13.5 |
Solubilities at 298 K. Values in the literature vary with temperature and with the hydration state of the solid, so use these for the trend and for comparison rather than as data-booklet quantities.
Solubility and pH are the same fact twice
Each formula unit that dissolves releases two hydroxide ions, so [OH−] = 2 × the solubility, and the pH follows straight from that. A hundredfold rise in solubility from magnesium to calcium is a rise of about two pH units — which is exactly the gap in the table. The model below does the arithmetic step by step.
Solubility of the sulfates10.1.5
The sulfates run the other way: they become less soluble down the group. Magnesium sulfate dissolves freely; barium sulfate is, for every practical purpose, insoluble.
| Sulfate | Solubility / mol dm−3 | Description | Where it shows up |
|---|---|---|---|
| MgSO4 | 1.8 | soluble | Epsom salts; magnesium reacts freely with sulfuric acid |
| CaSO4 | 4.7 × 10−3 | sparingly soluble | gypsum, plaster; the coating that stops calcium in sulfuric acid |
| SrSO4 | 7.1 × 10−4 | very sparingly soluble | celestine, the main ore of strontium |
| BaSO4 | 9.4 × 10−6 | insoluble for practical purposes | barite; the sulfate test; the barium meal |
Two trends, opposite directions — anchor them on barium
These are the two easiest facts in the topic to swap over. Fix them with one element: barium hydroxide is the soluble one; barium sulfate is the insoluble one. Everything else follows, because each trend is monotonic. If you can also remember why — barium sulfate is what makes the sulfate test work, and barium hydroxide is sold as a standard alkali — the pair becomes hard to get the wrong way round.
Why these two trends oppose each other
Whether an ionic solid dissolves is a contest between the lattice energy holding it together and the hydration energy released when its ions are surrounded by water — and both weaken as the cation gets larger. Which one weakens faster decides the trend, and that depends on the size of the anion. With a small anion such as OH−, lattice energy is very sensitive to cation size and falls away faster, so solubility rises. With a large anion such as SO42−, the lattice energy hardly notices the cation and the falling hydration energy wins, so solubility drops. 10.1.5 does not ask for this — the statement says only "state the variation" — but it explains why one trend is not simply the other written backwards.
Testing for sulfate ions10.1.5
Barium sulfate's insolubility is not just a fact to memorise — it is the basis of the standard test for a sulfate ion in solution.
The test, in the order an examiner wants it
1. Acidify the unknown solution with dilute hydrochloric acid.
2. Add barium chloride solution.
3. A white precipitate confirms a sulfate.
Ba2+(aq) + SO42−(aq) → BaSO4(s)
Trap: leaving out the acid
The acid is not a formality. Barium carbonate and barium sulfite are also white insoluble solids, so a carbonate or a sulfite in the unknown would give a white precipitate and be read as a sulfate. The acid destroys both before the barium is added — CO32− + 2H+ → H2O + CO2 and SO32− + 2H+ → H2O + SO2 — while barium sulfate is unaffected by acid. That difference is what makes the test specific.
Try each of those ions, with the acid and without it.
Where these compounds are used10.1.2, 10.1.5
Each of these uses is a property of the compound doing something useful, and in every case the property is one already on this page.
| Compound | Use | The property that makes it work |
|---|---|---|
| Mg(OH)2 | indigestion remedy (milk of magnesia) | a base, so it neutralises stomach acid — but so insoluble that the suspension itself is only mildly alkaline and cannot burn the stomach lining |
| Ca(OH)2 | liming fields to raise the pH of acidic soil | moderately soluble and alkaline enough to neutralise soil acidity without the pH overshooting |
| CaCO3 | neutralising acidity, and making cement and quicklime | a cheap insoluble base that reacts with acid; decomposes to CaO on heating |
| BaSO4 | the "barium meal" — X-ray imaging of the gut | opaque to X-rays, and so insoluble that essentially no toxic Ba2+ reaches the blood |
| acidified BaCl2 | the test for sulfate ions | barium sulfate is insoluble in water and in acid |
Why a barium meal is safe when barium is poisonous
It is the free Ba2+ ion that is toxic. Barium sulfate releases so few of them — about 10−5 mol dm−3 — and is unaffected by stomach acid, so it passes straight through. Barium carbonate would be lethal in the same situation, because stomach acid dissolves it and sets the ion free. Two compounds of the same metal, one safe to drink and one not, separated entirely by the solubility trend on this page.
Checking the compound trends10.1.2, 10.1.3, 10.1.5
The same exercise as before, for the compounds. Every blank is marked against the solubility table, the calculated pHs and the decomposition temperatures used above.
Predicting from the trends10.1.4
10.1.4 is the statement that asks you to do something with everything above: describe the trends, and make predictions from them. That second half is what separates a recall question from a thinking one, and it is where the marks are.
Before predicting anything, it is worth being able to write down the reactions themselves without hesitating. Every product on this page comes from one of five patterns.
The five patterns, in one place
metal + oxygen → the oxide · 2M + O2 → 2MO
metal + water → the hydroxide + hydrogen · M + 2H2O → M(OH)2 + H2
oxide + water → the hydroxide · MO + H2O → M(OH)2
carbonate, heated → the oxide + carbon dioxide · MCO3 → MO + CO2
metal + chlorine → the chloride · M + Cl2 → MCl2
How to answer a prediction question
Radium sits below barium, and nothing about it appears in this syllabus — which makes it the perfect test, because the only way to produce an answer is to extend a trend. The model fits a line through the four elements you do know and extends it one period further — and then, to show you what that is worth, it fits the same line to magnesium, calcium and strontium alone and checks it against the barium you can measure.
A number on its own rarely earns the mark
"Radium's first ionisation energy would be a little below barium's" is half an answer. The other half is the reason: because the outer electrons are in a shell further from the nucleus and shielded by more inner shells, so less energy is needed to remove one. State the trend, give the reason, then give the value — and give it as a rough figure or a direction, not a confident number.
A checklist for predicting anything about radium
Larger atom and larger ion · lower ionisation energies · lower electronegativity · more reactive than barium · reacts vigorously with cold water · its hydroxide is the most soluble and most alkaline · its sulfate is the least soluble · its carbonate and nitrate are the most thermally stable, needing the highest temperatures · it forms a 2+ ion like the rest.
Every item on that list is one of the trends above, extended by one step, and that is what an examiner is asking for. Notice how many of them run in opposite directions — which is why "more reactive, so everything is more extreme" is not a safe substitute for knowing which way each trend goes.
Where the real radium refuses to cooperate
Two items on that checklist are, in fact, wrong — and it is worth knowing which. Radium's measured first ionisation energy is 509 kJ mol−1, slightly higher than barium's 503, and its electronegativity is quoted at 0.90 against barium's 0.89. Both are the wrong side of the trend. The cause is relativistic: in an atom this heavy the innermost electrons move fast enough that the 7s orbital is pulled in and held more tightly than the group argument allows for.
This is not examinable, and the extrapolated answer is still the one that earns the mark. It is here because it is the cleanest demonstration of what a trend actually is: a summary of the elements you have measured, and a reasonable expectation about the ones you have not — never a guarantee. The model above will show you the gap for each property if you ask it to.
Trap: predicting from beryllium
A question sometimes asks you to predict a property of beryllium instead, and the trends above will mislead you. Be2+ is so small and so strongly polarising that its compounds have marked covalent character — beryllium hydroxide is amphoteric rather than simply basic, and beryllium does not react with water at all. If a question asks about beryllium, the answer it usually wants is why the trend fails, not the number the trend predicts.
Data10.1
Everything plotted or calculated on this page comes from the tables below. Beryllium is included so that the misfit is visible, and is marked wherever it is excluded from a trend.
The elements
| Element | Z | Atomic radius / pm | Ionic radius M2+ / pm | 1st I.E. / kJ mol−1 | 2nd I.E. / kJ mol−1 | m.p. / K | b.p. / K | χ | E° / V |
|---|---|---|---|---|---|---|---|---|---|
| Be (outside 10.1) | 4 | 112 | 45 | 900 | 1757 | 1560 | 2744 | 1.57 | −1.85 |
| Mg | 12 | 160 | 72 | 738 | 1451 | 923 | 1363 | 1.31 | −2.37 |
| Ca | 20 | 197 | 100 | 590 | 1145 | 1115 | 1757 | 1.00 | −2.87 |
| Sr | 38 | 215 | 118 | 549 | 1064 | 1050 | 1655 | 0.95 | −2.89 |
| Ba | 56 | 217 | 135 | 503 | 965 | 1000 | 2170 | 0.89 | −2.91 |
Atomic radii are metallic radii; ionic radii are the six-coordinate values. χ is the Pauling electronegativity. E° is the standard electrode potential for M2+(aq) + 2e− → M(s).
The compounds
| Metal | M(OH)2 solubility / mol dm−3 | pH of saturated solution | MSO4 solubility / mol dm−3 | MCO3 decomposes near / °C |
|---|---|---|---|---|
| Mg | 1.2 × 10−4 | 10.4 | 1.8 | 540 |
| Ca | 1.6 × 10−2 | 12.5 | 4.7 × 10−3 | 900 |
| Sr | 3.4 × 10−2 | 12.8 | 7.1 × 10−4 | 1290 |
| Ba | 1.5 × 10−1 | 13.5 | 9.4 × 10−6 | 1360 |
These are reference values, not the official data booklet
Solubilities vary with temperature and with the hydration state of the solid, and published figures differ between sources; the decomposition temperatures are indicative, because a carbonate decomposes over a range rather than at a point. Every number here is good enough to establish the trends and to compare one element with another, which is what this topic asks for — but in an exam, quote whatever data the question or the official booklet gives you, not these.
The pH column is calculated from the solubility beside it, not measured: [OH−] = 2s, pOH = −log10[OH−], pH = 14 − pOH. A measured value would sit a little lower, because these solutions are concentrated enough that activity effects matter.
Equations worth being able to write without thinking
| Reaction | Equation |
|---|---|
| metal + oxygen | 2Ca + O2 → 2CaO |
| metal + water | Ca + 2H2O → Ca(OH)2 + H2 |
| magnesium + steam | Mg + H2O → MgO + H2 |
| metal + hydrochloric acid | Ca + 2HCl → CaCl2 + H2 |
| metal + sulfuric acid | Mg + H2SO4 → MgSO4 + H2 |
| oxide + water | CaO + H2O → Ca(OH)2 |
| oxide + acid | MgO + 2HCl → MgCl2 + H2O |
| hydroxide + acid | Ca(OH)2 + 2HCl → CaCl2 + 2H2O |
| carbonate + acid | CaCO3 + 2HCl → CaCl2 + H2O + CO2 |
| carbonate, heated | CaCO3 → CaO + CO2 |
| nitrate, heated | 2Ca(NO3)2 → 2CaO + 4NO2 + O2 |
| the sulfate test | Ba2+(aq) + SO42−(aq) → BaSO4(s) |
Self-test10.1
Thirty-three questions across the whole of topic 10. Every one has an explanation attached, so a wrong answer is worth more than a right one.