Alkali metals—lithium, sodium, potassium, rubidium, cesium, and francium form highly reactive compounds. Their compounds generally show high solubility in water, strong alkaline nature, and characteristic flame colours. This article explains the characteristics, properties, reactions, and trends of alkali metal compounds, including hydroxides, oxides, carbonates, bicarbonates, nitrates, and chlorides, in a simple and exam-oriented format.
The compounds of alkali metals are mostly ionic in nature. Due to the low ionization enthalpies and large atomic sizes, the atoms of alkali metals form cations readily by losing the valence electrons. Consequently, they form ionic bonds with non-metals of p-block. General characteristics of some of the compounds of alkali metals are discussed in this article.
Oxides and Hydroxides of Alkali Metals
Alkali metals react with oxygen to form three types of oxides, namely oxides, peroxides and superoxides depending upon the metal.
Lithium forms only oxide, Li₂O and some peroxide (Li₂O₂), sodium forms the peroxide, Na₂O₂ (plus some superoxide Na₂O) while potassium, rubidium and cesium form the superoxides (MO₂).
Under appropriate conditions, pure compounds, M₂O, M₂O₂ or MO₂ may be prepared. It may be noted that superoxide ion, (O₂⁻) has three electron bond i.e., it has one unpaired electron. Therefore, superoxides are coloured. For example, LiO₂ and NaO₂ are yellow, KO₂ is orange, RbO₂ is brown and CsO₂ is orange. These are also paramagnetic. Sodium peroxide is also yellow in colour probably due to the presence of some amount of superoxide in it. However, the normal oxides of alkali metals are colourless and diamagnetic.
It may be noted that the oxides of the formula M₂O₃ called sesquoxides have also been prepared. These are dark coloured paramagnetic powders. These are prepared by thermal decomposition of MO₂ (M = K, Rb, Cs). These can also be prepared by oxidation of liquid ammonia solutions of metals or by controlled oxidation of the peroxides. These are considered to be peroxide disuperoxides having the formula [(M⁺)₄O₂²⁻(O₂⁻)₂].
Alkali metals also form oxides of the formula MO₃ (called ozonides); M = Na, K, Rb and Cs. These ozonides on standing decompose to superoxide and oxygen.
2MO₃ ➔ 2MO₂ + O₂
The normal oxides are basic because they dissolve in water to form alkali metal hydroxides. Therefore, the alkali metal oxides are basic in nature.
Metal oxide :
M₂O + H₂O ➔ 2M⁺ + 2OH⁻
Na₂O + H₂O ➔ 2NaOH (aq)
Peroxides and superoxides give hydrogen peroxide also.
Metal peroxide
M₂O₂ + 2H₂O ➔ 2M⁺ + 2OH⁻ + H₂O₂
Na₂O₂ + 2H₂O ➔ 2NaOH + H₂O₂
Metal superoxide
2MO₂ + 2H₂O ➔ 2M⁺ + 2OH⁻ + H₂O₂ + O₂
2KO₂ + 2H₂O ➔ 2KOH + H₂O₂ + O₂
The peroxides and superoxides also act as oxidising agents because they react with water and form H₂O₂ and O₂ respectively. Sodium peroxide is widely used as an oxidizing agent in inorganic chemistry.
All the alkali metals, their oxides, peroxides, and superoxides readily dissolve in water to produce the corresponding hydroxides, which are strong alkalis :
2Na + 2H₂O ➔ 2NaOH + H₂
Na₂O + H₂O ➔ 2NaOH
Na₂O₂ + 2H₂O ➔ 2NaOH + H₂O₂
2KO₂ + 2H₂O ➔ 2KOH + H₂O₂ + O₂
The hydroxides of all alkali metals are white crystalline solids. The hydroxides of alkali metals are strongly basic and their basic strength increases down the group. They are the strongest of all bases and readily dissolve in water with the evolution of much heat. A number of hydrates of the heavier alkali metal hydroxides (e.g., NaOH · nH₂O, where n = 1, 2, 3, 4, 5, 7) have been prepared from their aqueous solutions, but little is known about their structures.
Properties of Alkali Metal Hydroxides
Basic Strength :
The basic strength of these hydroxides increases as we move down the group from Li to Cs :
LiOH < NaOH < KOH < RbOH < CsOH
Explanation: Alkali metal hydroxides behave as strong bases due to their low ionization energies. The M–O bond in M–O–H can easily break to yield M⁺ and OH⁻ ions:
MOH ➔ M⁺ + OH⁻
As we move down the group, the ionization energy decreases and the ionic radius increases, causing the M–OH bond to cleave more easily. Hence, basic strength increases from LiOH to CsOH.
Solubility and Thermal Stability :
All these hydroxides are highly soluble in water and thermally stable, except lithium hydroxide (LiOH), which decomposes on heating :
2LiOH + (Heat) ➔ Li₂O + H₂O
Formation of Salts with Acids :
Being strongly basic, alkali metal hydroxides react with all acids to form salts:
NaOH + HCl ➔ NaCl + H₂O
2NaOH + H₂SO₄ ➔ Na₂SO₄ + 2H₂O
These salts are colorless, ionic solids that are highly soluble in water.
Halides of Alkali Metals
Alkali metals combine directly with halogens under appropriate conditions to form halides of the general formula MX. These halides can also be prepared by the action of aqueous halogen acids (HX) on metal oxides, hydroxides, or carbonates:
M₂O + 2HX ➔ 2MX + H₂O
MOH + HX ➔ MX + H₂O
M₂CO₃ + 2HX ➔ 2MX + CO₂ + H₂O
(where M = Li, Na, K, Rb, or Cs and X = F, Cl, Br, or I)
All these halides are colorless, high-melting crystalline solids having high negative standard enthalpies of formation (ΔHf°).
Enthalpy Trends of Alkali Metals Halides : For fluorides, the enthalpies of formation (ΔHf°) values become less and less negative as we move down the group. Conversely, for chlorides, bromides, and iodides, the enthalpies of formation (ΔHf°) values become more and more negative moving down the group as shown in Table.
| Element | MF | MCl | MBr | MI |
| Li | −612 | −398 | −350 | −271 |
| Na | −569 | −400 | −360 | −288 |
| K | −563 | −428 | −392 | −328 |
| Rb | −549 | −423 | −389 | −329 |
| Cs | −531 | −424 | −395 | −337 |
Relative Stability of Alkali Metals Halides : For any given metal, the negative enthalpy of formation decreases in the order :
Fluoride > Chloride > Bromide > Iodide
Thus, fluorides are the most stable, while iodides are the least stable. The trends in melting points, boiling points, and solubility of alkali metal halides can be understood in terms of polarization effects, lattice energy, and hydration enthalpy of the ions.
Fazan’s Rule of Polarization Effects – Comparison of Ionic and Covalent Character of Alkali Metal Halides
When a cation approaches an anion, the electron cloud of the anion is attracted towards the cation and hence gets distorted. This effect is called polarization. The power of the cation to polarize the anion is called its polarizing power and the tendency of the anion to get polarized is called its polarizability. The greater the polarization produced more is the concentration of electrons between the two atoms thereby decreasing the ionic character or increasing the covalent character of the bond.
The covalent character of any compound, in general, depends upon the following factors:
(i) Size of the Cation
Smaller the cation, greater is its polarizing power and hence larger is the covalent character. For example, the covalent character decreases as the size of the cation increases:
LiCl > NaCl > KCl > RbCl > CsCl
Thus, LiCl is more covalent than KCl.
(ii) Size of the Anion
Larger the anion, greater is its polarizability. This explains why covalent character of lithium halides is in the order:
LiI > LiBr > LiCl > LiF
Thus, Li⁺ being small in size, polarises the anion and results in decrease of the positive charge on Li⁺ ion. Therefore, lithium halides are covalent in nature. For example, LiI is covalent. For the same reason, dipole moment of LiI (6.25 D) is much less than the theoretically expected value of 11.5 D if it were 100% ionic.
(iii) Charge on the Ion
Cation Charge : Greater the charge on the cation, greater is its polarizing power and hence larger is the covalent character. That is why the covalent character of some of the halides increases in the order Na⁺Cl⁻ < Mg²⁺Cl₂ < Al³⁺Cl₃.
Anion Charge and Size : Similarly, greater the charge on the anion, more easily it gets polarized thereby imparting more covalent character to the compound formed. For example, covalent character increases in the order: NaCl < Na₂SO₄ < Na₃PO₄ as the size of the anion increases and hence larger is the covalent character. Thus, the covalent character decreases as the size of the anion decreases.
(iv) Electronic Configuration of the Cation
If two cations have the same charge and size, the one with a pseudo noble gas configuration, i.e., having 18 electrons in the outermost shell has greater polarizing power than a cation with noble gas configuration, i.e., having 8 electrons in the outermost shell. For this reason, CuCl is more covalent than NaCl (although Na⁺ and Cu⁺ have same charge viz +1 and nearly the same size viz Na⁺ (1.02 Å) and Cu⁺ (0.96 Å)).
Lattice Enthalpy of Alkali Metals Halides
Lattice Enthalpy: Lattice enthalpy is defined as the amount of energy required to separate one mole of solid ionic compound into its gaseous ions. Evidently greater the lattice enthalpy, higher is the melting point of the alkali metal halide and lower is its solubility in water.
Hydration Enthalpy of Alkali Metals Halides
Hydration Enthalpy is the amount of energy released when one mole of gaseous ions combine with water to form hydrated ions.
M⁺(g) + aq ➔ M⁺(aq) + hydration enthalpy
X⁻(g) + aq ➔ X⁻(aq) + hydration enthalpy
Higher the hydration enthalpy of the ions, greater is the solubility of the compound in water. Based on the hydration enthalpy values for the alkali metal cations and halide anions, here is the total sum of the hydration enthalpies (ΔhydH°(cation) + ΔhydH°(anion)) for each combination of alkali metal halide (in kJ mol⁻¹).
- Cations: Li⁺ (-506), Na⁺ (-406), K⁺ (-330), Rb⁺ (-310), Cs⁺ (-276)
- Anions: F⁻ (-513), Cl⁻ (-370), Br⁻ (-339), I⁻ (-294)
| Cation \ Anion | F⁻ (-513) | Cl⁻ (-370) | Br⁻ (-339) | I⁻ (-294) |
| Li⁺ (-506) | -1019 | -876 | -845 | -800 |
| Na⁺ (-406) | -919 | -776 | -745 | -700 |
| K⁺ (-330) | -843 | -700 | -669 | -624 |
| Rb⁺ (-310) | -823 | -680 | -649 | -604 |
| Cs⁺ (-276) | -789 | -646 | -615 | -570 |
Further, the extent of hydration depends upon the size of the ion. Smaller the size of the ion, more highly it is hydrated and hence greater is its hydrated ionic radius and less is its ionic mobility (conductance). Thus the order of their ionic radii and mobility are in order as follows :
- Ionic radii (in pm): Li⁺ (76) < Na⁺ (102) < K⁺ (138) < Rb⁺ (152) < Cs⁺ (167)
- Hydration number* : Li⁺ (25.3) > Na⁺ (16.6) > K⁺ (10.5) > Rb⁺ (10.0) > Cs⁺ (9.9)
- Hydrated radius (in pm): Li⁺ (340) > Na⁺ (276) > K⁺ (212) > Rb⁺ (228) = Cs⁺ (228)
- Ionic mobility (at infinite dilution): Li⁺ (33.5) < Na⁺ (43.5) < K⁺ (64.5) < Rb⁺ (67.5) < Cs⁺ (68.0)
*The hydration number is the average nurnber of water molecules associated with the metal ion.
Solubility and Melting Point Trends of Alkali Metals Halides
In the light of above arguments, let us now explain the melting points of alkali metal halides and their solubility in water or organic solvents:
(i) A delicate balance between lattice enthalpy and hydration enthalpy determines the ultimate solubility of a compound in water. For example, low solubility of LiF (0.27 g / 100 g H₂O) is due to its high lattice energy (-1005 kJ mol⁻¹) whereas the low solubility of CsI is due to small hydration energy of the two ions (-670 kJ mol⁻¹).
(ii) The solubility of most of the alkali metal halides except those of fluorides decrease on descending the group since the decrease in hydration enthalpy is more than the corresponding decrease in lattice enthalpy. For example, difference in lattice enthalpy between NaCl and KCl is 67 kJ mol⁻¹ but difference in hydration enthalpy of Na⁺ and K⁺ ions is 76 kJ mol⁻¹. Thus KCl is less soluble in H₂O than NaCl.
(iii) Because of the small size and higher electronegativity, lithium halides except LiF are predominantly covalent and hence are soluble in organic solvents such as alcohol, acetone, ethyl acetate; LiCl is also soluble in pyridine. In contrast, NaCl being ionic is insoluble in organic solvents.
(iv) Due to high hydration enthalpy of Li⁺ ion, lithium halides are soluble in water except LiF which is sparingly soluble due to its high lattice enthalpy. However, as we move down the group, the solubility of alkali metal fluorides increases regularly as we move from LiF to CsF since the decrease in lattice enthalpy more than compensates the decrease in hydration enthalpy.
(v) For the same alkali metal, the melting points decrease in the order :
fluoride > chloride > bromide > iodide
Explanation : For the same alkali metal ion, the lattice energies decrease as the size of the halide ion increases.
For example, lattice enthalpies of NaF, NaCl, NaBr and NaI are 919, 776, 745 and 685 kJ mol⁻¹. As the lattice enthalpies decrease, energy required to break the lattice (melting point) decreases and hence the melting points of sodium halides decrease from NaF (1261 K) ↣ NaCl (1084 K) ↣ NaBr (1028 K) ↣ NaI (944 K).
(vi) For the same halide ion, the melting points of lithium halides are lower than those of the corresponding sodium halides and thereafter they decrease as we move down the group from Na to Cs.
Explanation: The low melting point of LiCl (887 K) as compared to that of NaCl is probably because LiCl is covalent in nature while NaCl is ionic (because Li atom is smaller in size than Na atom). Thereafter, the melting points decrease as we move from NaCl (1084 K) ↣ KCl (1039 K) ↣ RbCl (988 K) ↣ CsCl (925 K) because the lattice enthalpies decrease as the size of the alkali metal atom increases.
Salts of Oxoacids
Since the alkali metals are highly electropositive, therefore, their hydroxides are very strong bases and hence they form salts with all oxoacids (H₂CO₃, H₃PO₄, H₂SO₄, HNO₃, HNO₂ etc.). They are generally soluble in water and stable towards heat. Here, we shall discuss only carbonates and bicarbonates.
The carbonates (M₂CO₃) of alkali metals are remarkably stable up to 1273 K, above which they first melt and then eventually decompose to form oxides. Li₂CO₃, however, is considerably less stable and decomposes readily:
Li₂CO₃ ➔ Li₂O + CO₂
Being strongly basic, alkali metals also form solid bicarbonates. No other metals form solid bicarbonates, though NH₄HCO₃ also exists as a solid. Lithium, however, does not form solid bicarbonate though it does exist in solution. All the bicarbonates on gentle heating undergo decomposition to form carbonates with the evolution of CO₂ :
2 MHCO₃ ➔ M₂CO₃ + CO₂ + H₂O
As the electropositive character increases down the group, the stability of carbonates and bicarbonates increases.
All the carbonates and bicarbonates are soluble in water and their solubilities increase rapidly on descending the group. This is due to the reason that their lattice energies decrease more rapidly than their hydration energies on moving down the group.
Physical Properties of Selected Alkali Metal Halides
The lattice enthalpies, hydration enthalpies, solubilities, and melting points of various alkali metal halides are summarized in the table below:
| Compound | Lattice Enthalpy(kJ mol⁻¹) | Hydration Enthalpy(kJ mol⁻¹) | Solubility(g / 100 g H₂O) | Melting Point(K) |
| LiCl | −845 | −876 | 63.7 | 887 |
| NaCl | −770 | −776 | 35.7 | 1084 |
| KCl | −703 | −700 | 34.7 | 1039 |
| RbCl | −674 | −671 | 77.0 | 988 |
| CsCl | −644 | −646 | 162.0 | 925 |
| NaF | −893 | −919 | 4.22 | 1261 |
| NaBr | −730 | −745 | 116.0 | 1028 |
| NaI | −685 | −685 | 184.0 | 944 |
| LiF | −1005 | −1019 | 0.27 | 1115 |
| CsI | −582 | −670 | 44.0 | 894 |
Questions, Answers, and Conceptual Explanations
This section provides conceptual short questions and answers based on the characteristics of compounds of alkali metals (Group 1), covering their oxides, hydroxides, halides, carbonates, bicarbonates, nitrates, and other important compounds.
When a cation is highly polarising ? Which alkali metal ion has the highest polarising power ?
A cation is highly polarising if its charge/radius ratio is very high. Li⁺ ion has the highest polarising power among the alkali metal ions because it has highest charge/ radius ratio.
Arrange the following alkali metal ions in decreasing order of their mobility: Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺
Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺
Arrange the following in order of increasing covalent character: MCl, MBr, MF, and MI (where M = alkali metal)
With increasing size of the anion, covalent character increases and, hence, the order is:
MF < MCl < MBr < MI
Given E° values: Cl⁻ / Cl₂ = +1.36 V, I⁻ / I₂ = +0.53 V, Ag⁺ / Ag = +0.79 V, Na⁺ / Na = –2.71 V, Li⁺ / Li = –3.04 V
Arrange the following ionic species in decreasing order of reducing strength: I⁻, Ag⁺, Cl⁻, Li⁺, Na⁺
The lesser the E° value, the stronger is the reducing agent. Therefore, the decreasing order of reducing strength is:
Li⁺ > Na⁺ > I⁻ > Ag⁺ > Cl⁻
Why is KO₂ paramagnetic?
The superoxide ion, O₂⁻, contains one unpaired electron and therefore, it is paramagnetic. The molecular orbital electronic configuration of O₂⁻ is:
KK(σ2s)²(σ*2s)²(σ2pz)²(π2px)²(π2py)²(π*2px)²(π*2py)¹
This also shows that it has one unpaired electron and hence is paramagnetic.
Why does table salt get wet in the rainy season?
Pure NaCl is only hygroscopic, but table salt is impure NaCl containing impurities of Na₂SO₄, CaCl₂, MgCl₂, and CaSO₄. All of these being hygroscopic absorb moisture from air in the rainy season. As a result, table salt gets wet.
Among alkali metals in aqueous solution, Li⁺ ion has the lowest mobility. Why?
Because of its small size among alkali metals, Li⁺ ions are most highly hydrated in aqueous solution. As a result, among alkali metals, the mass of hydrated lithium ion is the highest and thus it has the lowest ionic mobility.
Lithium has the highest ionization enthalpy in group 1 elements, yet it is the strongest reducing agent. Why?
Because lithium has the highest reduction potential.
Why LiCl is more covalent than KCl ?
Due to its smaller size, Li⁺ is more polarising than Na⁺ (and K⁺) and hence LiCl is more covalent than NaCl (and KCl).
Why LiI has greater covalent character than LiF ?
Due to bigger size, I⁻ is more polarizable than F⁻ and hence LiI is more covalent than LiF.
Why MgCl₂ is more covalent than NaCl ?
Due to higher charge, Mg²⁺ is more polarising than Na⁺ and hence MgCl₂ is more covalent than NaCl.
Why CuCl is more covalent than NaCl ?
Due to pseudo inert gas configuration, Cu⁺ is more polarising than Na⁺ and hence CuCl is more covalent than NaCl.
When is a cation highly polarising? Which alkali metal has the highest polarising power? [NCERT]
A cation is highly polarising if its charge/size ratio is very high. Li⁺ ion has the highest polarising power among the alkali metal ions.
The enthalpy of formation of hypothetical CuCl(s) is theoretically found to be –188 kJ mol⁻¹ and ΔH° for CuCl₂(s) is –795 kJ mol⁻¹. Calculate ΔH° for the disproportionation reaction:
2CuCl(s) ➔ CuCl₂(s) + Cu(s)
ΔH° for the above reaction = ∑ΔH°(products) – ∑ΔH°(reactants)
= ΔH°(CuCl₂) + ΔH°(Cu) – 2ΔH°(CuCl)
= –795 + 0 – (2 × –188 kJ mol⁻¹)
= –795 + 376 = –419 kJ mol⁻¹
Why does the reaction >C–Cl + KF ➔ >C–F + KCl proceed better with KF than with NaF? [NCERT]
The standard enthalpies of formation of alkali metal chlorides become more and more negative as we move down the group, i.e., ΔH° of KCl is more negative (–436 kJ mol⁻¹) than that of NaCl (–400 kJ mol⁻¹). Therefore, the above reaction proceeds better with KF than with NaF.
Why is it that on being heated in excess supply of air, K, Rb, and Cs form superoxides in preference to oxides and peroxides? [NCERT]
K⁺, Rb⁺, and Cs⁺ are large cations and superoxide (O₂⁻) is larger than oxide (O²⁻) and peroxide (O₂²⁻) ion. Since a large cation stabilizes a large anion, therefore, these metals form superoxides in preference to oxides and peroxides.
Why LiCl is soluble in organic solvents ?
Li⁺ has very high polarizing power and, therefore, LiCl is covalent in nature. Being covalent in nature, it is soluble in organic solvents.
Name the metal which floats on water without any apparent reaction with it.
Lithium
Why are lithium salts commonly hydrated and those of other alkali ions usually anhydrous ?
Because of small size, Li⁺ has high charge density and therefore, lithium salts are extensively hydrated.
Why is LiF almost insoluble in water whereas LiCl is soluble not only in water but is also in acetone?
Both Li⁺ and F⁻ have small size and therefore LiF has very high lattice enthalpy. Although hydration enthalpy of LiF is also high but it is less than lattice enthalpy (Lattice enthalpy of LiF = –1045 kJ, hydration enthalpy of LiF = –1034 kJ/mol). But LiCl has lower lattice enthalpy than corresponding hydration enthalpy and therefore, is soluble in water. LiCl has also some covalent character because of greater polarization hence it is also readily soluble in nonpolar solvents such as acetone.
Lithium is the only alkali metal to form nitride directly. Explain.
Li⁺ being small in size can easily stabilize a small nitride ion (N³⁻). Other alkali metal ions of larger size cannot stabilize a small nitride ion as effectively.
An aqueous solution of sodium carbonate gives alkaline tests. Why?
Sodium carbonate in aqueous solution gets hydrolysed to give OH⁻ ions and, therefore, gives alkaline tests:
CO₃²⁻ + H₂O ➔ HCO₃⁻ + OH⁻
Why is Li₂CO₃ decomposed at a lower temperature whereas Na₂CO₃ at a higher temperature?
Li⁺ ion being small in size cannot stabilize the large CO₃²⁻ ion effectively; therefore, Li₂CO₃ is relatively unstable on heating. On the other hand, sodium is larger in size and can easily stabilize large CO₃²⁻ ions to form stable Na₂CO₃. This is due to lattice energy effects.
LiH, LiF, and Li₃N show exceptional thermal stabilities. Explain.
Li⁺ is small in size and has an exceptionally high charge–radius ratio. Therefore, its salts with small anions such as F⁻, H⁻, and N³⁻ are exceptionally stable due to their high lattice enthalpies.
Which out of the following can be used to store an alkali metal and why? (H₂O, C₂H₅OH, Benzene)
Benzene can be used to store an alkali metal because water and ethanol react violently with alkali metals:
Na + H₂O ➔ NaOH + ½H₂
Na + C₂H₅OH ➔ C₂H₅ONa + ½H₂
Why are alkali metals not found free in nature?
Alkali metals are highly reactive and therefore are not found free in nature. They are present in the combined state in the form of halides, oxides, silicates, borates, nitrates, etc.
Complete the following reactions:
O₂²⁻ + 2H₂O ➔
2O₂⁻ + 2H₂O ➔
O₂²⁻ + 2H₂O ➔ 2OH⁻ + H₂O₂
2O₂⁻ + 2H₂O ➔ 2OH⁻ + H₂O₂ + O₂