S-block: Alkali Metals

ABBREVIATIONS:
- bpt = boiling point, e.c. = electron arrangement/configuration, Pd = period of the Periodic Table, Gp = group of the Periodic Table, max. = maximum, mpt = melting point, ox. state = oxidation state or oxidation number.
INFORMATION on the NOTES
1. The oxidation states quoted are numerically equal to the element's valency. The oxidation state of an element in its normal stable state is 0. The information given assumes this and any reference to the oxidation states of the element refers to its value in compounds.
2. The periodicity of melting/boiling points, atomic radii, 1st ionisation energy and electronegativity and other data for the elements are also tabulated and graphically shown on a separate web pages:
  - Z=1-20 Periods 1-3 + start Pd4, Z=1-38 Periods 1-4 + start Pd5, Z=1- 20 Periods 1-6 + start of Pd 7

7.1. Introduction to Group 1 Alkali Metals and Group 2 Alkaline Earth Metals

The first two vertical columns of the Periodic Table, Groups 1 and 2, are called the s-block metals, because they only have 1 or 2 electrons in their outer shell.
These outer electrons are of an s-orbital type (s sub-shell or sub-quantum level) and the chemistry of the metals, with relatively their low ionisation energies, is dominated by the loss of these s electrons to form a cation.
The outer s¹ electron loss by the Group 1 Alkali Metals to gives the M⁺ ion and the s² electron loss by the Group 2 Alkaline Earth Metals gives the M²⁺ ion.
Consequently most of the compounds of Group 1-2 elements tend to be ionic in nature.
For introduction-revision GCSE notes on Alkali Metals and GCSE Quiz on Alkali Metals and an ASA2 Quiz on the basics of s-block metal chemistry.

7.2. Information and Data Table GROUP 1 ALKALI METALS

property\_Z symbol, name	₃Li Lithium	₁₁Na Sodium	₁₉K Potassium	₃₇Rb Rubidium	₅₅Cs Caesium	₈₇Fr Francium
melting point/^oC	181	98	64	39	29	27
boiling point/^oC	1347	883	774	688	679	677
density/gcm^-3	0.53	0.97	0.86	1.48	1.87	>1.87
1st IE/kJmol^-1	513	496	419	403	376	400
2nd IE/kJmol^-1	7298	4562	3051	2632	2420	2100
atomic radius/pm	152	186	231	244	262	270
M⁺ ionic radius/pm	78	98	133	149	165	180
electronegativity	0.98	0.93	0.82	0.82	0.79	0.70
electron configuration	2,1	2,8,1	2,8,8,1	2,8,18,8,1	2,8,18,18,8,1	2,8,18,32,18,8,1
electron configuration	[He]2s¹	[Ne]3s¹	[Ar]4s¹	[Kr]5s¹	[Xe]6s¹	[Rn]7s¹
Electrode potential M/M⁺	-3.04V	-2.71V	-2.92V	-2.92V	-2.92V	-2.92V
Symbol - flame colour	Li - red/crimson	Na - yellow	K - lilac/purple	Rb - red	Cs - blue	Fr - na

Some of the data is tabulated and plotted on a separate web-page.
Typical metals in some ways e.g. silvery grey lustrous solids*, very good conductors of heat and electricity, relatively high boiling points.
- * When freshly cut, but they rapidly tarnish by reaction with oxygen to form an oxide layer, which is why they are stored under oil.
Untypical in other ways e.g. relatively soft, low density (Li-K float on water before reacting ...), and very low melting points.
Any metal flame colour is due to electronic transitions in the atom or cation.
- Electrons are promoted to higher quantum levels via collisions of the high thermal kinetic energy particles in the hot flame. When the promoted electron 'relaxes' or 'falls' back to its more stable electronic level, energy is emitted (this is the basis of an emission spectrum). If the frequency/wavelength/energy of the photons emitted is in the visible region of the electromagnetic spectrum, a 'flame colour' results e.g. as observed in fireworks.
  - Note: Because the set of quantum level energies are unique for each atom, it means the quantum level difference varies from atom to atom, therefore the frequency of emitted photons is different, hence you see different flame colours in the visible region of light from each Group 1/2 metal.
    - Planck's Equation: ΔE = hν, where
    - ΔE = E₂ - E₁, the energy difference between e.g. the outer s level E₁ and a higher level E₂,
    - h = Planck's constant and ν = frequency of light of the emitted photons.
Oxidation state or oxidation number is always +1 in compounds.
- Only the single outer s-electron is easily lost, the 2nd, and subsequent ionisation energies are far too high to form chemically stable cations of 2+ etc.
  - Detailed notes on oxidation state and redox reaction theory.
- The stable Group 1 cation has the electron configuration of a noble gas,
  - e.g. sodium ion, Na⁺, is 2,8 or 1s²2s²2p⁶ or [Ne]
See also section 4. on Group trends and comparison with Group 2 metals.
- and for help notes on data, see the inorganic data page. (opens in new window)
PLEASE NOTE that Francium is highly radioactive and therefore difficult and dangerous to study BUT all its known physical and chemical properties fit in with it being at the foot of Group 1 and other properties could be inferred from the properties and group trends of Li to Cs.

7.3. Information and Data Table GROUP 2 ALKALINE EARTH METALS

property\_Zsymbol, name	₄Be Beryllium	₁₂Mg Magnesium	₂₀Ca Calcium	₃₈Sr Strontium	₅₆Ba Barium	₈₈Ra Radium
melting pt./^oC	1278	649	839	769	729	700
boiling pt./^oC	2970	1090	1484	1384	1637	1140
density/ gcm^-3	1.85	1.74	1.55	2.54	3.51	5.0
1st IE/ kJmol^-1	900	738	590	550	503	509?
2nd IE/kJmol^-1	1757	1451	1145	1064	965	979
3rd IE/kJmol^-1	14848	7733	4910	4210	3600	3300
atomic radius/ pm	111	160	197	215	217	223
M²⁺ ionic radius/ pm	34	78	106	127	143	152
electronegativity	1.57	1.31	1.00	0.95	0.89	0.89
electron config.	2,2	2,8,2	2,8,8,2	2,8,18,8,2	2,8,18,18,8,2	2,8,18,32,18,8,2
electron config.	[He]2s²	[Ne]3s²	[Ar]4s²	[Kr]5s²	[Xe]6s²	[Rn]7s²
Electrode pot'ial M/M²⁺	-1.97V	-2.36V	-2.84V	-2.89V	-2.92V	-2.92V
Symbol - flame colour (see chemical tests new window)	Be - na	Mg - na	Ca - brick red	Sr - crimson	Ba - apple green	Ra - na

Some of the data is tabulated and plotted on a separate web-page.
Very typical metals, silvery grey lustrous solids, relatively high melting and boiling points, good conductors of heat and electricity.
Compared to adjacent Group 1 metal on same period:
- The melting and boiling points are higher, and they are harder, stronger and more dense than the adjacent Group 1 metal on the same period. This is because their are two delocalised electrons per ion in the crystal lattice giving an overall stronger electrical attraction with the more highly charged M²⁺ ions.
- Chemically very similar e.g. form mainly ionic compounds but different formulae and less reactive because the 1st ionisation energies are higher (due to extra nuclear charge) and a 2nd ionisation energy input to form the stable M²⁺ ion.
Oxidation state or oxidation number is always +2 in compounds.
- The two outer s-electrons are readily lost. The 3rd, and subsequent ionisation energies are far too high to form chemically stable cations of 3+ etc.
- The stable Group 2 cation has electron configuration of noble gas,
  - e.g. calcium ion, Ca²⁺, is 2,8,8 or 1s²2s²2p⁶3s²3p⁶ or [Ar]
See also section 4. below on Group trends and comparison with Group 1 metals.
- and for help notes on data, see the inorganic data page.
PLEASE NOTE that Radium is highly radioactive and therefore difficult and dangerous to study BUT all its known physical and chemical properties fit in with it being at the foot of Group 2 and other properties could be inferred from the properties of Mg to Ba.

7.4. General Trends down groups 1 & 2 with increasing atomic number and formula patterns

Some of the data discussed here, is tabulated and plotted on a separate web-page. (opens in new window)
The 1st ionisation energy (IE) or 2nd etc. decrease: (important to link to reactivity trend)
- because as you go down the group from one element down to the next, on the next period, the atomic radius gets bigger due to an extra filled electron shell. The outer electrons are further and further from the nucleus and are also shielded by the extra full electron shell of negative charge. Therefore the outer electrons are less and less strongly held by the positive nucleus and so less and less energy is needed to remove them.
- Successive ionisation energies always increase e.g. ... 3rd > 2nd > 1st, because the same nuclear charge is attracting fewer electrons and on average closer to the nucleus. BUT note the 2nd IE for Group 1, and the 3rd IE for Group 2, show a particularly significant increase in IE compared to the previous ionisation energy or energies. This is due to removing an electron from an electronically highly stable full inner shell and puts an upper limit on the chemically stable oxidation state.
- Note (probably not needed for exams, but a Q students might raise): Despite the significant increase in atomic number i.e. positive nuclear charge down the group, this effect is outweighed by 'shielding effects' of inner full electron shells and also the nuclear charge is spread over an increasingly larger surface as the atomic radius increases. The effective nuclear charge is NOT what it seems as given by the atomic number and is more related to the number of outer electrons and the size of the atom.
Atomic or ionic radius increases:
- Because from one element to the next, an extra shell of electrons is added, increasing the electron 'bulk' and the outer electrons are increasingly less strongly held (see above).
- The radii of the adjacent Gp 2 atom is smaller than Gp 1 atom on the same period, because the nuclear charge has increased by one unit (L to R in PT), but is attracting electrons in the same shell.
- Similarly the radii of Gp 2 M²⁺ ion is smaller than the adjacent Gp 1 M⁺ ion on the same period, because the nuclear charge has increased by one unit (L to R in PT), but is attracting the same number of electrons in the same shells. (see data tables in section 2. and section 3.)
Generally (but not always), the melting and boiling points fall steadily:
- This is because the ionic radii increase down the group increasing charge separation between the metal cations of the lattice and the free delocalised electrons. This weakens the electrical attractive bonding force and so less thermal KE is needed to weaken the lattice to the 'collapse point' i.e. melting. BUT the situation is not as simple as might be expected, e.g. the metal ions do not always have the same crystal lattice packing arrangement.
The 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. The outer electron is further and further from the nucleus and is also shielded by the extra full electron shell of negative charge. Therefore the outer electron is less and less strongly held by the positive nucleus. This combination of factors means the outer electron is more easily lost, the M⁺ ion more easily formed, and so the element is more reactive as you go down the group. The reactivity argument mainly comes down to increasingly lower ionisation energy down the group* and a similar argument applies to the Gp 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 enthalpies of (i) atomisation, (ii) ionisation, (iii) hydration of gaseous ion (BUT not here!).
The electronegativity tends to decrease:
- The electronegativity values are the lowest of the elements, but there is still a group trend. They get lower because the effective nuclear attractive force on the outer electron charge decreases down the group. You can explain it along the lines of the decreasing 1st IE argument (above), by merely changing the last part of the argument from 'easier to lose electron' to 'weaker attraction of electron charge'.
Formula patterns:
- The general formulae are written in the summary tables in 'simple' format M₂O or ionic formulae (M⁺)₂O^2- where M represents Li to Fr or Be to Ra.
- Since all compounds can be considered ionic, most formulae needed are readily derived in principle by knowing the formula and charge of 10 ions!
- All formulae derive from these 10 ions by equating the total positive charge of the cation with the total negative charge of the anion, and expressing the formula as the simplest whole number ratio.
- The Group 7 halide ion, X^-, can be fluoride F^-, chloride Cl^-, bromide Br^- and iodide I^-.
- The ethanoate ion is included as an illustration of carboxylic acid salts (RCOOH acid ==> RCOO^- in salt) that some GCSE courses introduce.
- The oxides, hydroxides and carbonates and hydrogencarbonates are usually white ionic solids and the Gp1/2 salts listed in the 2nd table are usually white/colourless crystalline ionic solids.
  - The Group 2 hydrogencarbonates do not exist as stable solids and are only modestly stable in aqueous solution at room temperature.
- All relevant equations showing their formation and reactions are in subsequent sections.

cation \ anion	oxide	hydroxide	carbonate	hydrogencarbonate
cation \ anion	O^2-	OH^-	CO₃^2-	HCO₃^-
formula derived from Group1 cation M⁺	M₂O	MOH	M₂CO₃	MHCO₃
formula derived from Group1 cation M⁺	(M⁺)₂O^2-	M⁺OH^-	(M⁺)₂CO₃^2-	M⁺HCO₃^-
formula derived from Group 2 cation M²⁺	MO	M(OH)₂	MCO₃	M(HCO₃)₂
formula derived from Group 2 cation M²⁺	M²⁺O^2-	M²⁺(OH^-)₂	M²⁺CO₃^2-	M²⁺(HCO₃^-)₂

There are also Group 1 hydrogen sulphates of formula MHSO₄, i.e. half neutralised sulphuric acid which are rarely encountered but can be crystallised.
The oxidation numbers-states in the compounds listed in the two tables above and below:
- +1 for metal cation of group 1, +2 for metal cation of group 2
- oxygen -2, hydrogen +1, carbon +4 in table above and below
  - (except ethanoate, carbon oxidation numbers are awkward in organic compounds, leave em' in AS is my advice!)
- halogens e.g. Cl -1, nitrogen +5, sulphur +6
- you need to be able to analyse an anion to understand the relationship between the constituent oxidation states and the charge on the anion, sum of oxidation states = overall charge on ion e.g.
  - carbonate, CO₃^2-, C is +4, 3 O's at -2, sum of ox. states = +4 and -6 = -2 = charge on anion
  - hydrogencarbonate, HCO₃^-, H is +1, C is +4, 3 O's at -2, sum of ox. states = +1 and +4 and -6 = -1 = charge on anion
  - nitrate(V), NO₃^-, N is +5, 3 O's at -2, sum of ox. states = +5 and -6 = -1 = charge on anion
  - sulphate(VI), SO₄^2-, S is +6, 4 O's at -2, sum of ox. states = +6 and -8 = -2 = charge on anion

_cation\^anion	halide	nitrate(V)	sulphate(VI)	ethanoate
_cation\^anion	X^-	NO₃^-	SO₄^2-	CH₃COO^-
formula derived from Group1 cation M⁺	MX	MNO₃	M₂SO₄^2-	CH₃COOM
formula derived from Group1 cation M⁺	M⁺X^-	M⁺NO₃^-	(M⁺)₂SO₄^2-	CH₃COO^-M⁺
formula derived from Group 2 cation M²⁺	MX₂	M(NO₃)₂	MSO₄	(CH₃COO)₂M
formula derived from Group 2 cation M²⁺	M²⁺(X^-)₂	M²⁺(NO₃^-)₂	M²⁺SO₄^2-	(CH₃COO^-)₂M²⁺

The three strong acids mentioned in reactions in sections 5. to 9. are ...
- hydrochloric acid, HCl ==> chloride salts
- nitric acid, HNO₃ ==> nitrate salts
- sulphuric (or sulfuric) acid ==> sulphate salts (or sulfates)
- and the 4th acid is the weak organic carboxylic acid ethanoic acid*, CH₃COOH ==> ethanoates
  - *old name 'acetic acid', and the salts were called 'acetates'

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

The oxides and hydroxides are white ionic solids.

Group 1 metals: 4M_(s) + O_2(g) ==> 2M₂O_(s) (redox change)
- shows the formation of the 'simple' oxide expected from their position in the periodic table when the element is heated 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) (see 6. below for hydroxide chemistry)
  - ionically: (M⁺)₂O^2-_(s) + H₂O_(l) ==> 2M⁺_(aq) + 2OH^-_(aq) (not a redox change)
    - This is an acid-base reaction, the O^2- 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) [a 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) 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

Group 2 metals: 2M_(s) + O_2(g) ==> 2MO_(s) (redox change)
- shows the formation of the oxide expected from their position in the periodic table when the element is heated 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)
    - ionically: M²⁺O^2-_(s) + H₂O_(l) ==> M(OH)_2(s _{or 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!
  - All the oxides are basic and readily neutralised by acids (not a redox change).
    - MO_(s) + 2HCl_(aq) ==> MCl_2(aq) + H₂O_(l) to give the soluble chloride salt
      - ionically: M²⁺O^2-_(s) + 2H⁺_(aq) ==> M²⁺_(aq) + H₂O_(l)
      - This applies to all four examples, acid proton donation to the oxide ion base.
      - The chloride Cl^-, nitrate NO₃^- and sulphate SO₄^2- are spectator ions.
  - MO_(s) + 2HNO_3(aq) ==> M(NO₃)_2(aq) + H₂O_(l)
    - to give the soluble nitrate salt
  - MO_(s) + H₂SO_4(aq) ==> MSO_4(aq _{or s)} _{or s)} + H₂O_(l)
    - 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) to give the ethanoate salt
  - Beryllium oxide BeO is amphoteric (another Be Gp 2 anomaly) and dissolves in strong bases like sodium hydroxide: 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.

Group 1 metal hydroxides: 2M_(s) + 2H₂O_(l) ==> 2M⁺OH^-_(aq) + H_2(g) (redox change)
- 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). The reactivity trend is explained in section 4. above.
  - * 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) 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)
    - to give the soluble nitrate salt
  - 2MOH_(aq) + H₂SO_4(aq) ==> M₂SO_4(aq) + 2H₂O_(l)
    - to give the soluble sulphate salt
  - MOH_(aq) + CH₃COOH_(aq) ==> CH₃COOM_(aq) + H₂O_(l)
    - to give the soluble ethanoate salt
  - * The hydroxide solutions are readily titrated with standardised hydrochloric acid (burette) using phenolphthalein indicator, the colour change is from pink to colourless.

Group 2 metal hydroxides: M_(s) + 2H₂O_(l) ==> M(OH)_2(aq _{or s)} + H_2(g) (redox reaction)
- 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 changes from +1 to 0 in H₂.
- M = Be (no reaction, anomalous), Mg (very slow reaction), Ca, Sr, Ba (fast to very fast).
  - The reactivity trend for Group 2, and its explanation, are similar to that above for the Group 1 Alkali Metals.
  - 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 high temperature:
  - M(OH)_2(s) ==> MO_(s) + H₂O_(g)
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, the solubility increases down the group.
- M(OH)_2(aq _{or s)} + 2HCl_(aq) ==> MCl_2(aq) + 2H₂O_(l) to give the soluble chloride salt*
  - all base (OH^-) - acid (H⁺) reactions; ionically if soluble: OH^-_(aq) + H⁺_(aq) ==> H₂O_(l)
  - ionically if insoluble: M²⁺(OH^-)_2(s) + 2H⁺_(aq) ==> M²⁺_(aq) + 2H₂O_(l)
- M(OH)_2(aq _{or s)}+ 2HNO_3(aq) ==> M(NO₃)_2(aq) + 2H₂O_(l)
  - to give the soluble nitrate salt
- M(OH)_2(aq _or_s) + H₂SO_4(aq) ==> M₂SO_4(aq _{or s)}+ 2H₂O_(l)
  - to give the sulphate salt
- M(OH)_2(aq _or_s) + 2CH₃COOH_(aq) ==> (CH₃COO)₂M_(aq) + 2H₂O_(l)
  - 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 more or less insoluble, 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) barium nitrate + sodium hydroxide ==> sodium nitrate + barium hydroxide
    - Ba(NO₃)_2(aq) + 2NaOH_(aq) ==> 2NaNO_3(aq) + Ba(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 9.
For the thermal decomposition of nitrates see section 11.

7.7. The reaction of s-block metals and 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 4.).
- M_(s) + 2HCl_(aq) ==> MCl_2(aq) + H_2(g) (redox reaction) to form soluble chloride salt
  - 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 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 _{or 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 'nearest neighbours', albeit weakly bonded.
  - 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, 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).

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

The salts are white/colourless crystalline solids

Group 1 metals readily react with halogens:
- e.g. heating the metal in chlorine will cause it to burn forming the chloride
- 2M_(s) + Cl_2(g) ==> 2MCl_(s) (redox reaction)
  - 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:
- e.g. heating in chlorine the chloride is formed
- M_(s) + Cl_2(g) ==> MCl_2(s)
- 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.

7.9. The properties and chemistry of the carbonates (CO₃^2-)

and also the hydrogencarbonates (HCO₃^-)

The carbonates and hydrogencarbonates are white ionic solids

Group 1 carbonates M₂CO₃: Formed on bubbling carbon dioxide into excess hydroxide solution
- 2MOH_(aq) + CO_2(g) ==> M₂CO_3(aq) + H₂O_(l)
- ionic equation: 2OH^-_(aq) + CO_2(g) ==> CO₃^2-_(aq) + H₂O_(l)
- The carbonates are white solids, quite soluble in water, and, apart from lithium, thermally stable to red-heat. (see section 11.)
- Hydrated sodium carbonate, Na₂CO₃.10H₂O, is known as washing soda and is used to soften water by precipitating magnesium and calcium salts as their less soluble carbonates (see section 10).
Group 1 hydrogencarbonates MHCO₃: Formed on bubbling excess carbon dioxide into the hydroxide solution
- The reaction above happens first and then:
  - M₂CO_3(aq) + H₂O_(l) + CO_2(g) ==> 2MHCO_3(aq)
  - ionic equation: CO₃^2-_(aq) + H₂O_(l) + CO_2(g) ==> 2HCO₃^-_(aq)
  - They are white solids, slightly soluble in water, weakly alkaline and readily decompose on heating to form the carbonate and carbon dioxide gas.
  - e.g. at 270^oC: 2NaHCO_3(s) ==> Na₂CO_3(s) + H₂O_(l) + CO_2(g)
  - An alternative to yeast in baking is to use sodium hydrogencarbonate ('sodium bicarbonate' or 'baking soda'). The rising action is also due to carbon dioxide gas formation from reacting with an acid (e.g. an organic acid like tartaric acid) and nothing to do with enzymes:
    - self-raising baking powder = carbonate base + a solid organic acid, giving
    - sodium hydrogencarbonate + acid ==> sodium salt of acid + water + carbon dioxide
Group 1 carbonates and hydrogencarbonates are readily neutralised by acids:
- these are base(proton accepting CO₃^2-or HCO₃^-)-acid(H⁺ from HCl etc.) reactions giving a salt, water and carbon dioxide e.g.
- M₂CO_3(aq) + 2HCl_(aq) ==> 2MCl_(aq) + H₂O_(l) + CO_2(g) to give the chloride salt*
  - ionically for any soluble carbonates: CO₃^2-_(aq) + 2H⁺_(aq) ==> H₂O_(l) + CO_2(g)
- M₂CO_3(aq) + 2HNO_3(aq) ==> 2MNO_3(aq) + H₂O_(l) + CO_2(g) to give the soluble nitrate salt
- M₂CO_3(aq) + H₂SO_4(aq) ==> M₂SO_4(aq) + H₂O_(l) + CO_2(g) to give the soluble sulphate salt
- M₂CO_3(aq) + 2CH₃COOH_(aq) ==> CH₃COOM_(aq) + H₂O_(l) + CO_2(g) to give the soluble ethanoate salt
- MHCO_3(aq) + HCl_(aq) ==> MCl_(aq) + H₂O_(l) + CO_2(g) to give the soluble chloride salt
  - ionically for any soluble hydrogencarbonates: HCO₃^-_(aq) + H⁺_(aq) ==> H₂O_(l) + CO_2(g)
- MHCO_3(aq) + HNO_3(aq) ==> MNO_3(aq) + H₂O_(l) + CO_2(g) to give the nitrate salt
- 2MHCO_3(aq) + H₂SO_4(aq) ==> M₂SO_4(aq) + 2H₂O_(l) + CO_2(g) to give the sulphate salt
- MHCO_3(aq) + CH₃COOH_(aq) ==> CH₃COOM_(aq) + H₂O_(l) + CO_2(g) to give the ethanoate salt
- * The group 1 carbonates e.g. sodium carbonate can be titrated with standardised hydrochloric acid using methyl orange indicator (red in acid, yellow in alkali, the end point is a sort of 'pinky orange').
Group 2 carbonates MCO₃: formed on bubbling carbon dioxide into the hydroxide solution or 'slurry', but beryllium carbonate is not stable (another anomaly). Non of them are very soluble.
- M(OH)_2(aq) + CO_2(g) ==> MCO_3(s) + H₂O_(l)
- When M = Ca, this the reaction of limewater when positively testing for carbon dioxide gas.
- They can also be prepared by a double decomposition precipitation reaction (see section 10.).
- The carbonates decompose on heating to give the oxide and carbon dioxide and exhibit a clear thermal stability trend (see section 11.).
Group 2 hydrogencarbonates M(HCO₃)₂: formed when excess carbon dioxide is bubbled through a slurry of the carbonate and they are only stable in solution and their reaction with acids is not important:
- MCO_3(s) + H₂O_(l) + CO_2(g) M(HCO₃)_2(aq)
  - This the reaction of 'temporary hard water' formation from calcium and magnesium carbonate minerals like limestone and dolomite. Boiling the solution reverses the reaction, so removing the metal cations from solution, and referred to as removing 'temporary hardness'.
Group 2 carbonates MCO₃ readily neutralised by acids to form salt, water and carbon dioxide:
- These are Bronsted-Lowry base (proton accepting CO₃^2-)-acid (H⁺ from HCl etc.) reactions giving a salt, water and carbon dioxide e.g.
- MCO_3(s) + 2HCl_(aq) ==> MCl_2(aq) + H₂O_(l) + CO_2(g) to give the chloride salt
  - ionically for all four examples: M²⁺CO₃^2-_(s) + 2H⁺_(aq) ==> M²⁺_(aq) + H₂O_(l) + CO_2(g)
- MCO_3(s) + 2HNO_3(aq) ==> M(NO₃)_2(aq) + H₂O_(l) + CO_2(g) to give the nitrate salt
- MCO_3(s) + H₂SO_4(aq) ==> M₂SO_4(aq) + H₂O_(l) + CO_2(g) to give the sulphate salt
- MCO_3(s) + 2CH₃COOH_(aq) ==> (CH₃COO)₂M_(aq) + H₂O_(l) + CO_2(g) to give the ethanoate salt
The thermal decomposition of carbonates and nitrates is covered in detail in section 11.

7.10. SOLUBILITY TRENDS of Group 2 compounds - linked to preparations

All the nitrates, M(NO3)₂, are soluble in water.
The hydroxides get more soluble down the group:
- If more or less insoluble, they can be made by adding sodium hydroxide solution to a solution of a soluble salt of M e.g.
  - magnesium chloride + sodium hydroxide ==> sodium chloride + magnesium hydroxide
  - MgCl_2(aq) + 2NaOH_(aq) ==> 2NaCl_(aq) + Mg(OH)_2(s)
    - or ionically: Mg²⁺_(aq) + 2OH^-_(aq) ==> Mg(OH)_2(s)
The sulphates get less soluble down the group:
- If more or less insoluble, they can be made by adding dilute sulphuric acid or sodium sulphate solution to a solution of a soluble salt of M.
- This reaction is used as a test for a sulphate by adding an acidified barium chloride/dil. hydrochloric acid or barium nitrate/dil. nitric acid solution to a solution of the suspected sulphate. A dense white precipitate of barium sulphate forms in a positive result and also illustrates the preparation too e.g.
- barium chloride + sodium sulphate ==> sodium chloride + barium sulphate
- BaCl_2(aq) + Na₂SO_4(aq) ==> 2NaCl_(aq) + BaSO_4(s)
  - or ionically Ba²⁺_(aq) + SO₄^2-_(aq) ==> BaSO_4(s)
The carbonates get less soluble down the group:
- If insoluble, they can be made by adding sodium carbonate solution to a solution of a soluble salt of M e.g. the 'double decomposition' ...
- magnesium chloride + sodium carbonate ==> sodium chloride + magnesium carbonate
- MgCl_2(aq) + Na₂CO_3(aq) ==> 2NaCl_(aq) + MgCO_3(s)
  - or ionically: Mg²⁺_(aq) + CO₃^2-_(aq) ==> MgCO_3(s)
  - You can also use the nitrate and in the case of magnesium, its sulphate too.
  - The spectator ions are Na⁺ and the chloride or sulphate etc. anion from the original group II salt.
  - Hydrated sodium carbonate, Na₂CO₃.10H₂O, is known as washing soda and is used to soften water by using the above reaction.
  - e.g. calcium sulphate (gypsum) + sodium carbonate ==> sodium sulphate (soluble) + calcium carbonate (insoluble)
  - CaSO_4(aq) + Na₂CO_3(aq) ==> Na₂SO_4(aq) + CaCO_3(s) (ionic equation similar to above)
Explanation of solubility trends (usually dealt with later in course e.g. in A2)
- The simplest approach is to consider the two enthalpy change trends.
  - The process of dissolving can be analysed in terms of two theoretical stages e.g. for simple cation-anion ionic compound.
    - In the arguments outlined below Mⁿ⁺ could be Gp1 or Gp2 metal cation etc., X^n- could be halide, oxide, hydroxide, sulphate, carbonate anion etc., and n is the charge on ion - the n's may be different or the same):
  - (1) Mⁿ⁺_aX^n-_b(s) ==> aMⁿ⁺_(g) + bX^n-_(g) (breaking the lattice apart into its constituent ions)
  - This process is always endothermic, and is called the lattice enthalpy. Its usually defined in the opposite direction by saying it is 'the energy released when 1 mole of an ionic lattice is formed from its constituent gaseous ions' (at 298K, 1 atmos./101kPa pressure).
  - * The lattice enthalpy decreases down the group as the cation radius increases (anion radius constant for a particular series e.g. sulphates). Therefore, energetically, the solvation in terms of lattice energy is increasingly favoured down the group.
  - (2) Mⁿ⁺_(g) + aq ==> Mⁿ⁺_(aq) and X^n-_(g) + aq ==> X^n-_(aq)
    - Representing the solvation-hydration of ions.
  - The equations above represent to the two 'hydration enthalpies', the heat released when an isolated gaseous ion becomes solvated by water to form an aqueous solution (1M, 298K, 1 atmos./101kPa pressure)
  - * The hydration enthalpy for the cation decreases down the group as the radius gets larger. Therefore, energetically, the solvation is less favoured down the group as the cation radius increases.
- * In both cases the numerical enthalpy value increases the smaller the radii as charges closer, and the greater the ionic charge (constant for a series), both factors increase the electrical attraction of either cation-anion in the crystal or ion-water in aqueous solution.
We therefore have two competing trends! So, one approach is to say which 'energy change' trend outweighs the other to explain the solubility trend. e.g. for Group 2 hydroxides, energetically, the decrease in lattice enthalpy more than compensates for the decrease in the hydration enthalpy of the M²⁺ cation as it gets larger down the group so leading to greater solubility.
Unfortunately the above is hardly an explanation of a correct prediction! and neither is entropy taken into consideration. The explanations offered are argued after the fact and unsatisfactory! There is no simple explanation possible and ultimately the solubility is dependent on the entropy changes, a notoriously difficult concept area.
If there was an appropriate AS-A2 answer, it would be in the textbooks by now! See below on Jim Clarks website for an intelligent discussion on the matter. Jim's Group 2 pages solubility descriptions and trends and discussions and theory of solubility

7.11. Thermal decomposition & stability trends of Group 1 & 2 compounds

Theory of thermal instability:
- The lower down the metal in the group the more thermally stable is its nitrate or carbonate etc.
- This is because the polarising power* of the cation increases up the group with the smaller ionic radius,
  - AND, in most cases discussed here, the smaller the cation the greater the lattice enthalpy of the oxide formed on decomposition (meaning the oxide is more thermodynamically stable up the group).
  - For particularly the tiny Li⁺ and Be²⁺ ion, the polarising effect considerably reduces the stability of their compounds (does BeCO₃ exist?)
- *The 'polarising power' of a cation is a measure of its electric field effect to attract and distort electron charge on a neighbouring anion:
  - The cation polarising power increases with increase in charge or decreasing radius, both of which increase the intensity of the electric field effect.
- The 'polarisability of an anion' is how easily the electron charge clouds are 'distorted' by a neighbouring cation.
  - The anion is more easily distorted the larger the anion radius and the higher its charge.
The general 'polarising effect' is shown in the diagram below.
- Think of the Mⁿ⁺ cation as the Gp1 or Gp2 cation and the XO₃^n- anion as the nitrate ion or the carbonate ion etc.

The electrical field of the cation distorts or polarises the anion, and at the decomposition temperature, a 'residual' oxide ion is attracted to the cation and the rest of the original larger anion is released as a gas or gases.

Notes: (i) The residual oxide ion is smaller and less polarisable,

(ii) the smaller oxide ion means the resulting oxide has a higher lattice enthalpy than the carbonate or nitrate,

and (iii) these reactions become more favoured at higher temperature because of the large increase in the 'systems' entropy when gases formed.

Trends in thermal stability:
- In all cases, for a particular set of e.g. Gp1 or Gp2 compounds, the thermal stability increases down the group as the ionic radius of the cation increases, and its polarising power decreases.
- Group 1 compounds tend to be more thermally stable than group 2 compounds because the cation has a smaller charge and a larger ionic radius, and so a lower polarising power, particularly when adjacent metals on the same period are compared.
Group 1 Carbonates:
- lithium carbonate readily decomposes: Li₂CO_3(s) ==> Li₂O_(s) + CO_2(g)
- but the others are quite stable to red heat, so again lithium is anomalous by its comparative carbonate instability.
Group 2 Carbonates:
- The carbonates thermally decompose into the metal oxide and carbon dioxide gas.
- MCO₃ ==> MO_(s) + CO_2(g)
- This is the reaction that converts calcium carbonate (limestone) into calcium oxide (quicklime) in a limekiln at about 900^oC. BeCO₃ unstable at room temperature, MgCO₃ decomposes at about 400^oC, SrCO₃ at 1280^oC and BaCO₃ at 1360^oC.
Group 1 Nitrates:
- lithium nitrate is the least stable and decomposes readily on heating to form lithium oxide, nitrogen dioxide and oxygen.
- 4LiNO_3(s) ==> 2Li₂O(s) + 4NO_2(g) + O_2(g)
- The other group 1 Alkali Metal nitrates [NO₃^-, nitrate(V)] decompose to form the nitrite [NO₂^-, nitrate(III)] salt and oxygen gas. Lithium is anomalous due to the particularly high polarising power of the Li⁺ ion.
- 2MNO_3(s) ==> 2MNO_2(s) + O_2(g) where M = Na, K, Rb, Cs
- The nitrites, or nitrate(III)'s, are very thermally stable white solids, soluble in water giving neutral solutions in water.

Group 2 Nitrates:
- For M = Mg, Ca, Sr, Ba the nitrate decompose to form the metal oxide, nasty brown nitrogen dioxide [nitrogen(IV) oxide] gas and oxygen gas when strongly heated.
- 2M(NO₃)_2(s) ==> 2MO_(s) + 4NO_2(g) + O_2(g)
Comparing the stabilities of Group 1 and Group 2 compounds:
- Group 1 compounds are more stable than group 2 compounds, because the polarising effect of the group I cation M⁺, is much less than the polarising power of the smaller and more highly charged group II M²⁺ ion, particularly when comparing adjacent metals on the same period.

7.12. some examples of the uses of Group 1 and 2 Metals and their Compounds.

The chlorides are used as sources of metal extraction by electrolysis.
Sodium and magnesium are then used to extract titanium from its chloride by displacement.
Sodium vapour is used in the yellow-orange street lamps.
Sodium chloride 'common salt' is used as a food flavouring and preservative.
Sodium carbonate is used in the manufacture of glass and the treatment of hard water.
Sodium hydroxide (and chlorine and hydrogen) are manufactured by the electrolysis of sodium chloride solution.
Sodium hydroxide is used in the manufacture of soaps, detergents, bleaches, rayon.
Potassium nitrate is used in fertilisers.
Magnesium is used in the manufacture of alloys, particularly those of aluminium.
Magnesium hydroxide is used in antacid indigestion powders to neutralise excess stomach (hydrochloric) acid (equation in section 6.).
Calcium carbonate (limestone) and calcium oxide (quicklime, from thermal decomposition of limestone in kiln) are both used in agriculture to reduce the acidity of soil to improve its fertility.
Limestone is used directly as building and road foundation material.
Limestone is heated with clay (aluminium silicates) to make cement.
Magnesium hydroxide mixed with water ('milk of magnesia') is used as an antacid remedy in indigestion tablets.
Barium sulphate is used in X-ray colonoscopy of the bowel.
and there are lots of other examples if you dig around.

ABBREVIATIONS:
- bpt = boiling point, e.c. = electron arrangement/configuration, Pd = period of the Periodic Table, Gp = group of the Periodic Table, max. = maximum, mpt = melting point, ox. state = oxidation state or oxidation number.
INFORMATION on the NOTES
1. The oxidation states quoted are numerically equal to the element's valency. The oxidation state of an element in its normal stable state is 0. The information given assumes this and any reference to the oxidation states of the element refers to its value in compounds.
2. The periodicity of melting/boiling points, atomic radii, 1st ionisation energy and electronegativity and other data for the elements are also tabulated and graphically shown on a separate web pages:
  - Z=1-20 Periods 1-3 + start Pd4, Z=1-38 Periods 1-4 + start Pd5, Z=1- 20 Periods 1-6 + start of Pd 7