alkali earth metals

7.5. The reaction of s-block metals and oxygen & their oxide (O^2-) chemistry

The oxides and hydroxides are white ionic solids.

• The reaction of Group 1 metals with oxygen (a redox reaction)
• Group 1 metals: 4M_(s) + O_2(g) ==> 2M₂O_(s)    (M = Li, Na, K, Rb, Cs)
  • shows the formation of the 'simple' oxide expected from their position in the periodic table when the element is heated or burned in air.
  • Oxidation state changes: M is 0 to +1, Oxygen is 0 to -2 in the oxide ion O^2-.
    • ionically: 4M_(s) + O_2(g) ==> 2(M⁺)₂O^2-_(s)
      • the metal is oxidised (0 to +1), electron loss, increase in oxidation state
      • oxygen molecules are reduced (0 to -2), electron gain, decrease in oxidation state
  • The oxides are soluble in water forming the strongly alkaline hydroxide:
    • M₂O_(s) + H₂O_(l) ==> 2MOH_(aq)
    • ionically: (M⁺)₂O^2-_(s) + H₂O_(l) ==> 2M⁺_(aq)  +  2OH^-_(aq) (not a redox change)
      • This is an acid-base reaction, the O^2- ion is a strong Bronsted-Lowry base and accepts a proton from water (acting as the Bronsted-Lowry acid).
  • Unfortunately, except for lithium (an anomaly), 'higher' oxides can be formed e.g.
  • 2M_(s) + O_2(g) ==> M₂O_2(s)  [redox change, M (0 to +1), O (0 to -1)]
    • shows the formation of the yellow-orange peroxide by Na, K, Rb and Cs
    • each oxygen is in the -1 oxidation state in the peroxide ion O₂^2- 
    • they readily hydrolyse with water forming hydrogen peroxide
      • M₂O_2(s) + 2H₂O_(l) ==> 2MOH_(aq) + H₂O_2(aq)  (not a redox change)
  • M_(s) + O_2(g) ==> MO_2(s)  shows the formation of the 'superoxide' by K, Rb and Cs
    • oxidation number changes are M from 0 to +1 as expected, but on average each oxygen changes from 0 to  -¹/₂  in the superoxide ion O₂^- 
    • 2MO_2(s) + 2H₂O_(l) ==> 2MOH_(aq) + H₂O_2(aq) + O_2(g)  (redox change)
    • oxidation state changes: M and H no change (+1), four O's change from  -¹/₂  in superoxide ions to two of -1 in the peroxide molecule and two at zero in the oxygen molecule.
      • This is a case of disproportionation where the oxidation state of an element gives a higher and lower state product from the same 'original species'.
• The simple oxides readily dissolve in acids and are neutralised to form salts.
  • M₂O_(s) + 2HCl_(aq) ==> 2MCl_(aq) + H₂O_(l)   (M = Li, Na, K, Rb, Cs)
    • to give the soluble chloride salt
    • ionically: (M⁺)₂O^2-_(s) + 2H⁺_(aq) ==> 2M⁺_(aq) + H₂O_(l)  (not a redox change)
      • Acid-base reaction, acid donates proton the oxide ion base, applies to all four examples.
      • The chloride Cl^-, nitrate NO₃^- and sulphate SO₄^2- are spectator ions.
  • M₂O_(s) + 2HNO_3(aq) ==> 2MNO_3(aq) + H₂O_(l) to give the soluble nitrate salt
  • M₂O_(s) + H₂SO_4(aq) ==> M₂SO_4(aq) + H₂O_(l) to give the soluble sulphate salt
  • M₂O_(s) + 2CH₃COOH_(aq) ==> 2CH₃COOM_(aq) + H₂O_(l) to give the soluble ethanoate salt

• The reaction of Group 2 metals with oxygen (a redox reaction)
• Group 2 metals:  2M_(s) + O_2(g) ==> 2MO_(s)    (M = Be, Mg, Ca, Sr, Ba)
  • shows the formation of the oxide expected from their position in the periodic table when the element is heated or burned in air. Oxidation state changes: M from 0 to +2, and oxygen from 0 to -2.
  • The oxide, apart from beryllium, is slightly soluble in water forming the alkaline hydroxide, which increases in strength of basic character down the group.
    • MO_(s) + H₂O_(l) ==> M(OH)_2(s/aq)   (not a redox change, M = Be, Mg, Ca, Sr, Ba)
      • ionically: M²⁺O^2-_(s) + H₂O_(l) ==> M(OH)_2(s/aq)
        • if the hydroxide is soluble: M²⁺O^2-_(s) + H₂O_(l) ==> M²⁺_(aq) + 2OH^-_(aq)
        • Bronsted-Lowry acid-base reaction, the oxide base accepts proton from the water.
        • The mixture of magnesium hydroxide and water is sometimes called milk of magnesia.
        • The formation of calcium hydroxide (slaked lime) when water is added to calcium oxide (quicklime) is very exothermic!
        • The pH of the resulting solution ranges from ~pH 10 to ~pH 13 for Mg(OH)₂ to Ba(OH)₂
    • All the oxides are basic and readily neutralised by acids (not a redox change).
      • MO_(s) + 2HCl_(aq) ==> MCl_2(aq) + H₂O_(l)   (M = Be, Mg, Ca, Sr, Ba)
        • to give the soluble chloride salt
        • ionically: M²⁺O^2-_(s) + 2H⁺_(aq) ==> M²⁺_(aq) + H₂O_(l) 
        • This applies to all four acid reactions examples in this section, acid proton donation to the oxide ion base.
        • In each case the chloride Cl^-, nitrate NO₃^- and sulphate SO₄^2- are spectator ions.
    • MO_(s) + 2HNO_3(aq) ==> M(NO₃)_2(aq) + H₂O_(l)   (M = Be, Mg, Ca, Sr, Ba)
      • to give the soluble nitrate salt
    • MO_(s) + H₂SO_4(aq) ==> MSO_4(aq/s) + H₂O_(l)    (M = Be, Mg, Ca, Sr, Ba)
      • to form the sulphate salt (soluble => insoluble)
      • but reaction increasingly slower for calcium oxide ==> barium oxide as the sulphate becomes less insoluble.
    • MO_(s) + 2CH₃COOH_(aq) ==> (CH₃COO)₂M_(aq) + H₂O_(l)  (M = Be, Mg, Ca, Sr, Ba)
      • to give the ethanoate salt
    • Beryllium oxide BeO is amphoteric (another Be Gp 2 anomaly) and dissolves in strong bases like sodium hydroxide.
      • The equation below shows the formation of a hydroxo beryllate complex ion (not a redox change).
      • BeO_(s) + 2NaOH_(aq) + H₂O_(l) ==> Na₂[Be(OH)₄]_(aq) (beryllate salt)
      • ionically: Be²⁺O^2-_(s) + 2OH^-_(aq)  + H₂O_(l) ==> [Be(OH)₄]^2-_(aq)  

---

7.6. Reaction of s-block metals and water & their hydroxide (OH^-) chemistry

The oxides and hydroxides are usually white ionic solids.

• The reaction of group 1 metals with water (a redox reaction)
• Group 1 metal hydroxide formation
  • 2M_(s) + 2H₂O_(l) ==> 2M⁺OH^-_(aq) + H_2(g)   (M = Li, Na, K, Rb, Cs)
  • shows the formation of the alkaline metal hydroxide and hydrogen.
    • Oxidation state changes: M from 0 to +1, one H per water remains unchanged in oxidation number and one changes from +1 to 0 in H₂.
  • M = Li (slow at first), Na (fast), K (faster - may ignite hydrogen to give a lilac coloured flame* from hot potassium atoms), Rb, Cs, Fr (very explosive) i.e. the reactivity increases down the group.
    • The reactivity trend is explained in section 7.4
    • * 2H_2(g) + O_2(g) ==> 2H₂O_(l)  i.e. the chemistry of the lit splint pop!
  • The hydroxides, MOH, are white ionic solids, all very soluble (except  LiOH), strong bases, getting stronger down the group.
    • All Group 1 hydroxides are soluble in water giving strongly alkaline solutions,
    • and their aqueous solutions readily neutralised by acids (not a redox change) e.g.
    • MOH_(aq) + HCl_(aq) ==> MCl_(aq) + H₂O_(l)   (M = Li, Na, K, Rb, Cs)
      • to give the soluble chloride salt*
      • ionically: OH^-_(aq) + H⁺_(aq) ==> H₂O_(l)  
      • an acid-base reaction, same for all four examples in this section
    • MOH_(aq) + HNO_3(aq) ==> MNO_3(aq) + H₂O_(l)   (M = Li, Na, K, Rb, Cs)
      • to give the soluble nitrate salt
    • 2MOH_(aq) + H₂SO_4(aq) ==> M₂SO_4(aq) + 2H₂O_(l)   (M = Li, Na, K, Rb, Cs)
      • to give the soluble sulphate salt
    • MOH_(aq) + CH₃COOH_(aq) ==> CH₃COOM_(aq) + H₂O_(l)   (M = Li, Na, K, Rb, Cs)
      • to give the soluble ethanoate salt CH₃COO^-M⁺
    • * The hydroxide solutions are readily titrated with standardised hydrochloric acid (burette) using phenolphthalein indicator, the colour change is from pink to colourless.

• The reaction of group 1 metals with water (a redox reaction)
• Group 2 metal hydroxide formation
  • M_(s) + 2H₂O_(l) ==> M(OH)_2(aq/s) + H_2(g)    (M = Mg, Ca, Sr, Ba)
  • shows the formation of the hydroxide and hydrogen with cold water.
  • ionically: M_(s) + 2H₂O_(l) ==> M²⁺_(aq) + 2OH^-_(aq) + H_2(g)
  • oxidation number changes, M is 0 to +2, for one H per water it changes from +1 to 0 in H₂.
  • M = Be (no reaction, anomalous), Mg (very slow reaction), Ca, Sr, Ba (fast to very fast).
    • i.e. the reactivity increases down the group.
    • The reactivity trend for Group 2, and its explanation, are similar to that above for the Group 1 Alkali Metals.
    • The reactivity trend for s-block metals is explained below
    • Magnesium hydroxide and calcium hydroxide (limewater) are sparingly soluble, but the solubility increases down the group, so barium hydroxide is moderately soluble.
    • As previously mentioned, a mixture of magnesium oxide/hydroxide and water is sometimes called milk of magnesia and the saturated aqueous solution of calcium hydroxide is called limewater.
• If the metal is heated in steam the oxide is formed:
  • e.g. Mg_(s) + H₂O_(g) ==> MgO_(s) + H_2(g) 
    • NOT an experiment you would do with Alkali Metals! but beryllium gives little reaction.
  • The oxide is formed because the hydroxide is thermally unstable at higher temperatures
    •   M(OH)_2(s) ==> MO_(s) + H₂O_(g)    (M = Be, Mg, Ca, Sr, Ba)
• REACTIVITY TREND THEORY: The Group 1/2 metal gets more reactive down the group because ...
  • When an alkali metal atom reacts, it loses an electron to form a singly positively charged ion.
    • e.g. Na ==> Na⁺ + e^-
    • in terms of electrons 2.8.1 ==> 2.8 and so forming a stable ion with a noble gas electron arrangement.
  • As you go down the group from one element down to the next the atomic radius gets bigger due to an extra filled electron shell as you go down from one period to the next one.
  • This means the outer electron is further and further from the nucleus.
  • This also means the outer electron is also shielded by the extra full electron shell of negative charge.
  • Due to this shielding the effective nuclear charge on the external electron is ~ +1 (~ proton number - number of noble gas inner core electrons).
  • Further more, the effective nuclear charge of ~+1 is acting over a larger 'surface area' as the atomic radius increases.
  • Therefore both of these factors combine to make the outer electron less and less strongly held by the positive nucleus as the atomic number increases (down the group).
  • So, the outer electron is more easily lost, and the M⁺ ion more easily formed, and so the element is more reactive as you go down the group - best seen in the laboratory with their reaction with water.
  • The reactivity argument mainly comes down to increasingly lower ionisation energy down the group (i.e. ease of ion formation) and a similar argument applies to the Group 2 metals, but two electrons are removed to form the cation.
  • The enthalpy change in forming the hydrated cation from the solid metal does not appear to be as important here.
    • At a more advanced and detailed level, this change can be theoretically split into the
    • enthalpies of (i) atomisation, (ii) ionisation, (iii) hydration of gaseous ion ... (BUT not here!).
  • The reactivity trend is also paralleled by the increasingly negative half-cell potential (E^θ_M/M+ and E^θ_M/M2+) down groups, 1 and 2 i.e. increasing potential to acts as a reducing agent - an electron donor.
  • As with water, the reaction of a group 1/2 metal with oxygen or halogens gets more vigorous as you descend the group.
• All the hydroxides are basic with increasing strength down the group and readily neutralised by acids (not redox reactions). Magnesium hydroxide is sparingly soluble in water but the solubility increases down the group.
  • M(OH)_2(aq/s) + 2HCl_(aq) ==> MCl_2(aq) + 2H₂O_(l)    (M = Be, Mg, Ca, Sr, Ba)
    • to give the soluble chloride salt*
    • all base (OH^-) ... acid (H⁺) reactions
    • ionically if soluble the reaction is: OH^-_(aq) + H⁺_(aq) ==> H₂O_(l)  
    • ionically if insoluble: M²⁺(OH^-)_2(s) + 2H⁺_(aq) ==> M²⁺_(aq) + 2H₂O_(l)
  • M(OH)_2(aq/s) + 2HNO_3(aq) ==> M(NO₃)_2(aq) + 2H₂O_(l)    (M = Be, Mg, Ca, Sr, Ba)  
    • to give the soluble nitrate salt
  • M(OH)_2(aq/s) + H₂SO_4(aq) ==> M₂SO_4(aq/s) + 2H₂O_(l)    (M = Be, Mg, Ca, Sr, Ba)  
    • to give the sulphate salt
  • M(OH)_2(aq/s) + 2CH₃COOH_(aq) ==> (CH₃COO)₂M_(aq) + 2H₂O_(l)    (M = Be, Mg, Ca, Sr, Ba)  
    • to give the ethanoate salt
  • * Saturated calcium hydroxide solution (limewater) can be titrated with standardised hydrochloric acid (burette, low molarity) to determine its solubility. You normally use phenolphthalein indicator and the end-point colour change is from pink to colourless.
• The Group 2 hydroxides, M(OH)₂, get more soluble down the group:
  • If the hydroxide more or less insoluble (e.g. for Be and Mg), they can be made by adding excess sodium/potassium hydroxide solution to a solution of a soluble salt of a Group 2 metal e.g. three 'double decompositions' are shown below ...
    • (i) calcium chloride + sodium hydroxide ==> sodium chloride + calcium hydroxide
      • CaCl_2(aq) + 2NaOH_(aq) ==> 2NaCl_(aq) + Ca(OH)_2(s) 
    • (ii) magnesium sulphate + potassium hydroxide ==> potassium sulphate + magnesium hydroxide
      • MgSO_4(aq) + 2KOH_(aq) ==> K₂SO_4(aq) + Mg(OH)_2(s) 
    • (iii) beryllium nitrate + sodium hydroxide ==> sodium nitrate + beryllium hydroxide
      • Be(NO₃)_2(aq) + 2NaOH_(aq) ==> 2NaNO_3(aq) + Be(OH)_2(s) 
    • or ionically: M²⁺_(aq) + 2OH^-_(aq) ==> M(OH)_2(s) for any Group 2 metal M
    • All the hydroxides are white powders or white gelatinous precipitates.
• Beryllium hydroxide is amphoteric (an anomaly in the group), because apart from the reactions above, it dissolves in strong alkalis like sodium hydroxide to form a hydroxo-complex ion salts called 'beryllates' e.g.
  • Be(OH)_2(s) + 2NaOH_(aq) ==> Na₂[Be(OH)₄]_(aq)  (not a redox change)
  • ionically: Be(OH)_2(s) + 2OH^-_(aq) ==> [Be(OH)₄]^2-_(aq) showing formation of a complex ion
• For the reaction of Group 1 and 2 hydroxides with carbon dioxide to form the carbonates and hydrogen carbonates, see section 7.9
• For the thermal decomposition of nitrates see section 7.11
• See also VOLUMETRIC TITRATION QUESTIONS involving acids and group1 /2 alkaline hydroxides

---

7.7. The reaction of s-block metals with acids

• Group 1 metals are far too reactive to contemplate adding them to acids in a school laboratory!
• Group 2 metals, apart from beryllium (another anomaly), readily react with acids, with increasing vigour down the group (explanation in section 7.4). A redox reaction to form the soluble chloride salt.
  • M_(s) + 2HCl_(aq) ==> MCl_2(aq) + H_2(g)   (M = Mg, Ca, Sr, Ba)
    • ionically for all four examples: M_(s) + 2H⁺_(aq) ==> M²⁺_(aq) + H_2(g) 
    • oxidation state changes: one M at (0) and two H's at (+1) ==> one M (+2) and two H's at (0)
      • the metal is oxidised, electron loss, increase in oxidation state
      • hydrogen ions are reduced, electron gain, decrease in oxidation state
  • M_(s) + 2HNO_3(aq) ==> M(NO₃)_2(aq) + H_2(g)
    • to form the soluble nitrate salt
    • Looks ok in principle, and does this with Mg and very dilute nitric acid, but rarely this simple, the nitrate(V) ion can get reduced to nasty brown nitrogen(IV) oxide gas (nitrogen dioxide, NO₂) and other products, NO gas?, NO₂^- ion?
  • M_(s) + H₂SO_4(aq) ==> MSO_{4(aq/ s)} + H_2(g)
    • to form soluble ==> insoluble sulphate salt
    • The reaction from magnesium to barium becomes increasingly slower as the sulphate becomes less soluble, it coats the metal, inhibiting the reaction.
  • M_(s) + 2CH₃COOH_(aq) ==> (CH₃COO)₂M_(aq) + H_2(g)  to form soluble ethanoate salt
    • This reaction is much slower than the previous three because ethanoic acid is a weak acid (about 2% ionised, so the fizzing appears a lot less vigorous than the other three acids using solutions of similar molarity).
• In aqueous solutions the metal cations formed are hydrated to aqa-complex ions.
  • not quite the simple isolated ions Mⁿ⁺_(aq) which we use in most equations for brevity.
  • e.g. [M(H₂O)₆]ⁿ⁺_(aq) where n=1 for Gp 1 and n=2 for group 2.
    • There may be several layers of water molecules around the ion, so the six is not the whole story, but is typical for the number of 'nearest neighbours', albeit weakly dative covalently bonded water molecules in this case.
    • The six is called the co-ordination number and each water molecule (or anything else attached to the central metal ion) is called a ligand. 
    • The shape of such an ion is 'octahedral' and its simplified structure is shown above on the right. The middle 'blob' is the metal ion and the six outer 'blobs' are the water molecules.
    • (I will replace with proper diagrams later) 
    • However lithium and beryllium are anomalous (M = Li n = 1, or Be n = 2), because of electronic quantum level restrictions, they can have a maximum co-ordination number of four, so their aqueous cations should be written as [M(H₂O)₄]ⁿ⁺_(aq) which has a tetrahedral shape (shown on the left).
• As described above, The soluble groups 1/2 salt solutions contain the hydrated cations derived from the metal:
  • tetra-aqua cations [Li(H₂O)₄]⁺(aq) and [Be(H₂O)₄]²⁺(aq)
  • or the hexa-aqua ions [M(H₂O)₆]⁺(aq) M = Na, K etc. for Group 1
  • and [M(H₂O)₆]²⁺(aq) where M = Mg, Ca etc. for Group 2
  • The tetraaqua beryllium ion and the hexaaqua magnesium ions generate a slight acidity in their salt solutions due to the significant polarising power of the ions (Be²⁺ very small and double charged, Mg²⁺ double charged) e.g.
  • for beryllium: [Be(H₂O)₄]²⁺(aq) + H₂O(l) [Be(H₂O)₃(OH)]⁺(aq) + H₃O⁺(aq) 
  • or magnesium: [Mg(H₂O)₆]²⁺(aq) + H₂O(l) [Mg(H₂O)₅(OH)]⁺(aq) + H₃O⁺(aq) 

---

7.8. The reaction of s-block metals with chlorine & halide (X^-) salts

The salts are white or colourless crystalline solids

• Group 1 metals readily react with halogens (a redox reaction)  
  • e.g. heating the metal in chlorine will cause it to burn forming the chloride
  • 2M_(s) + Cl_2(g) ==> 2MCl_(s)  (redox reaction, M = Li, Na, K, Rb, Cs)
    • Oxidation state changes: M from 0 to +1, X = F, Cl, Br & I from 0 to -1
    • The salt products, M⁺X^-, are white-colourless crystalline ionic solids that dissolve in water to give neutral solutions of about pH 7. The crystalline solids have high melting and boiling points.
    • The solids do not conduct electricity (no mobile ions or electrons) but will conduct and undergo electrolysis when molten or dissolved in water when ions are free to move to electrodes.
  • The halogen is in the -1 oxidation state in the halide ion X^- 
  • The halides of groups 1-2 are important raw materials e.g.
    • sodium chloride ==> sodium hydroxide from rock salt by electrolysis of aqueous solution
    • potassium bromide/iodide ==> elemental bromine/iodine from seawater by oxidation
    • calcium chloride ==> calcium metal by electrolysis of molten chloride
• Group 2 metals (except Be) readily react on heating with halogens (a redox reaction)   
  • e.g. heating in chlorine the chloride is formed
  • M_(s) + Cl_2(g) ==> MCl_2(s)    (M = Mg, Ca, Sr, Ba)
    • Oxidation state changes: M from 0 to +2, X = F, Cl, Br & I from 0 to -1
    • The salt products, M²⁺(X^-)₂,  are similar in properties to the Group 1 M⁺X^- compounds.
    • However, beryllium chloride has a polymeric covalent structure, due to the high polarising influence of beryllium in its +2 oxidation state and the smaller difference in electronegativity between Be-Cl compared to chlorine and the other group 1 and 2 metals.