Alkali metals (Group 1) — physical - chemical properties, anomaly of Li, diagonal Li-Mg

Physical Properties of Alkali Metals

1. Electronic Configuration

All alkali metals have one electron in the outermost s-orbital:

Li: [He] 2s¹
Na: [Ne] 3s¹
K: [Ar] 4s¹
Rb: [Kr] 5s¹
Cs: [Xe] 6s¹
Fr: [Rn] 7s¹

WHY this matters: This single valence electron determines ALL their properties—reactivity, bonding, metalic nature.

2. Atomic and Ionic Radii

Formula Trend Down the Group

$r_{\text{atom}} : \quad \text{Li} < \text{Na} < \text{K} < \text{Rb} < \text{Cs}$ $r_{\text{ion}} : \quad \text{Li}^+ < \text{Na}^+ < \text{K}^+ < \text{Rb}^+ < \text{Cs}^+$

Derivation from first principles: Each successive element adds a new electron shell. Lithium has 2 shells (K, L), sodium has 3 (K, L, M), potassium has 4, etc.

Effective nuclear charge felt by the outermost electron: $Z_{\text{eff}} = Z - \sigma$ where $Z$ = nuclear charge, $\sigma$ = shielding constant.

As we go down: $Z$ increases (+1 proton each step), BUT $\sigma$ increases MORE (full inner shells shield strongly). Net result: $Z_{\text{eff}}$ increases only slightly, but distance to nucleus increases dramatically. Atomic radius increases.

When they form M⁺ ions (lose the ns¹ electron), one shell is completely removed. Same trend persists.

WHY ionic radius is smaller than atomic radius:

Neutral atom: [Core] ns¹
After ionization: [Core] (the ns¹ shell is gone!)
Fewer electron-electron repulsions, and the remaining electrons are pulled closer by the unshielded nucleus.

3. Density

4. Ionization Enthalpy (I.E.)

Consequence: Cesium is the most reactive metal (easiest to lose electron), lithium the least reactive in Group 1.

5. Melting and Boiling Points

Alkali metals are soft (easily cut with knife) for the same reason!

6. Flame Colors

Chemical Properties of Alkali Metals

1. Reaction with Water

Worked example Worked Example: Sodium in Water

Given: 2.3 g of Na dropped in water. Find volume of H₂ at STP. Step 1: Write balanced equation $2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2$

Why this step? We need stoichiometry—mole ratio of Na to H₂.

Step 2: Find moles of Na $n_{\text{Na}} = \frac{2.3\text{ g}}{23\text{ g/mol}} = 0.1\text{ mol}$

Why? Moles connect mass to the balanced equation.

Step 3: Use stoichiometry From equation: 2 mol Na → 1 mol H₂ So: 0.1 mol Na → 0.05 mol H₂ Step 4: Volume at STP $V = n \times 22.4\text{ L/mol} = 0.05 \times 22.4 = 1.12\text{ L}$

Answer: 1.12 L of H₂ gas.

2. Reaction with Oxygen

Formula Oxygen Reactions (Complex!)

Lithium: Forms only normal oxide $4\text{Li} + \text{O}_2 \rightarrow 2\text{Li}_2\text{O}$
Sodium: Forms mainly peroxide $2\text{Na} + \text{O}_2 \rightarrow \text{Na}_2\text{O}_2$
K, Rb, Cs: Form superoxides $\text{K} + \text{O}_2 \rightarrow \text{KO}_2$ WHY this trend: Larger cations (K⁺, Rb⁺, Cs⁺) can stabilize larger anions (O₂⁻, the superoxide ion). Small Li⁺ can only stabilize small O²⁻.

This is a lattice energy argument: $U \propto \frac{z^+ z^-}{r^+ + r^-}$ For Li⁺ (small): $r^+$ is tiny → can pack tightly with O²⁻ → stable oxide. For Cs⁺ (large): $r^+$ is huge → needs large anion for stable lattice → superoxide.

3. Reaction with Halogens

4. Reaction with Hydrogen

5. Reducing Nature

Alkali metals are the strongest reducing agents in aqueous solution.

Formula Standard Reduction Potentials

$E^\circ: \quad \text{Li}^+/\text{Li} = -3.04\text{ V} \quad (\text{most negative!})$ $E^\circ: \quad \text{Na}^+/\text{Na} = -2.71\text{ V}$ $E^\circ: \quad \text{K}^+/\text{K} = -2.93\text{ V}$ More negative $E^\circ$ = stronger reducing agent (easier to donate electrons).

WAIT: Shouldn't Li be the weakest (highest I.E.)? Why is it the STRONGEST?

The key: In aqueous solution, we must account for hydration enthalpy!

Full thermodynamic cycle: $\text{M}(s) \rightarrow \text{M}(g) \quad [\Delta H_{\text{sub}}]$ $\text{M}(g) \rightarrow \text{M}^+(g) + e^- \quad [\text{I.E.}]$ $\text{M}^+(g) \xrightarrow{\text{H}_2\text{O}} \text{M}^+(aq) \quad [\Delta H_{\text{hyd}}]$ Net: $\Delta H = \Delta H_{\text{sub}} + \text{I.E.} + \Delta H_{\text{hyd}}$

Li⁺ is TINY → extremely high hydration enthalpy (water molecules stronglyolvate it). This OVERCOMPENSATES for its high I.E.

Result: Li is the strongest reducing agent in solution, even though it has the highest I.E.!

The Anomaly of Lithium

Anomalous Properties

Hardness: Li is the hardest alkali metal (still soft, but relatively harder).
- Why: Smallest size → strongest metalic bonds → harder to deform.
Melting & Boiling Point: Highest in the group.
- Why: Stronger metallic bonding (atoms packed closer).
Hydration Enthalpy: Exceptionally high.
- Why: Tiny Li⁺ (radius ≈ 76 pm) → huge charge density → water molecules strongly attracted.
Solubility of Salts: LiF is nearly insoluble; other Li salts are less soluble than corresponding Na, K salts.
- Why: High lattice energy for LiF (both ions tiny) cannot be overcome by hydration.
Covalent Character: Li compounds have more covalent character than other alkali compounds.
- Why: Fajans' rules! Small cation (Li⁺) polarizes the electron cloud of anions → covalent character.
- Example: LiCl is soluble in organic solvents (alcohol, acetone), but NaCl is not.
Deliquescence: LiCl, LiBr are strongly deliquescent (absorb moisture from air).
- Why: High hydration enthalpy → strongly attracts water molecules.
Reaction with N₂: Only Li forms nitride directly!
- $6\text{Li} + \text{N}_2 \rightarrow 2\text{Li}_3\text{N}$
- Why: Li⁺ is small enough to stabilize the small N³⁻ ion. Other alkali metals can't do this.
Decomposition of Compounds: Li₂CO₃, LiOH decompose on heating; other alkali carbonates/hydroxides don't.
- $\text{Li}_2\text{CO}_3 \xrightarrow{\Delta} \text{Li}_2\text{O} + \text{CO}_2$
- Why: Small Li⁺ → high polarizing power → destabilizes CO₃²⁻ → easier decomposition.

Diagonal Relationship: Lithium and Magnesium

Why Li and Mg are Similar

Similarities Between Li and Mg

Hardness: Both are harder than their respective group members.
Reaction with O₂: Both form only normal oxides (Li₂O, MgO), not peroxides/superoxides.
Reaction with N₂: Both form nitrides directly.
- $6\text{Li} + \text{N}_2 \rightarrow 2\text{Li}_3\text{N}$
- $3\text{Mg} + \text{N}_2 \rightarrow \text{Mg}_3\text{N}_2$
Carbonate/Hydroxide Decomposition: Both decompose on heating.
- $\text{Li}_2\text{CO}_3 \xrightarrow{\Delta} \text{Li}_2\text{O} + \text{CO}_2$
- $\text{MgCO}_3 \xrightarrow{\Delta} \text{MgO} + \text{CO}_2$
Halide Solubility: LiF and MgF₂ are sparingly soluble. LiCl and MgCl₂ are deliquescent.
Covalent Character: Both form compounds with covalent character (organolithium, Grignard reagents).
Nitrate Decomposition:
- $4\text{LiNO}_3 \xrightarrow{\Delta} 2\text{Li}_2\text{O} + 4\text{NO}_2 + \text{O}_2$
- $2\text{Mg(NO}_3)_2 \xrightarrow{\Delta} 2\text{MgO} + 4\text{NO}_2 + \text{O}_2$
- (Other alkali nitrates form nitrites: $2\text{NaNO}_3 \rightarrow 2\text{NaNO}_2 + \text{O}_2$ )

Worked example Worked Example: Comparing LiCl and NaCl

Question: Why is LiCl soluble in ethanol but NaCl is not?

Step 1: Identify the key difference LiCl has significant covalent character; NaCl is purely ionic.

Why this matters? Covalent compounds dissolve in nonpolar/less polar solvents; ionic compounds need highly polar solvents. Step 2: Apply Fajans' rules

Li⁺ is very small (high charge density)
Cl⁻ is relatively large (easily polarizable)
Li⁺ polarizes Cl⁻ electron cloud → partial covalent bond Why? Cations with high charge density distort anion electron clouds.

Step 3: Solubility principle "Like dissolves like"

Ethanol is less polar than water
LiCl (partial covalent) can interact with ethanol via dipole interactions
NaCl (purely ionic) cannot—it needs the strong polarity of water

Answer: LiCl's covalent character makes it soluble in ethanol.

Summary Table

Property	Trend Down Group 1	Li Anomaly
Atomic radius	Increases	Smallest
Ionization energy	Decreases	Highest
Reactivity with water	Increases	Least reactive
Melting point	Decreases	Highest
Metalic character	Increases	Least metallic
Hydration enthalpy	Decreases	Highest
Covalent character	Decreases	Highest
Reducing power (aq)	Li > K > Na > Rb > Cs	Li is strongest!

Recall Feynman: Explain to a 12-year-old

Imagine you have five friends: Li, Na, K, Rb, and Cs. They all have one toy they don't really want (that's the valence electron). Li is the smallest kid, holds his toy closest to himself, but when someone offers to trade (like water), he gets SO excited because he's tiny and people really like him (high hydration). So even though he holds his toy tighter than the others, he's actually the FIRST to give it away in a trade!

Cs is the biggest kid, barely holding onto his toy—it's so far from his body that it just falls off! He gives it away even without a good trade.

Now here's the funny part: Li is so small that he acts more like his diagonal neighbor Mg (who's in a different group) than his own family! It's like Li and Mg are both tiny kids who share toys the same way, even though Mg is technically from a different family. They both hate nitrogen (they grab it and make nitrides), they both only like normal oxygen toys (not the fancy peroxide ones), and their carbonates break easily when heated.

The lesson? Being small and highly charged (like Li⁺) makes you DIFFERENT from your family!

Connections 3.106-Periodic-trends-in-s-block — Why I.E. decreases down the group

3.1.08-Alkaline-earth-metals-Group-2 — Compare with Group 2 trends
3.1.09-Diagonal-relationship-general — Why it happens
4.2.03-Ionic-vs-covalent-character — Fajans' rules for Li compounds
2.5.04-Lattice-energy-Born-Haber-cycle — Why LiF is insoluble
5.3.02-Standard-reduction-potentials — Li's reducing power explained
6.1.05-Hydration-enthalpy — Why tiny ions have huge hydration energies

#flashcards/chemistry

What is the electronic configuration pattern of alkali metals? :: All alkali metals have one electron in the outermost s-orbital (ns¹ configuration). Example: Na is [Ne] 3s¹.

How does atomic radius change down Group 1 and why?

Atomic radius increases down the group (Li < Na < K < Rb < Cs) because each element adds new electron shell, increasing the distance from the nucleus. Shielding also increases, weakening the effective nuclear pull.

Why is potassium less dense than sodium?

Potassium has an unusually large jump in atomic radius (3rd to 4th shell). The volume increase outpaces the mass increase, causing density to drop temporarily. This is anomaly in the otherwise increasing density trend.

What is the ionization energy trend in Group 1?

Ionization energy decreases down the group (Li > Na > K > Rb > Cs). As atomic size increases, the outermost electron is farther from the nucleus and easier to remove despite increasing nuclear charge.

Why are alkali metals stored in kerosene?

Alkali metals react vigorously with oxygen and moisture in air. Kerosene is a nonpolar hydrocarbon that prevents contact with air and water, keeping the metals stable.

What products form when sodium reacts with water?

2Na + 2H₂O → 2NaOH + H₂. Sodium hydroxide (a strong base) and hydrogen gas are produced. The reaction is exothermic and can ignite the hydrogen.

Why does lithium form only normal oxide while potassium forms superoxide?

Small Li⁺ can only stabilize small O²⁻ anions (forming Li₂O). Large K⁺ can stabilize larger O₂⁻ superoxide anions (forming KO₂). This is due to lattice energy considerations—small cations need small anions for stable crystal structures.

Which alkali metal is the strongest reducing agent in aqueous solution and why?

Lithium is the strongest reducing agent (E° = -3.04 V) because its tiny Li⁺ ion has extremely high hydration enthalpy. The energy released by solvation overcompensates for its high ionization energy.

Name three anomalous properties of lithium.

(1) Highest melting and boiling points in Group 1; (2) Forms nitride (Li₃N) directly with N₂, unlike other alkali metals; (3) Its compounds (like LiCl) have significant covalent character and dissolve in organic solvents.

What is the diagonal relationship between Li and Mg?

Li (Group 1, Period 2) and Mg (Group 2, Period 3) have similar charge density (charge/radius ratio), giving them similar chemical properties: both form nitrides, decompose carbonates/hydroxides on heating, form only normal oxides, and have compounds with covalent character.

Why does LiF have low solubility while other alkali halides are highly soluble?

LiF has extremely high lattice energy because both Li⁺ and F⁻ are very small ions. The strong electrostatic attraction in the crystal lattice cannot be overcome by hydration enthalpy, making LiF sparingly soluble.

What flame color does sodium produce and why?

Sodium produces a yellow flame. When heated, the3s¹ electron jumps to a higher energy level; when it falls back, it emits yellow light (589 nm) corresponding to the specific energy gap in sodium.

Why do alkali metals have low melting points?

Alkali metals have weak metallic bonding because they have only one valence electron to contribute to the delocalized electron sea. As atomic size increases down the group, metalic bonding weakens further, and melting points decrease.

Write the reaction of lithium with nitrogen.

6Li + N₂ → 2Li₃N. Lithium nitride forms. This is unique—other alkali metals do not react directly with nitrogen at ordinary temperatures due to the stability of the N≡N triple bond and their larger ionic size.

Why does Li₂CO₃ decompose on heating but Na₂CO₃ does not?

Small Li⁺ has high polarizing power, which distorts the electron cloud of CO₃²⁻, weakening the C-O bonds and making decomposition easier: Li₂CO₃ →₂O + CO₂. Larger Na⁺ has lower polarizing power, so Na₂CO₃ is thermally stable.

Concept Map

Hinglish (regional understanding)

Intuition Hinglish mein samjho

Hinglish (regional understanding)

Intuition Hinglish mein samjho

Dekho beta, alkali metals (Group 1) ki sabse badi baat samajhna ho toh ek chiz yaad rakho — inke paas outermost shell mein sirf ek lonely electron hota hai, jo ns¹ configuration mein bahut door baitha hota hai nucleus se. Isko socho jaise kisi crowded room ke bilkul kinare pe khada banda — bilkul loosely attached, easily nikal sakta hai. Isiliye ye metals itne reactive hote hain ki water mein daalo toh explode kar jaate hain, aur inko kerosene mein store karna padta hai! Ye ek electron hi inki saari properties decide karta hai — reactivity, bonding, metallic nature, sab kuch.

Ab trend ki baat karein toh jaise-jaise hum group mein neeche jaate hain (Li → Na → K → Rb → Cs), naye shells add hote jaate hain, toh atomic radius badhta jaata hai aur outer electron nucleus se aur door hota jaata hai. Yahan ek important concept hai Z_eff = Z - σ — matlab nuclear charge toh badhta hai har step pe, par shielding usse zyada badhti hai, toh net mein outer electron ka pakad kamzor hota jaata hai. Isi wajah se ionization enthalpy neeche jaate-jaate ghatti hai, aur reactivity aur badhti jaati hai. Isiliye Cesium sabse reactive hai aur Lithium sabse kam. Density mein ek chhota sa twist hai — Potassium anomaly dikhata hai kyunki uska radius jump itna bada hota hai ki volume mass se zyada badh jaata hai, toh density temporarily gir jaati hai.

Ye sab why-it-matters isliye hai kyunki exam mein direct puchte hain "yeh trend kyun?" aur agar tumhe first principles se — Z, distance, aur shielding ka interplay — samajh aa gaya, toh tum har trend khud derive kar sakte ho, ratta maarne ki zaroorat hi nahi. Melting/boiling points bhi isi metallic bonding logic se aate hain — jitna chhota atom aur strong pull, utna strong bond, utna high M.P. Toh basically ek hi core idea — "lonely outer electron aur uski distance from nucleus" — se poora Group 1 ka behaviour unlock ho jaata hai. Isko pakad lo, baaki sab automatically fit ho jaayega!

Alkali metals (Group 1) — physical - chemical properties, anomaly of Li, diagonal Li-Mg

Physical Properties of Alkali Metals

1. Electronic Configuration

2. Atomic and Ionic Radii

3. Density

4. Ionization Enthalpy (I.E.)

5. Melting and Boiling Points

6. Flame Colors

Chemical Properties of Alkali Metals

1. Reaction with Water

2. Reaction with Oxygen

3. Reaction with Halogens

4. Reaction with Hydrogen

5. Reducing Nature

The Anomaly of Lithium

Anomalous Properties

Diagonal Relationship: Lithium and Magnesium

Why Li and Mg are Similar

Similarities Between Li and Mg

Summary Table

Connections 3.106-Periodic-trends-in-s-block — Why I.E. decreases down the group

Concept Map

Hinglish (regional understanding)

Hinglish (regional understanding)

Go deeper — visual, from zero

Test yourself — Hydrogen and s-Block

Connections