Alkali and alkaline earth metalsThe metals are all very reactive* when pure. –strongly reducing (very powerful reducing agents) In nature, these elements only ever appear as their

The s-block elements:

Alkali and alkaline earth metals

oThe alkali family is found in the first column of

the periodic table include; Li, Na, K, Rb, Cs, Fr.

o Metallic elements.

oAmong the alkali metals;

sodium and potassium are abundant and

lithium, rubidium and caesium have much

lower abundances.

Francium is highly radioactive; its longest-

lived isotope 223Fr has a half-life of only 21

minutes.

The s-block elements:alkali and alkaline earth metals

Alkali metals

oAtoms of the alkali metals have a single

electron in valence shell (1 valence

electron) and electron configuration of 1s1.

oSoft metals (easily cut with a knife).

oThey are the most reactive metals.

oThey react violently with water.

oAlkali metals are never found as free

elements in nature. They are always

bonded with another element.


Alkali metals

The metals are all very reactive* when pure.

– strongly reducing (very powerful reducing agents)

In nature, these elements only ever appear as their +1 cations.

Their cations (with the exception of lithium) have low charge densities.

– often leads to good solubility and an ability to stabilize anions that can not

be formed along with any other cation.

Cations occur in wide variety of important chemicals.

Form ionic compounds with highly electronegative (nonmetals) found in

Groups 6 (O, S, Se and 7 (F, Cl, Br, I, At) , as well as P and N from Group 5.


Alkali metals

*Most alkali metals are stored under oil. But lithium is so light that it floats on oil. So

a coat of petroleum jelly is applied to lithium before it is stored.

Physical Properties

• Lowest ionization energies and electronegativities in periodic table

- easily ionized to form ions with +1 charge/+1 oxidation state.


Physical Properties


• They are good conductors of electricity.

Their ability to conduct electricity is due to the availability of their outer electrons.

• Low densities.

• Low melting points and boiling points.

Due to the presence of a single valence electron.

Flame tests

oAlkali metal ions are identied by flame tests.

o In flame ions in excited states return to their ground states emit light.


Element Li Na K Rb Cs

Flame colour crimson yellow lilac purple blue

o The emission spectrum depends

on a number of factors;

- the flame temperature;

- but is most strongly is the

energy levels of the emitter, leading

to characteristic colours.

o Overall reactions:

LiCl(l) Li(l) + ½ Cl2(g)

NaCl(l) Na(l) + ½ Cl2(g)


Production of alkali metals

o Lithium sodium and potassium metal are produced by electrolysis of

the molten chlorides.

Reactivity towards Bronsted acids

The alkali metals react even with weak Bronsted acids such as water

– The rate of reaction depends upon the acidity of the medium

2Na + 2H2O 2NaOH + H2

2Na + 2EtOH 2NaOEt +H2

2Na + 2NH3 2NaNH2 +H2

– very slow reaction in absence of catalyst

Chemical reactions


Oxides of the alkali metals

oAlkali metal on combustion in Oxygen

M + O2 → depends on the alkali metal

Li2O (oxide)

Na2O2 (peroxide)

KO2, RbO2, CsO2 (superoxide)

o On combustion in excess of air, lithium forms mainly the oxide, Li2O (plus

some peroxide Li2O2), sodium forms the peroxide, Na2O2 (and some

superoxide NaO2) whilst potassium, rubidium and caesium form the

superoxides, MO2.

oThe increasing stability of the peroxide or superoxide, as the size of the

metal ion increases, is due to the stabilisation of large anions by larger

cations through lattice energy effects.


The oxidation state of O?


o These oxides and peroxides are colourless when pure, but the superoxides

are yellow or orange in colour. The superoxides are also paramagnetic.

Sodium peroxide is widely used as an oxidising agent in inorganic chemistry.

oAll oxides compounds react violently with H2O (hydrolysis) under

formation of hydroxide

Li2O(s) + H2O(l) 2 LiOH(s)

Na2O2(s) + 2 H2O(l) 2 NaOH(s) + H2O2(l)

2 KO2(s) + 2 H2O(l) 2 KOH(s) + H2O2(l) + O2(g)



oThe oxides are all basic. They react readily with acids (even weak acids)

– Na2O + H2O 2NaOH (the water acts as an acid)

– Na2O + HCl 2NaCl

The alkali metal hydroxides are the strongest of all bases.

oThey react with CO2 to form carbonates:

Li2O(s) + CO2(g) 2 Li2CO3(s)

2 Na2O2(s) + 2 CO2(g) 2 Na2CO3(s) + O2(g)

4 KO2(s) + 2 CO2(g) 2 K2CO3(s) + 3 O2(g)

KO2 is used as CO2 and H2O absorber in diving gear.


Reactivity of Alkali Metals with Group 5, Elements (N, P)

o With Phosphorus

12 M + P4 → 4 M3P (phosphides)

o With Nitrogen

6 Li + N2 → 2 Li3N (nitride)

Only Li forms a nitride.

The lattice energy for nitrides with larger alkali metal cations is not

sufficient to makeΔGrxn<0.


Reactivity of Alkali Metals with Group 7 Elements “Halogens”

o The reactivity of alkali metals increases as the ionization energy decreases

Cs > Rb > K> Na > Li

o Reactivity increases moving down the group.

oExample:

o The resulting compounds are colorless, crystalline ionic salts called halides.


Reactivity of Alkali Metals with Group 7 Elements “Halogens”

o The melting and boiling points always follow the trend:

ofluoride > chloride > bromide > iodide.

oAll these halides are soluble in water.

oThe low solubility of LiF in water is due to its high lattice enthalpy.

oThe low solubility of CsI is due to smaller hydration enthalpy of its two

ions. Other halides of lithium are soluble in ethanol and acetone; LiCl is

soluble in pyridine also.


Some typical reactions

Reaction with water:

All alkali metals react with water to produce hydrogen and the metal

hydroxide.

M(s) + H2O(l) M+(aq) + OH-

(aq) + ½ H2(g)

The metals are called alkali metals because the resulting solution is alkaline

owing to the presence of the hydroxide ion.

Reaction of lithium with oxygen:

2Li(s) + ½ O2(g) Li2O(s)

Alkali metal compounds are almost ionic.


Some typical reactions

Reaction with halogens (group 17: F2, Cl2, Br2, I2)

• All alkali metals react with halogens to form ionic halides.

M(s) + 1/2 X2 MX(s)

Ex. 2Na(s) + Cl2(g) 2NaCl(s)

• The most vigorous reaction occurs between the elements which are furthest

apart on the Periodic Table:

- the most reactive alkali metal, francium (Fr), at the bottom of Group 1,

with the most reactive halogen, fluorine, at the top of Group 7.


Activity Series for Alkali Metals

The most active metals are the most easily oxidized. Consequently, they are the

best reducing agents.

Activity decreases going down the group; ease of oxidation.


Basicity of oxides, sulfides nitrides

Salts containing high charge density anions

like O2-, S2-, N3-, P3- with low charge

density cations (alkali metal and most

alkaline earth metals) are basic

– S2- + H2O HS- + OH-

– N3- + 3H2O NH3 + 3 OH-


Hydration

oAn ion in solution is surrounded by water molecules.


o Hydration enthalpy ( hydrH): Enthalpy change

for the transfer of an ion from the gas phase to

solution

M+(g) M+

(aq)

Solubility of alkali metal compounds

oAlkali metals have relatively large, negative enthalpies of hydration.

o Because they carry a single charge, the forces holding their crystals

together, are less strong than those holding together crystals of more highly

charged ions.

oAs a consequence, almost all alkali metal compounds are extremely soluble

in water.

(solubilities often reaching several hundred grams per litre).

o Exception: some lithium compounds with highly charged anions,

example, Lithium phosphate: 0.39 g L-1


Summery of the general reactions of alkali metals

Chemical reactions


Lithium

Lithium is the most reducing of all the alkali metals

– but not the most reactive.

Unlike the other alkali metals; it reacts with nitrogen to form a nitride.



ANOMALOUS PROPERTIES OF LITHIUM

The anomalous behaviour of lithium is due to the:

(i) exceptionally small size of its atom and ion, and

(ii) high polarising power (i.e., charge/ radius ratio).

As a result,

i. there is increased covalent character of lithium compounds which is

responsible for their solubility in organic solvents.

ii. Further, lithium shows diagonal relationship to magnesium which has

been discussed later.


Points of Difference between Lithium and other Alkali Metals

(i) Lithium is much harder. Its m.p. and b.p. are higher than the other alkali

metals.

(ii) Lithium is least reactive but the strongest reducing agent among all the

alkali metals.

On combustion in air it forms mainly monoxide, Li2O and the nitride, Li3N

unlike other alkali metals.

(iii) LiCl is deliquescent and crystallises as a hydrate, LiCl.2H2O whereas

other alkali metal chlorides do not form hydrates.


(iv) Lithium hydrogencarbonate (LiHCO3) is not obtained in the solid form

while all other elements form solid hydrogencarbonates.

(v) Lithium nitrate when heated gives lithium oxide, Li2O, whereas other

alkali metal nitrates decompose to give the corresponding nitrite.

4LiNO3 → 2Li2O + 4 NO2 + O2

2 NaNO3 → 2NaNO2 + O2

(vi) LiF and Li2O are comparatively much less soluble in water than the

corresponding compounds of other alkali metals.


Points of Similarities between Li and Mg

o The similarity between Li and Mg arises

because of their similar sizes:

atomic radii, Li = 152 pm, Mg = 160 pm;

ionic radii : Li+ = 60 pm, Mg2+= 65 pm.


Points of Similarities between Lithium and Magnesium

oThe main points of similarity are:

(i) Both lithium and magnesium are harder and lighter than other elements in

the respective groups.

(ii) Lithium and magnesium react slowly with water.

-Their oxides and hydroxides are much less soluble and their hydroxides

decompose on heating.

-Both form a nitride, Li3N and Mg3N2, by direct combination with

nitrogen.

(iii) The oxides, Li2O and MgO do not combine with excess oxygen to give

any superoxide.


Points of Similarities between Lithium and Magnesium

(iv) The carbonates of Li and Mg decompose easily on heating to form the

oxides and CO2.

-Solid hydrogencarbonates are not formed by Li and Mg.

(v) Both LiCl and MgCl2 are soluble in ethanol.

(vi) Both LiCl and MgCl2 are deliquescent and crystallise from aqueous

solution as hydrates, LiCl·2H2O and MgCl2·8H2O.


Uses

• The most important use of lithium is in rechargeable batteries for mobile

phones, laptops, digital cameras and electric vehicles.

• Lithium metal is used to make useful alloys,

for example with lead to make ‘white metal’ bearings for motor engines,

with aluminium to make aircraft parts.

• Used in anti-depression medication.

• Organometallics of lithium are important.


Li- bipolar disorder

About two million people in the US experience the extreme mood swings of

bipolar disorder.

During a manic phase, they think they can conquer the world. During a

depression, they may feel hopeless.

The ionic compound lithium carbonate often is used to control these

symptoms.

Lithium ions probably have an effect on the transmission of messages

between brain cells.


In 1991, Sony Corporation released the first commercial lithium ion

battery with a carbon anode. Nowadays, many devices use rechargeable

lithium ion batteries and several cathode chemistries have been

developed.

Why Li ??

Lithium is the lightest of all metals;

has the greatest electrochemical potential and

provides the largest specific energy per weight.

Charge

Anode

(Carbon / graphite)

Electrolyte

(LiPF6 in EC:DMC)

Cathode

LiCoO2

Li- ion battery

Anode:

Li Li+ + e-

Cathode:

Co4+ + e- Co3+

The cathode reactions are solid state redox reactions.


Documents

Alkali and alkaline earth metalsThe metals are all very reactive* when pure. –strongly reducing (very powerful reducing agents) In nature, these elements only ever appear as their