2.

0 GROUP I – ALKALI METALS

 Have one valence electron which is easily lost. The elements are typically soft, and
highly reactive. Lithium as first member shows considerable differences.
2.1 Occurrence

Li occurs in aluminosilicates e.g. LiAl(SiO3)2. Na occurs as rock salt (NaCl), Chile salt petre
(NaNO3), cryolite (Na2AlF6) and in soda ash (NaHCO3/Na2CO3). K occurs in carnallite mineral
(KCl.MgCl2.6H2O), as kainaite (KCl.MgSO4.3H2O) and chile salt petre (KNO3). Rb and Cs
are less common, but they can be obtained in small amounts during crystallization of the other
compounds (i.e. Li, Na and K). Fr is produced during the radioactive decay of actinium. It has
a short half-life so that its chemistry is not well known.
227 223
89𝐴𝑐 ⟶ 87𝐹𝑟 + 29𝐻𝑒

2.2 Extraction

Extraction is by the electrolysis of their fused chlorides. Other chlorides are added to lower
the melting point e.g. addition of 1:4 CaCl2 lowers the melting point of NaCl from 800 0C to
about 600 0C. Graphite is used as an anode to discharge Cl2 and steel is used as cathode to
discharge Na.

𝐴𝑛𝑜𝑑𝑒: 2𝐶𝑙 − ⟶ 𝐶𝑙2(𝑔) + 2𝑒 −

𝐶𝑎𝑡ℎ𝑜𝑑𝑒 2𝑁𝑎+ + 2𝑒 − ⟶ 2𝑁𝑎(𝑠)

−
Overall: 2𝐶𝑙(𝑔) + 2𝑁𝑎+ ⟶ 2𝑁𝑎(𝑠) + 𝐶𝑙2(𝑔)

Cl2 is a valuable co-product.

2.3 Physical properties

 They are soft silvery-grey and good conductors of heat and electricity. They are
highly malleable (sheets)and ductile (wires).
 Are very reactive and have to be stored under oil in order to maintain their physical
properties.
 Exhibit marked photo-electricity i.e. emit electrons when irradiated with light e.g. Cs
is used in photoelectric cells. Electrons are loosely held.
 Metal ions burn with colored flames.
 Table 1. Some physical properties of group I elements
Metal Li Na K Rb Cs
I I I
Electronic [He] 2s [Ne] 3s [Ar] 4s [Kr] 5sI [Xe] 6sI
configuration
Atomic 1.52 1.86 2.27 2.47 2.68
radius (Ǟ)
Ionic radius 0.9 1.18 1.52 1.66 1.81
(Ǟ)
Density 0.53 0.97 0.86 1.53 1.88
(g/cm3)
Ionization 520 496 419 403 376
energy
(kJ/mol)
M.pt (0C) 181 98 63 39 28
Hydration -515 -405 -312 -296 -263
energy
(kJ/mol)
Colour of Crimson Yellow Purple Red Blue
flame, ƛ(nm) red (589.2) (766.5) violet (455.5)
(670.8) (780.0)

 Hydration energies/enthalpy- is the enthalpy change when one mole of isolated

gaseous ions is completely dissolved in water to form one mole of aqueous ions.
Na+(g) Na+(aq) + ΔHθhyd
 All the simple salts of group I metals are soluble in water. They are surrounded with
water molecules (hydrated). Li+ being the smallest, has high charge density and is
heavily hydrated. For this reason, electrical conductivity increases of the ions in
aqueous solution is in the order Cs+ > Rb+ > K+ > Na+ > Li+
 The high hydration energy of Li explains why lithium halides are hydrated (LiX.3H2O),
while others form anhydrous halides (MX, M = Na, K, Rb, Cs).
2.3 Chemical properties

The metals are very reactive and their reactivity increases down the line. With water, Li
reacts quietly, Na and K react vigorously, Rb and Cs very violently.

𝑀(𝑠) + 𝐻2 𝑂(𝑙) ⟶ 𝑀(𝑂𝐻)(𝑎𝑞) + 𝐻2 (𝑔) (M = Li, Na, K, Rb, Cs)

2.3.1 Reaction with air (oxygen)

When heated in air they react to form oxides

4𝐿𝑖 + 𝑂2 ⟶ 2𝐿𝑖2 𝑂 𝑂2− − 𝑂𝑥𝑖𝑑𝑒

𝑁𝑎 + 𝑂2 ⟶ 𝑁𝑎2 𝑂 𝑙𝑖𝑚𝑖𝑡𝑒𝑑 𝑎𝑖𝑟

𝑁𝑎 + 𝑂2 ⟶ 𝑁𝑎2 𝑂2 𝑠𝑢𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑡 𝑎𝑖𝑟 𝑂22− − 𝑝𝑒𝑟𝑜𝑥𝑖𝑑𝑒

𝐾
𝑅𝑏} + 𝑂2 ⟶ 𝑀𝑂2 𝑂2− − 𝑠𝑢𝑝𝑒𝑟𝑜𝑥𝑖𝑑𝑒
𝐶𝑠

Explanation. In size O2- < O22-< O2-. Similarly, Li+ <Na+………..<Cs+. Because of its small
size Li+ is not able to surround itself with large anions like O22- and O2- to give a stable
crystal lattice and consequently only the monoxide exists. However large K, Rb and Cs+ are
able to form structures with the superoxide, the largest of the three oxides.

Note: Large cations stabilize large anions and vice versa.

All the oxides are unstable in water.

O2- + H2O 2OH- e.g Li2O(s) + H2O(l) 2LiOH(aq)

O22- + H2O 2OH- + H2O2 e.g Na2O2(s) + 2H2O(l) 2 NaOH(aq) +

H2O2(l)

O2- + H2O 2OH- + H2O2 + O2 e.g 2KO2(s) + 2H2O(l) 2KOH(aq) + H2O2(l) +

O2(g)

The hydroxides form strongly alkaline solutions. Alkalinity increases down the group.

2.3.2 Reaction with nitrogen

Only Li reacts directly with N2 to form a nitride.

6𝐿𝑖 + 𝑁2 ⟶ 2𝐿𝑖3 𝑁

Both Li+ and N3- are small so that Li3N has high lattice energy. Na+, K+ Rb+ and Cs+ are
large and do not form stable M3N lattice. Li3N reacts with water liberating NH3

𝐿𝑖3 𝑁 + 3𝐻2 𝑂 ⟶ 3𝐿𝑖𝑂𝐻 + 𝑁𝐻3

2.3.3. Reaction with halogens

 The alkali metals readily and vigorously react with halogens to form onic halides.
M(s) + X2(g)→MX(s) (M = Alkali metals; X = F, Cl, Br, I)

 NB: lithium halides are somewhat covalent. It is because of the high polarizing power
of lithium ion (The distortion of electron cloud of the anion by the cation is called
polarisation). The Li+ ion is very small in size and has high charge density, hence very
polarizing. This explains why LiCl is soluble in some organic solvents such as ethanol.
2.3.4. Oxysalts (CO32-, HCO3-, NO3-)

The carbonates are remarkably stable except Li2CO3.

The thermal stability of salts containing a common ion increases down the group e.g.
Li2CO3 Li2O + CO2

Na, K, …….., Cs carbonates are stable

The crystal contained O2- will be more stable (Li2O) than Li2CO3 because the charge density
on O2- ions is greater. The smaller O2- ion can get closer to the small cation (Li+)

Similarly for nitrates

LiNO3 Li2O + NO2 + 1/2 O2

1
NaNO3 NaNO2 + /2 O2

Bicarbonates

The metals form solid bicarbonates which readily decompose to the carbonates.

2MHCO3 M2CO3 + H2O+ CO2

Li is exceptional in that it does not form a solid HCO3-(only exists in solution)

2.3.5 Solutions in liquid ammonia

The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting
in nature. M + (x + y)NH3 [M(NH3)x ]+ + [e(NH3)y ]−

The blue colour of the solution is due to the ammoniated electron ([e(NH3)y]−) which absorbs
energy in the visible region of light and thus imparts blue colour to the solution.

The solutions are paramagnetic and on standing slowly liberate hydrogen resulting in the
formation of amide.

M+(am) + e− + NH3(1) MNH2(am) + ½H2(g) (where ‘am’ denotes solution in

ammonia.).

In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.
Anomalous behavior of lithium

The properties of lithium compounds differ from those of other group one elements and tend
to resemble those of magnesium (diagonal relationship). This is due to;

 The very small atomic and ionic size of lithium,

 The high polarising power of the Li+ ion.

Notable comparison include:

i. LiH is stable to approximately 900 0C where as NaH decomposes at 350 0C

ii. LiCl, LiBr, LiI, and LiClO4 are soluble in some organic solvents e.g. ethanol. This is
because the compounds are more covalent than those of the others members of group
IA. (MgCl2 also soluble in ethanol-diagonal relationship)
iii. Lithium sulfate in contrast to other metal sulfates salts does not form alums. Because
of its size.
The specific compound is the hydrated potassium aluminium sulfate (potassium alum) with
the formula KAl(SO4)2·12H2O. The wider class of compounds known as alums have the related
empirical formula, AB(SO4)2·12H2O.
iv. Forms a monoxide as the only oxide just like magnesium when heated in oxygen.
v. Like group IIA forms a nitride with nitrogen
vi. The carbonates, hydroxides , phosphates and fluorides of Li are sparingly soluble in
water (like Mg)
vii. It’s nitrate decomposes to the normal oxide
viii. Neither lithium nor magnesium form solid bicarbonates (exist only in solution).
ix. When acetylene is passed though heated Li, it does not form acetylide. Others form
acetylides.
x. When ammonia is passed through heated Li, it forms Li2NH, while others form MNH2.
Uses

i) NaCl used as table salt

(i) Because sodium remains liquid over a wide temperature range (97.8–883°C), it is
used as a coolant in specialized high-temperature applications, such as nuclear
reactors.
(ii) Cesium, because of its low ionization energy, is used in photosensors, photocells
and other electronic devices. In these devices, cesium is ionized by a beam of
 visible light, thereby producing a small electric current; blocking the light
interrupts the electric current and triggers a response.
(iii) NaOH/KOH used in a wide variety of industrial processes such as manufacture of
soap; Na2CO3, used in the manufacture of glass; K2O, used in porcelain glazes;
and Na4SiO4, used in detergents. KCl used as fertilizer.
(iv) Li2CO3 is one of the most effective treatments available for manic depression or
bipolar disorder.
(v) Lithium is used in lithium ion batteries. The positive electrode (cathode) half-
reaction in the lithium-doped cobalt oxide substrate is;

CoO2 + Li+ + e- LiCoO2

The negative electrode (anode) half-reaction for the graphite is; LiC6 C6 + Li+ +
e- The full reaction (left to right: discharging, right to left: charging) being;

LiC6 + CoO2 C6 + LiCoO2