We studied the occurrences, extraction, utilizes and physical properties of the alkali metals. In this chapter we will be learning the chemical properties of the alkali metals. We will as well study the oxides, hydroxides sulphides, hydrides carbides and thermal stability of salts of the alkali metals.
Compounds of alkali metal:
In accordance with their highly electropositive character such metals are extremely reactive and are powerful reducing agents reacting with water or most non-metals. They form crystalline ionic salts with high melting and boiling points. Such salts are generally soluble in water giving conducting solution. Now we shall discuss several of the important classes of such salts.
Oxides and Hydroxides:
Since alkali metals are extremely reactive their luster is lost in air due to the formation of oxide with atmospheric oxygen. Three types of oxides are formed through the alkali metals viz, normal oxides having 02 ion and the peroxides having 022-[0 - 0] ion, both of that are dia magnetic and colourlesss. The 3rd one that is coloured and para magnetic is super oxide containing 02- ion. Controlled oxidations of such metals are shown below
2 Li → (1/2 O2 ) + Li2 O→(1/2O2) + Li2O2
2 Na → 1/2 O2 + Na2O →1/2O2 + Na2O2
2K →1/2 O2 + K2O→ 1/2O2 + K2O2 → (O2) +2KO2
2 Rb→ (1/2O2) + Rb2O→ (1/2O2) + Rb2O2 →O2 +2RbO2
2 Cs →1/2O2 + CS2 O→ (1/2O2) + CS2O2 → (O2) + 2CSO2
In the method shown above the products underlined are the main products whenever the metals are burnt in a free supply of air. We might notice in the above method that Lithium forms usual oxide, sodium forms peroxide whereas potassium, rubidium and caesium form superoxide as the main product. All the group one metal oxides are powerfully basic or react vigorously giving hydroxide.
Li O + H2O → 2 Li OH (less violent)
Na2O2 + 2H2 O2 → Na OH + H2O2
2KO2 + 2H2O → 2KOH + H2O2 + O2
As we observe, the peroxides or the super oxides on reaction with water give H2O2 that in turn is a powerful oxidizing agent. Therefore the peroxides and the super oxides are as well oxidizing in nature. The vital strength of the hydroxides raise down the group. As the charge density (charge/size ratio) of the cation decreases between M + and OH- as well decreases. So, OH- can be liberated readily into the solution as we go down the group.
Alkali metals reacts with sulphur to form 2 types of sulphides; simple sulphides of Na2S and polysulphides like Na2Sn, where n - 2, 3, 4, or 6.
These polysulphides contain a zig-zag chain structure as shown below.
Alkali metals react with hydrogen and form ionic hydrides, M+H-. Such hydrides on reaction with water liberate hydrogen. Therefore they are a useful source of hydrogen: MH+ H2O → MOH + H2
Lithium hydride on reaction with AlC13 in other solution forms lithium aluminium hydride which is a useful reducing agent in organic chemistry.
4LiH + AlCl3 → LiAlH4 + 3LiCl
Likewise, sodium hydride forms sodium borohydride that is as well used as a reducing agent.
Lithium reacts with carbon to form ionic carbides, while alike carbides of other metals are not formed on reacting with carbon. They can, though, be formed on heating the metal with acetylene and whenever acetylene is passed through a solution of the metal in liquid ammonia:
Na +C2 H2 → (LiQ.NH3) + Na HC2 → Na2 C2-
Such carbides contain the carbide ion (C = C) 2- on hydrolysis they give acetylene. Therefore, are termed as acetylides;
Na2C2 + 2 H2O → 2 NaOH + C2 H2
Alkali metals as well form covalent compounds such as methylithium, LiCH3 and ethylsodium NaC2H5. These come under the separate class of organometallic compounds.
The main reactions of group 1 elements are sumarized in Table.
→ (2Li + O2) excess +Li2O The higher metals from Na2O2, K2O2, KO2, RBO2, CsO2
2M + S → M2S Very vigorous reaction. Polysulphides are also formed
M + H2O → MOH + 1/2 H2 with Li fairly slow, whereas K explodes
M ROH→MOR + 1/2 H2 Vigorous (R = alkyl, aryl). With Li fair slow.
M + 1/2 H2→ MH At high temperatures. Ionic hydrides. LiH is the most stable
M + 1/2 X2 → MX X = halogen. The higher members can form polyhalides, e.g., KI3.
3Li + 1/2 N2→ Li3N Slow at room temperature: rapid at elevated temperatures.
M + NH3 (1) → [M (NH3) n+ + e-(NH3) → (catalyst) + M+NH2-+1/2H2
2M + C (orC2H2) →M2C2 (acetylides) → H2O+C2H2
M + Hg → amalgams
Table: The reaction of the Group 1 elements
Thermal Stability of Salts:
The case with that a salt decomposes is related to the enthalpy of formation of the salt. (The standard enthalpy of formation of a compound is the enthalpy transform whenever 1 mole of the compound in the standard state is formed from the elements in the standard state).
The enthalpy of formation Hf of a salt MA (M is the metal, A the anion) is given by
Hf = (H+ I) + (H - EA) - Hlatt
Where H is the enthalpy of atomisation, I is the ionisation energy or EA is the electron affinity. Because for any salt in a particular group the terms involving the anion alone remain steady the value of .H°F for these compounds is dependent upon the sum of the enthalpy terms of the particular metal (H atom + Imetal) and the lattice energy .H latt . The larger the lattice energy, the more negative the enthalpy of formation or thus, the more stable is the compound. All such terms become smaller on descending the series from lithium - caesium. The relative stabilities of the salts are thus decided through the parameter that decreases more rapidly - the lattice energy and the sum of the metal enthalpies.
In the salts having tiny anions of high charge density, for example F-, N3- , OH, O2- etc the transform in lattice energy is much dependent on the size of the cation or decreases rapidly on descending the group. Therefore, as the size of the cation amplifies, lattice energy decreases more than the change in the sun of the metal enthalpies. Therefore as we go down the group the constancy of such salts having small anions decreases. Therefore in alkali metal fluorides, the stability reduces in the order LiF > Na F> KF > RbF > CsF.
The opposite trend is examined in the stability of the salts containing large anions of low charge density, for example Br -, I- , NO- etc. In these cases, the lattice energy is relatively insensitive to the transform in cation size and there is more rapid decrease in ionization energy or atonisation enthalpy on descending the group. The lower values of these favour the stability of the compounds. Therefore, the stability of the compounds having large anions amplifies as we move down the group from lithium - caesium.
The stability of the compounds can as well be illustrated via using the theory of polarizing power. The simple scheme is that, as the charge density of the metal ion enhances, the thermal stability of the salts of large polarizing anions, relative to several decomposition product decreases. In common, the least polarizing metal ion is those of the most electropositive metal ions and these form the most stable salts with large anions. In another words small cations form stable salts with small anions or large cations form stable with large anions. Let us obtain for instance carbonate of Group one metals. The carbonates of sodium potassium or caesium are resistant to the heat of a burner flame. Though, lithium carbonate decomposes to its oxide and carbon dioxide under similar conditions.
Li2 CO3 → Li2O + CO2
The tendency of Li2CO3 to undergo thermal decomposition might be illustrated in terms of the achieve in electrostatic attraction which occurs whenever extremely small Li+ ion combines with the smaller oxide ion rather than the much larger carbonate ion. The other carbonates of group (Na - Cs) are more stable since the cations contain a lower charge density and are considerably larger in size or so their decomposition is less favourable energetically.
All the metals except lithium form constant bicarbonates (Lithium bicarbonate is formed only in aqueous solution and has not been isolated). Whenever we heat the alkali metal bicarbonates, they are decomposed to carbonates and simultaneously, carbon dioxide or water is liberated.
2 NaHCO3→ Na2CO3 + CO2 + H2O
The thermal stability of group one hydroxides as well follows a alike trend as that of carbonates. Therefore with the exception of LiOH which on heating d decomposes to Li2O, all other group one hydroxides are stable. Likewise, lithium nitrate as well decomposes on heating to give Li2O NO2 and O2 but all other alkali metal nitrates decompose on strong heating to nitrates liberating oxygen.
4 LiNO3 → 2 Li2O+O2+4NO2
2 NaNO3 → 2 NaNO2 + O2
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