Group 1 Elements: Alkali Metals

Q: Why are Group 1 elements called alkali metals?

Their oxides and hydroxides dissolve in water to give strongly alkaline solutions, for example sodium hydroxide and potassium hydroxide. The word alkali comes from the Arabic for plant ashes, which were rich in potassium carbonate. All six elements, lithium to francium, share one outer s electron, which is why they behave so alike.

Q: Why is lithium the strongest reducing agent even though it has the highest ionisation enthalpy?

Reducing power in water depends on the total energy to turn the solid metal into the hydrated ion. The tiny lithium ion releases 506 kJ per mole on hydration, far more than any other alkali ion. This outweighs its higher ionisation enthalpy, so the net cost is lowest and its electrode potential is the most negative, minus 3.04 volts.

Q: Why is potassium lighter than sodium?

Density is mass divided by volume. Going from sodium to potassium, the atomic radius jumps from 186 to 227 pm, which increases the atomic volume by more than the gain in mass. So potassium, at 0.86 grams per cubic centimetre, is lighter than sodium at 0.97, and both float on water.

Q: Why does a solution of sodium in liquid ammonia turn blue?

Sodium gives its electron to the ammonia, where it is trapped in a cavity surrounded by ammonia molecules. This ammoniated electron absorbs light at the red end of the visible spectrum, so the solution looks deep blue. The unpaired electron also makes the dilute solution paramagnetic and electrically conducting.

Q: Why are alkali metals stored in kerosene?

They react quickly with oxygen, water vapour and carbon dioxide in air, so they must be kept away from it. Kerosene is an unreactive hydrocarbon that does not dissolve them. Lithium is so light that it floats on kerosene, so it is usually wrapped in paraffin wax or petroleum jelly instead.

Q: Why do alkali metals give colours to a flame?

The heat of the flame excites the loosely held outer electron to a higher energy level. When it falls back, it emits light of a definite wavelength: lithium crimson red, sodium golden yellow, potassium lilac, rubidium red violet and caesium blue. The colours identify the metals in flame tests and flame photometry.

Q: How are alkali metals tested in JEE Advanced?

JEE Advanced usually sets statement-based questions with one or more correct options: electrode potential versus reactivity with water, oxide formed on burning, colours and magnetism of liquid ammonia solutions, and density or mobility orders. Learn every trend together with its reason and its exception, because the wrong options are built from the exceptions.

Q: Which alkali metal exceptions are most asked in JEE Advanced?

Five come up again and again in JEE Advanced: potassium is lighter than sodium, lithium has the most negative electrode potential but reacts slowest with water, sodium has the least negative potential, lithium alone forms a nitride and a normal oxide, and the lithium ion is the least mobile in water despite being the smallest.

ChemistryThe s-Block ElementsFor JEE aspirants

The alkali metals (Li, Na, K, Rb, Cs and radioactive Fr) make up Group 1 of the s-block. Each has a single $n s^{1}$ electron outside a noble gas core, so it is lost easily to give $M^{+}$ ions, and these are the most electropositive metals known. This page explains the trends in size, ionisation and hydration enthalpy, the physical properties and flame colours, and how alkali metals react with air, water, hydrogen, halogens and liquid ammonia. JEE Advanced asks the trends with their reasons, and above all the exceptions.

On this page1Group 1 basics2Size and energies3Physical properties4Flame tests5Air and water6

H_{2}

X_{2}

N_{2}

7Reducing nature8Liquid

N H_{3}

9Uses10Solved examples

Key Formulas - Quick Reference

★ Must learnConfiguration [noble gas] $n s^{1}$ ; oxidation state always $+ 1$ (the second IE is huge, e.g. Na 4562 kJ mol $^{- 1}$ ).
★ Must learnDown the group Li → Cs: radius ↑, ionisation enthalpy ↓, hydration enthalpy ↓ in size, m.p. ↓, reactivity ↑.
Hydration enthalpy: $L i^{+} > N a^{+} > K^{+} > R b^{+} > C s^{+}$ ; ionic mobility in water is the reverse.
Density: $Li < K < Na < Rb < Cs$ (K is lighter than Na).
★ Must learnBurning in air: $4Li + O_{2} \to 2L i_{2} O$ , $2Na + O_{2} \to N a_{2} O_{2}$ , $K + O_{2} \to K O_{2}$ (also Rb, Cs).
★ Must learnWater: $2M + 2 H_{2} O \to 2MOH + H_{2}$ ; hydrides $2M + H_{2} \to 2MH$ (ionic, $H^{-}$ ).
★ Must learnReducing power in water ( $E^{\circ}$ ): Li $- 3.04$ V most negative, Na $- 2.71$ V least; $Δ H$ for $M (s) \to M^{+} (aq)$ = sublimation + IE + hydration.
Liquid ammonia: $M + (x + y) N H_{3} \to [M (N H_{3})_{x}]^{+} + [e (N H_{3})_{y}]^{-}$ : blue, conducting, paramagnetic.
Flame colours: Li crimson, Na golden yellow, K lilac, Rb red-violet, Cs blue.

1. The s-Block and Group 1

The s-block contains the elements whose last electron enters the outermost $s$ orbital. An $s$ orbital holds only two electrons, so the block has just two groups: Group 1 ( $n s^{1}$ , alkali metals) and Group 2 ( $n s^{2}$ , alkaline earth metals). Group 1 is lithium, sodium, potassium, rubidium, caesium and francium. The name alkali comes from their hydroxides, which are strongly alkaline.

Figure 1: Group 1 on one card. Size grows, ionisation enthalpy, hydration enthalpy and melting point fall from Li to Cs; density rises except for the K dip (0.86 below Na 0.97). Fr (measured IE 393 kJ mol

^{- 1}

) is radioactive.

Abundance: Na (2.27% by mass of the lithosphere) and K (1.84%) are abundant; Li (18 ppm), Rb (78 ppm) and Cs (2-6 ppm) are rare.
Francium is highly radioactive: its longest-lived isotope, $^{223} Fr$ , has a half-life of only about 22 minutes, so its chemistry is rarely examined.
First member is different: Li differs from the rest of Group 1 and resembles Mg, the second element of Group 2. This diagonal relationship comes from similar ionic size (Li $^{+}$ 76 pm, Mg $^{2 +}$ 72 pm); it is covered on the Lithium page.
In living cells $N a^{+}$ and $K^{+}$ control ion balance and nerve impulses (details on the Sodium page).

Property	Li	Na	K	Rb	Cs	Fr
Atomic number	3	11	19	37	55	87
Atomic mass / g mol $^{- 1}$	6.94	22.99	39.10	85.47	132.91	(223)
Configuration	[He] $2 s^{1}$	[Ne] $3 s^{1}$	[Ar] $4 s^{1}$	[Kr] $5 s^{1}$	[Xe] $6 s^{1}$	[Rn] $7 s^{1}$
Ionisation enthalpy / kJ mol $^{- 1}$	520	496	419	403	376	393
Hydration enthalpy / kJ mol $^{- 1}$	$- 506$	$- 406$	$- 330$	$- 310$	$- 276$	-
Metallic radius / pm	152	186	227	248	265	-
Ionic radius $M^{+}$ / pm	76	102	138	152	167	(180)
m.p. / K	454	371	336	312	302	-
b.p. / K	1615	1156	1032	961	944	-
Density / g cm $^{- 3}$	0.53	0.97	0.86	1.53	1.90	-
$E^{\circ} (M^{+} /M)$ / V	$- 3.04$	$- 2.714$	$- 2.925$	$- 2.930$	$- 2.927$	-

2. Configuration, Size and Energies

2.1 Electronic configuration

Every alkali metal has one valence electron, $n s^{1}$ , outside a noble gas core. This loosely held electron makes them the most electropositive metals: they lose it readily to form $M^{+}$ , and they are therefore never found free in nature, only as compounds (NaCl, KCl, carnallite, spodumene). Removing a second electron would break into the noble gas core, so the second ionisation enthalpy is enormous and the only oxidation state is $+ 1$ .

2.2 Atomic and ionic radii

Alkali metal atoms are the largest in their period: a new shell starts with just one electron and the smallest nuclear charge of that period. Down the group one shell is added each time, so both radii grow from Li to Cs. The $M^{+}$ ion is much smaller than its atom because the whole outer shell is lost.

Figure 2: Atoms and

M^{+}

ions to scale. Both radii increase down the group; Li

^{+}

(76 pm) is the smallest ion, which is the root of lithium's anomalies.

2.3 Ionisation enthalpy

First ionisation enthalpies are low and fall from Li (520) to Cs (376 kJ mol $^{- 1}$ ). Nuclear charge rises down the group, but the extra shells increase the distance and screen the valence electron, and this effect wins. Caesium therefore has the lowest ionisation enthalpy of the stable elements.

2.4 Hydration enthalpy

Hydration enthalpy is the energy released when a gaseous ion is surrounded by water molecules. A small ion attracts water most strongly, so hydration enthalpy decreases in size as the ion grows: $L i^{+} > N a^{+} > K^{+} > R b^{+} > C s^{+}$ . Because $L i^{+}$ holds water so tightly, most lithium salts crystallise as hydrates ( $LiCl \cdot 2 H_{2} O$ ) while the salts of the heavier metals are usually anhydrous.

Figure 3: Both energies fall in size from Li to Cs because the atom and ion grow. Li

^{+}

releases the most hydration energy (

- 506

kJ mol

^{- 1}

), so lithium salts are often hydrated (

LiCl \cdot 2 H_{2} O

Exam Trick Smallest ion, thickest coat. Li

^{+}

is the smallest bare ion but carries the largest shell of water, so its hydrated radius is the largest and it moves slowest in water. Ionic mobility and conductance:

L i^{+} < N a^{+} < K^{+} < R b^{+} < C s^{+}

Key idea

One idea runs through Group 1: size grows down the group. Gas-phase properties (radius, IE) follow size directly; properties in water follow the hydration that size controls.

Quick Recall: tap to check

Why is the second ionisation enthalpy of sodium so large?

The second electron must come out of the complete [Ne] core (

2 p^{6}

), which is much closer to the nucleus and not shielded; Na has IE

_{2}

= 4562 kJ mol

^{- 1}

against IE

_{1}

= 496.

Arrange the hydrated ions by size:

L i^{+}

N a^{+}

K^{+}

L i^{+} (aq) > N a^{+} (aq) > K^{+} (aq)

: the smallest bare ion is the most hydrated.

Why are lithium salts often hydrated but potassium salts not?

L i^{+}

has the highest hydration enthalpy (

- 506

kJ mol

^{- 1}

), enough to keep water in the crystal;

K^{+}

(

- 330

) is not.

3. Physical Properties

All alkali metals are silvery white, soft (they cut with a knife) and light. Each atom gives only one electron to the metallic bond, and the atoms are large, so the metallic bonding is weak: melting and boiling points are low and fall down the group (Cs melts at 302 K). Density rises down the group because atomic mass grows faster than volume, with one exception: K (0.86 g cm $^{- 3}$ ) is lighter than Na (0.97), because potassium's volume jumps more than its mass.

Figure 4: Density rises down the group except K (0.86) below Na (0.97); Li, Na and K are lighter than water. Melting points fall from 454 K (Li) to 302 K (Cs), which melts on a hot day.

3.1 Flame colours

The metals and their salts colour an oxidising flame. Heat promotes the valence electron to a higher level; when it drops back, the energy difference is emitted as visible light of a characteristic wavelength. The colour identifies the metal (flame test), and its intensity measures the amount (flame photometry, atomic absorption spectroscopy).

Metal	Li	Na	K	Rb	Cs
Colour	crimson red	golden yellow	violet (lilac)	red-violet	blue
$λ$ / nm	670.8	589.2	766.5	780.0	455.5

Figure 5: Flame test colours. Heat excites the ns electron; light is emitted as it falls back. The listed K and Rb lines lie at the edge of vision, so their weaker violet lines give the lilac and red-violet flames.

Exam Trick The names are the colours. Bunsen and Kirchhoff discovered Rb and Cs from their flame spectra: rubidus is Latin for deep red and caesius for sky blue. Remember Rb red-violet, Cs blue. To see K in the presence of Na, look through blue cobalt glass, which absorbs the sodium yellow.

3.2 Photoelectric effect

The ionisation enthalpies of K and Cs are so low that even visible light can knock out an electron from the metal surface. This is why potassium and caesium are used in photoelectric cells. Lithium holds its electron much more firmly and needs higher-energy light.

Key idea

Weak one-electron metallic bonding makes alkali metals soft and low melting; density rises except K below Na; flame colours come from the easily excited

n s^{1}

electron.

4. Chemical Properties

Large size and low ionisation enthalpy make the alkali metals highly reactive, and reactivity increases down the group. Every reaction below is the metal handing its $n s^{1}$ electron to something else.

Figure 6: Everything an alkali metal does starts with losing its

n s^{1}

electron. Only lithium reacts with

N_{2}

; the oxide type changes with cation size.

4.1 Reactivity towards air

In dry air the metals tarnish: a film of oxide forms, which takes up moisture to give the hydroxide (and then carbonate with $C O_{2}$ ). Heated in oxygen they burn vigorously, and the product depends on the metal:

4Li + O_{2} oxide 2L i_{2} O

2Na + O_{2} peroxide N a_{2} O_{2}

M + O_{2} superoxide, M = K, Rb, Cs M O_{2}

In every oxide the metal is $+ 1$ ; only the oxygen species changes ( $O^{2 -}$ , $O_{2}^{2 -}$ , $O_{2}^{-}$ ). A large anion such as superoxide is stable only next to a large cation. Lithium also combines directly with the nitrogen of air to form the nitride, $L i_{3} N$ . Because they react with both air and water, the metals are stored under kerosene; lithium (density 0.53) floats on kerosene and is kept in paraffin wax.

Exam Trick Big cation, big oxygen anion. Li (small) keeps the small

O^{2 -}

, Na takes

O_{2}^{2 -}

, and K, Rb, Cs hold the biggest,

O_{2}^{-}

. Matching sizes give the highest lattice energy. Full oxide chemistry is on the Compounds of Alkali Metals page.

4.2 Reactivity towards water

All alkali metals react with water to give the hydroxide and hydrogen:

2M + 2 H_{2} O M = alkali metal 2 M^{+} + 2O H^{-} + H_{2}

Lithium has the most negative $E^{\circ}$ , yet it reacts the least vigorously; sodium, with the least negative $E^{\circ}$ , reacts vigorously, and K, Rb and Cs react explosively. $E^{\circ}$ is a thermodynamic quantity: it tells how far a reaction can go, not how fast. Lithium stays solid (m.p. 454 K) while sodium melts on the water from the heat released, and the tiny, heavily hydrated $L i^{+}$ moves slowly away from the surface. The metals also react with other proton donors, such as alcohols, gaseous ammonia and terminal alkynes, releasing $H_{2}$ :

2Na + 2 C_{2} H_{5} OH 2 C_{2} H_{5} ONa + H_{2}

2Na + 2N H_{3} gaseous N H_{3}, heat 2NaN H_{2} + H_{2}

Figure 7: The lithium paradox. Li has the most negative

E^{\circ}

(thermodynamics) yet reacts the most slowly with water (kinetics): it stays solid (m.p. 454 K) and its tiny, heavily hydrated ion moves slowly. Na has the least negative

E^{\circ}

Thermodynamics:

E^{\circ}

- 3.04

V is the most negative, so lithium is the strongest reducing agent in water. Na

- 2.71

V is the least negative.

Kinetics: the rate you seeLi reacts slowest: high m.p. (no melting, small contact area) and a slow, heavily hydrated ion. K, Rb, Cs react explosively.

4.3 Reactivity towards dihydrogen

Heated in hydrogen at about 673 K (lithium at 1073 K), the metals form ionic hydrides $M^{+} H^{-}$ , white solids with high melting points. The hydride ion is proved by electrolysis of the molten salt: $H_{2}$ is set free at the anode. Hydrides react with water to give $H_{2}$ ( $NaH + H_{2} O \to NaOH + H_{2}$ ), and their stability falls from LiH to CsH as the lattice energy falls.

2M + H_{2} 673 K 2 M^{+} H^{-}

4.4 Reactivity towards halogens

The metals react vigorously with halogens to give ionic halides $M^{+} X^{-}$ . Lithium halides are somewhat covalent: the very small $L i^{+}$ has a high polarising power (ability to distort the electron cloud of an anion). A large anion is distorted most easily, so among the halides LiI is the most covalent.

2M + X_{2} X = F, Cl, Br, I 2MX

4.5 Reducing nature and electrode potential

Alkali metals are strong reducing agents. In water their reducing power is measured by the standard electrode potential $E^{\circ} (M^{+} /M)$ , which reflects the overall change $M (s) \to M^{+} (aq) + e^{-}$ . This change can be broken into three steps:

Sublimation: $M (s) \to M (g)$ , energy absorbed (enthalpy of atomisation).
Ionisation: $M (g) \to M^{+} (g) + e^{-}$ , energy absorbed (ionisation enthalpy).
Hydration: $M^{+} (g) + H_{2} O \to M^{+} (aq)$ , energy released (hydration enthalpy).

Lithium has the highest ionisation enthalpy, but its small ion releases so much hydration energy that the total is the smallest of the group. That is why lithium has the most negative $E^{\circ}$ and is the strongest reducing agent in water, while sodium is the weakest.

Figure 8:

M (s) \to M^{+} (aq) + e^{-}

in three steps. Li pays the most ionisation energy but gets back the most hydration energy, so its net cost (+173 kJ mol

^{- 1}

) is the lowest and its

E^{\circ}

the most negative; Na (+197) is the highest.

JEE Advanced

The three-step sum reproduces the whole $E^{\circ}$ order. Using the enthalpy of atomisation (Li 159, Na 107, K 89, Rb 81, Cs 77 kJ mol $^{- 1}$ ) with the ionisation and hydration enthalpies of the table:

Metal	Sublimation	IE	Hydration	Net $Δ H$	$E^{\circ}$ / V
Li	159	520	$- 506$	173	$- 3.04$
Rb	81	403	$- 310$	174	$- 2.930$
Cs	77	376	$- 276$	177	$- 2.927$
K	89	419	$- 330$	178	$- 2.925$
Na	107	496	$- 406$	197	$- 2.714$

The smaller the net enthalpy, the more negative $E^{\circ}$ : Li < Rb < Cs < K < Na in both columns. Entropy terms are small and similar, so enthalpy alone gives the right order. The same logic explains Exercise 10.26(c) on the Group 2 page.

4.6 Solutions in liquid ammonia

Alkali metals dissolve in liquid ammonia to give deep blue, conducting solutions:

M + (x + y) N H_{3} [M (N H_{3})_{x}]^{+} + [e (N H_{3})_{y}]^{-}

The blue colour comes from the ammoniated electron, which absorbs energy in the visible region (the red end), so the transmitted light is blue. Each solution contains an unpaired electron, so it is paramagnetic. On standing, the solution slowly liberates hydrogen and forms the amide ('am' = in ammonia):

M^{+} (am) + e^{-} + N H_{3} (l) MN H_{2} (am) + ½ H_{2} (g)

In concentrated solution the colour changes to bronze and the solution becomes diamagnetic, because the electrons pair up and metal-ion clusters form. The solvated electron makes these solutions powerful reducing agents: Na in liquid $N H_{3}$ reduces alkynes to trans-alkenes (Birch conditions).

Figure 9:

M + (x + y) N H_{3} \to [M (N H_{3})_{x}]^{+} + [e (N H_{3})_{y}]^{-}

. The unpaired ammoniated electron gives the blue colour and paramagnetism; in concentrated solution electrons pair up (bronze, diamagnetic).

Dilute solutionDeep blue; conducting; paramagnetic (free ammoniated electrons).

Concentrated solutionBronze, metallic lustre; diamagnetic (electrons paired, metal clusters).

4.7 Reactivity towards acids

With acids the reaction is violent (the metals react even with water), giving the salt and hydrogen, for example $2Na + 2HCl \to 2NaCl + H_{2}$ .

Key idea

Reactivity rises down the group, but in water Li is the strongest reducing agent (hydration) and the slowest to react (kinetics). The alkali metal always ends as

+ 1

4.8 Reading any trend question

Figure 10: Ask where the property lives before ranking. Gas phase follows size, water follows hydration, the bulk metal has the K density dip, and compounds follow cation-anion size match.

Quick Recall: tap to check

Why is KO

_{2}

formed but LiO

_{2}

is not?

The superoxide ion is large; it is stabilised only by a large cation (K

^{+}

, Rb

^{+}

, Cs

^{+}

) through a high lattice energy. Small Li

^{+}

forms the normal oxide

L i_{2} O

At which electrode is hydrogen evolved when molten NaH is electrolysed?

At the anode:

2 H^{-} \to H_{2} + 2 e^{-}

, which proves the hydride contains

H^{-}

What makes a sodium-ammonia solution paramagnetic?

The unpaired ammoniated electron

[e (N H_{3})_{y}]^{-}

5. Uses of the Alkali Metals

Metal	Uses
Lithium	alloys: with Pb ('white metal' bearings for motor engines), with Al (aircraft parts), with Mg (armour plates); lithium batteries and other electrochemical cells; thermonuclear reactions
Sodium	Na/Pb alloy to make $PbE t_{4}$ and $PbM e_{4}$ (old anti-knock additives, replaced by lead-free petrol); liquid sodium as coolant in fast breeder reactors
Potassium	vital in living systems; KCl as fertiliser; KOH for soft soap and as an absorbent of $C O_{2}$
Caesium	photoelectric cells (lowest ionisation enthalpy)

Figure 11: Group 1 on one page. Every branch traces back to one loosely held

n s^{1}

electron and to the small size of

L i^{+}

6. Solved Examples

Solved Example 1

What is the oxidation state of K in

K O_{2}

Solution:

$K O_{2}$ is potassium superoxide: it contains the superoxide ion $O_{2}^{-}$ . The compound is neutral, so K is +1. (Assigning $- 2$ to each O would wrongly give K = +4.)

Solved Example 2

The

E^{\circ}

values are:

C l_{2} /C l^{-}

+1.36 V,

I_{2} / I^{-}

+0.53 V,

A g^{+} /Ag

+0.79 V,

N a^{+} /Na

- 2.71

L i^{+} /Li

- 3.04

V. Arrange

I^{-}

, Ag,

C l^{-}

, Li, Na in decreasing order of reducing strength.

Solution:

The reduced form of a couple is a stronger reducing agent the more negative the $E^{\circ}$ of that couple. Ordering the couples from most negative to most positive: Li ( $- 3.04$ ), Na ( $- 2.71$ ), $I^{-}$ (+0.53), Ag (+0.79), $C l^{-}$ (+1.36).

Answer: $Li > Na > I^{-} > Ag > C l^{-}$ .

Solved Example 3

Which alkali metal has the most negative standard electrode potential but reacts the least vigorously with water?
(A) Na
(B) Li
(C) K
(D) Cs

Solution:

Answer: (B). Li has $E^{\circ} = - 3.04$ V (large hydration enthalpy of the small ion) but reacts slowly: it does not melt (m.p. 454 K) and its hydrated ion diffuses slowly. $E^{\circ}$ decides the extent, not the rate.

Solved Example 4

A dilute solution of sodium in liquid ammonia is prepared. Which statements are correct? (One or more options.)
(A) The blue colour is due to ammoniated electrons.
(B) The solution is paramagnetic.
(C) On concentrating, it turns bronze and becomes diamagnetic.
(D) The solution is an electrical insulator.

Solution:

Answer: (A), (B), (C). The ammoniated electron absorbs visible light (blue colour) and is unpaired (paramagnetic); in concentrated solution electrons pair and metal clusters form (bronze, diamagnetic). (D) is wrong: the solution conducts, carried mainly by the mobile solvated electrons.

Solved Example 5

How many of Li, Na, K, Rb and Cs are less dense than water (1.00 g cm

^{- 3}

)? (Integer answer.)

Solution:

Densities: Li 0.53, Na 0.97, K 0.86, Rb 1.53, Cs 1.90 g cm $^{- 3}$ . Three are below 1.00.

Answer: 3 (Li, Na, K).

Solved Example 6

For sodium: enthalpy of atomisation 107, ionisation enthalpy 496, hydration enthalpy of

N a^{+}

- 406

kJ mol

^{- 1}

. For lithium the values are 159, 520 and

- 506

kJ mol

^{- 1}

. Find

Δ H

for

M (s) \to M^{+} (aq) + e^{-}

for each and say which metal has the more negative

E^{\circ}

Solution:

Na: $107 + 496 - 406 = + 197$ kJ mol $^{- 1}$ . Li: $159 + 520 - 506 = + 173$ kJ mol $^{- 1}$ .

Less energy is needed to put lithium into solution as its ion, so lithium has the more negative $E^{\circ}$ ( $- 3.04$ V against $- 2.71$ V), even though its ionisation enthalpy is higher.

Solved Example 7

The correct order of ionic mobility in aqueous solution is
(A)

L i^{+} > N a^{+} > K^{+}

(B)

K^{+} > N a^{+} > L i^{+}

(C)

N a^{+} > L i^{+} > K^{+}

(D)

L i^{+} = N a^{+} = K^{+}

Solution:

Answer: (B). Mobility depends on the size of the hydrated ion. $L i^{+}$ is the most hydrated, so it is the largest in water and the slowest; $K^{+}$ is the least hydrated and the fastest.

Practice Questions

What are the common physical and chemical features of alkali metals? (NCERT 10.1)Answer: One $n s^{1}$ electron, +1 state, large atoms, low IE, highly electropositive; soft, light, low-melting, silvery; characteristic flame colours; strong reducing agents that react with $O_{2}$ , $H_{2} O$ , $H_{2}$ and halogens to give ionic compounds with basic oxides and hydroxides; dissolve in liquid $N H_{3}$ to give blue solutions.
Why are alkali metals not found in nature? (NCERT 10.3)Answer: Their single valence electron is lost so easily that they react with air, water and other substances; they occur only as compounds.
Explain why sodium is less reactive than potassium. (NCERT 10.5)Answer: K is larger and has a lower ionisation enthalpy (419 against 496 kJ/mol), so it loses its valence electron more easily.
Why are potassium and caesium, rather than lithium, used in photoelectric cells? (NCERT 10.9)Answer: Their ionisation enthalpies are low enough for visible light to eject electrons; lithium holds its electron too firmly.
When an alkali metal dissolves in liquid ammonia the solution can acquire different colours. Explain. (NCERT 10.10)Answer: Dilute: blue, from ammoniated electrons absorbing visible light (paramagnetic). Concentrated: bronze and diamagnetic, as electrons pair and metal-ion clusters form.
Comment: the mobilities of the alkali metal ions in water are $L i^{+} < N a^{+} < K^{+} < R b^{+} < C s^{+}$ . (NCERT 10.26a)Answer: Hydration decreases from $L i^{+}$ to $C s^{+}$ , so the hydrated ion shrinks down the group and moves faster.
Which alkali metal has the least melting point? (a) Na (b) K (c) Rb (d) Cs (NCERT 10.30)Answer: (d) Cs, 302 K: the largest atom has the weakest metallic bond.

Common Mistakes to Avoid

Watch out

Writing potassium as denser than sodium: the order is Li < K < Na < Rb < Cs.
Saying lithium reacts most vigorously with water because its $E^{\circ}$ is the most negative. $E^{\circ}$ is thermodynamic; lithium actually reacts the slowest.
Assuming reducing power in water rises smoothly down the group: Na has the least negative $E^{\circ}$ ( $- 2.71$ V).
Giving K an oxidation state of +4 in $K O_{2}$ : the anion is superoxide $O_{2}^{-}$ , so K is +1.
Calling $L i^{+}$ the most mobile ion in water because it is smallest; its hydrated ion is the largest.
Naming LiF the most covalent lithium halide. Polarisation grows with anion size: LiI is the most covalent.
Blaming the blue colour of Na in liquid $N H_{3}$ on $N a^{+}$ or Na atoms: it is the ammoniated electron.
Assuming every alkali metal burns in air to give a nitride. Only lithium forms $L i_{3} N$ .

Frequently Asked Questions

Why are Group 1 elements called alkali metals?

Their oxides and hydroxides dissolve in water to give strongly alkaline solutions, for example sodium hydroxide and potassium hydroxide. The word alkali comes from the Arabic for plant ashes, which were rich in potassium carbonate. All six elements, lithium to francium, share one outer s electron, which is why they behave so alike.

Why is lithium the strongest reducing agent even though it has the highest ionisation enthalpy?

Reducing power in water depends on the total energy to turn the solid metal into the hydrated ion. The tiny lithium ion releases 506 kJ per mole on hydration, far more than any other alkali ion. This outweighs its higher ionisation enthalpy, so the net cost is lowest and its electrode potential is the most negative, minus 3.04 volts.

Why is potassium lighter than sodium?

Density is mass divided by volume. Going from sodium to potassium, the atomic radius jumps from 186 to 227 pm, which increases the atomic volume by more than the gain in mass. So potassium, at 0.86 grams per cubic centimetre, is lighter than sodium at 0.97, and both float on water.

Why does a solution of sodium in liquid ammonia turn blue?

Sodium gives its electron to the ammonia, where it is trapped in a cavity surrounded by ammonia molecules. This ammoniated electron absorbs light at the red end of the visible spectrum, so the solution looks deep blue. The unpaired electron also makes the dilute solution paramagnetic and electrically conducting.

Why are alkali metals stored in kerosene?

They react quickly with oxygen, water vapour and carbon dioxide in air, so they must be kept away from it. Kerosene is an unreactive hydrocarbon that does not dissolve them. Lithium is so light that it floats on kerosene, so it is usually wrapped in paraffin wax or petroleum jelly instead.

Why do alkali metals give colours to a flame?

The heat of the flame excites the loosely held outer electron to a higher energy level. When it falls back, it emits light of a definite wavelength: lithium crimson red, sodium golden yellow, potassium lilac, rubidium red violet and caesium blue. The colours identify the metals in flame tests and flame photometry.

How are alkali metals tested in JEE Advanced?

JEE Advanced usually sets statement-based questions with one or more correct options: electrode potential versus reactivity with water, oxide formed on burning, colours and magnetism of liquid ammonia solutions, and density or mobility orders. Learn every trend together with its reason and its exception, because the wrong options are built from the exceptions.

Which alkali metal exceptions are most asked in JEE Advanced?

Five come up again and again in JEE Advanced: potassium is lighter than sodium, lithium has the most negative electrode potential but reacts slowest with water, sodium has the least negative potential, lithium alone forms a nitride and a normal oxide, and the lithium ion is the least mobile in water despite being the smallest.

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