The s-block elements are located on the left side of the Periodic Table, encompassing Group 1 and Group 2. These are elements in which the last electron enters the outermost s-orbital. Since an s-orbital can hold a maximum of two electrons, there are two groups in the s-block.

Group 1 consists of Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), and Francium (Fr). They are known as **alkali metals** because they react with water to form strongly alkaline solutions (hydroxides).

Group 2 includes Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). With the exception of Beryllium, they are called **alkaline earth metals**. This is because their oxides and hydroxides are alkaline, and these metal oxides are found in the Earth's crust.

Sodium and Potassium are relatively abundant alkali metals, while Lithium, Rubidium, and Caesium are less abundant. Francium is extremely rare and radioactive. Among alkaline earth metals, Calcium and Magnesium are more abundant than Strontium, Barium, and the rare, radioactive Radium.

The general electronic configuration is [noble gas]ns$^1$ for alkali metals and [noble gas]ns$^2$ for alkaline earth metals.

The first element of each group (Lithium and Beryllium) displays properties different from other members of their group. They also show similarities to the second element of the next group (Lithium resembles Magnesium, and Beryllium resembles Aluminium). This is called the **diagonal relationship**, arising from similar ionic sizes and/or charge-to-radius ratios.

Sodium, Potassium, Magnesium, and Calcium ions are biologically important, playing roles in ion balance and nerve impulse conduction.

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Group 1 Elements: Alkali Metals

The alkali metals exhibit characteristic trends in their physical and chemical properties down the group.

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Electronic Configuration

All alkali metals have a valence shell electronic configuration of ns$^1$ (Table 10.1), where n is the period number. This single, loosely held s-electron makes them highly electropositive and reactive. They readily lose this electron to form stable $+1$ ions (M$^+$). Consequently, they are not found in their elemental form in nature.

Property	Lithium (Li)	Sodium (Na)	Potassium (K)	Rubidium (Rb)	Caesium (Cs)	Francium (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)
Electronic configuration	[He] 2s$^1$	[Ne] 3s$^1$	[Ar] 4s$^1$	[Kr] 5s$^1$	[Xe] 6s$^1$	[Rn] 7s$^1$
Ionization enthalpy / kJ mol$^{-1}$	520	496	419	403	376	~375
Hydration enthalpy/kJ mol$^{-1}$ (M$^+$)	–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	–
Standard potentials E$^{\tiny\mathbf{0}}$ / V for (M$^+$ / M)	–3.04	–2.714	–2.925	–2.930	–2.927	–2.927
Occurrence in lithosphere†	18*	2.27**	1.84**	78-12*	2-6*	~ 10$^{-18}$ *

*ppm (part per million), ** percentage by weight; † Lithosphere: The Earth’s outer layer: its crust and part of the upper mantle

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Atomic And Ionic Radii

Alkali metal atoms are the **largest** in their respective periods due to their low effective nuclear charge experienced by the outermost electron. Both atomic and ionic (M$^+$) radii **increase** down the group from Li to Cs as the number of electron shells increases. The M$^+$ ion is always smaller than the parent atom due to the loss of the outermost electron shell.

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Ionization Enthalpy

Alkali metals have **very low ionization enthalpies** (energy to remove an electron) because the single valence electron is loosely held and effectively shielded by inner electrons. Ionization enthalpy **decreases** down the group as atomic size increases and the outermost electron is further from the nucleus.

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Hydration Enthalpy

When alkali metal ions (M$^+$) are dissolved in water, they get hydrated by polar water molecules, releasing energy (hydration enthalpy). The degree of hydration and hydration enthalpy **decrease** with increasing ionic size down the group (smaller ions are more effectively hydrated). Li$^+$ is the most hydrated and has the highest hydration enthalpy, leading to many lithium salts being hydrated (e.g., LiCl·2H$_2$O).

Li$^+$ > Na$^+$ > K$^+$ > Rb$^+$ > Cs$^+$ (decreasing hydration enthalpy)

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Physical Properties

• Appearance: Silvery white, lustrous.
• Softness: Very soft (can be cut with a knife). Softness increases down the group.
• Density: Low density, generally increasing down the group. Potassium is less dense than sodium.
• Melting and Boiling Points: Low melting and boiling points due to weak metallic bonding (only one valence electron involved). Melting and boiling points generally decrease down the group.
• Flame Colouration: Alkali metals and their salts impart characteristic colours to an oxidising flame (Li: Crimson red, Na: Yellow, K: Violet, Rb: Red-violet, Cs: Blue). This is due to the excitation and de-excitation of valence electrons, emitting light in the visible spectrum. This property is used in flame tests and photometry for identification.
• Photoelectric Effect: These metals readily emit electrons when irradiated with light (photoelectric effect), especially Cs and K, making them useful in photoelectric cells.
• Electrical and Thermal Conductivity: Good conductors, typical of metals.

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Chemical Properties

Alkali metals are **highly reactive** due to their large size and low ionization enthalpies. Reactivity **increases** down the group.

• Reactivity towards Air: They tarnish in dry air due to oxide formation and react with moisture to form hydroxides. They burn vigorously in oxygen: * Li forms mainly **monoxide** ($\text{Li}_2\text{O}$). * Na forms primarily **peroxide** ($\text{Na}_2\text{O}_2$). * K, Rb, Cs form **superoxides** ($\text{MO}_2$).

  $\text{4Li} + \text{O}_2 \rightarrow \text{2Li}_2\text{O}$

  $\text{2Na} + \text{O}_2 \rightarrow \text{Na}_2\text{O}_2$

  $\text{M} + \text{O}_2 \rightarrow \text{MO}_2$ (M = K, Rb, Cs)

  Lithium is unique as it also reacts directly with nitrogen in the air to form **nitride** ($\text{Li}_3\text{N}$). Alkali metals are stored in kerosene to prevent reaction with air/water.

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• Reactivity towards Water: React vigorously with water to form hydroxides and dihydrogen gas. Reactivity increases down the group. Lithium reacts less vigorously than expected from its reduction potential due to its high hydration energy.

  $\text{2M} + \text{2H}_2\text{O} \rightarrow \text{2M}^+ + \text{2OH}^- + \text{H}_2$

  They also react with other proton donors like alcohols and ammonia.
• Reactivity towards Dihydrogen: React with H$_2$ at high temperatures to form ionic hydrides (M$^+$H$^-$).

  $\text{2M} + \text{H}_2 \rightarrow \text{2MH}$

• Reactivity towards Halogens: React vigorously with halogens to form ionic halides (M$^+$X$^-$). Lithium halides are somewhat covalent due to the small, highly polarising Li$^+$ ion distorting the electron cloud of the halide anion. Covalent character increases with the size of the anion (LiI is most covalent).

  $\text{2M} + \text{X}_2 \rightarrow \text{2MX}$

• Reducing Nature: They are strong reducing agents (easily lose electrons), with lithium being the strongest. Their reducing power is measured by standard electrode potential ($E^0$). Although Li has the highest ionization energy, its exceptionally high hydration enthalpy of Li$^+$ gives it the most negative $E^0$ value, making it the strongest reducing agent in aqueous solution.

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Uses

Alkali metals and their compounds have various uses:

• **Lithium:** Used in alloys (with lead for bearings, aluminium for aircraft, magnesium for armour plates), in thermonuclear reactions, and electrochemical cells.
• **Sodium:** Used in Na/Pb alloy (historically for petrol additives), as a coolant in nuclear reactors, and in producing chemicals.
• **Potassium:** Essential in biological systems, KCl is a fertilizer, KOH is used for soft soap and CO$_2$ absorption.
• **Caesium:** Used in photoelectric cells.

General Characteristics Of The Compounds Of The Alkali Metals

Most common compounds of alkali metals are predominantly **ionic** due to the large size and low ionization enthalpy of the metals. However, lithium compounds show more covalent character due to the small size and high polarising power of Li$^+$.

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Oxides And Hydroxides

Reaction with oxygen yields different types of oxides: Li (monoxide, Li$_2$O), Na (peroxide, Na$_2$O$_2$), K, Rb, Cs (superoxides, MO$_2$). Higher oxides (peroxides, superoxides) are more stable with larger cations due to lattice energy effects. All these oxides are basic and react with water to form hydroxides, which are strong bases.

$\text{M}_2\text{O} + \text{H}_2\text{O} \rightarrow \text{2M}^+ + \text{2OH}^-$

$\text{M}_2\text{O}_2 + \text{2H}_2\text{O} \rightarrow \text{2M}^+ + \text{2OH}^- + \text{H}_2\text{O}_2$

$\text{2MO}_2 + \text{2H}_2\text{O} \rightarrow \text{2M}^+ + \text{2OH}^- + \text{H}_2\text{O}_2 + \text{O}_2$

Pure oxides and peroxides are colourless, while superoxides are yellow/orange and paramagnetic (due to unpaired electrons in O$_2^-$ ion).

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Alkali metal hydroxides are strong bases due to their high solubility and complete dissociation in water.

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Halides

Alkali metal halides (MX) are typically high melting, colourless ionic solids. They are generally soluble in water, although LiF has low solubility due to high lattice energy and CsI has low solubility due to low hydration energy. Lithium halides show some covalent character due to Li$^+$ polarising the anion, which is most prominent in LiI due to the large size of the I$^-$ anion.

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Salts Of Oxo-Acids

Alkali metals form salts with oxo-acids (acids with oxygen in the acidic group, like sulfates, nitrates, carbonates). These salts are generally soluble in water and thermally stable. The stability of carbonates and hydrogencarbonates increases down the group as the cation size increases. Lithium carbonate is less stable and decomposes easily on heating due to the small Li$^+$ ion stabilising the oxide product over the carbonate.

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Anomalous Properties Of Lithium

Lithium, the first element in Group 1, shows some properties that are significantly different from the other alkali metals. This anomalous behaviour is primarily due to its **exceptionally small atomic and ionic size** and its **high polarising power** (high charge/radius ratio).

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Points Of Difference Between Lithium And Other Alkali Metals

1. Hardness and melting/boiling points: Li is much harder and has higher m.p./b.p.
2. Reactivity and reducing power: Li is the least reactive (among alkali metals) but the strongest reducing agent (in aqueous solution).
3. Reaction with air: Forms mainly monoxide (Li$_2$O) and nitride (Li$_3$N). Others form peroxide/superoxides and do not form nitride directly.
4. Hydrates: LiCl is deliquescent and forms a hydrate (LiCl·2H$_2$O). Other alkali metal chlorides are usually anhydrous.
5. Hydrogencarbonates: Lithium hydrogencarbonate does not exist as a solid.
6. Reaction with ethyne: Lithium does not form ethynide.
7. Nitrate decomposition: LiNO$_3$ decomposes to oxide (Li$_2$O). Others decompose to nitrite.
8. Solubility: LiF and Li$_2$O are less soluble in water. LiCl is soluble in organic solvents.

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Points Of Similarities Between Lithium And Magnesium

Lithium shows similarities to magnesium (diagonal relationship) due to similar atomic and ionic sizes (Li 152pm, Li$^+$ 76pm; Mg 160pm, Mg$^{2+}$ 72pm). Key similarities:

1. Hardness: Both are harder and lighter than other elements in their respective groups.
2. Reactivity with water: Both react slowly with water.
3. Solubility and decomposition of hydroxides: Both form less soluble hydroxides that decompose on heating.
4. Reaction with nitrogen: Both form nitrides ($\text{Li}_3\text{N}, \text{Mg}_3\text{N}_2$) by direct combination with N$_2$.
5. Oxides and peroxides: Their oxides (Li$_2$O, MgO) do not form superoxides with excess oxygen.
6. Carbonate decomposition: Their carbonates decompose easily on heating. Neither forms solid hydrogencarbonates.
7. Halide solubility: Both LiCl and MgCl$_2$ are soluble in ethanol.
8. Hydrates: Both form hydrated chlorides ($\text{LiCl} \cdot \text{2H}_2\text{O}, \text{MgCl}_2 \cdot \text{8H}_2\text{O}$).

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Some Important Compounds Of Sodium

Important sodium compounds include sodium carbonate ($\text{Na}_2\text{CO}_3$), sodium chloride (NaCl), sodium hydroxide (NaOH), and sodium hydrogencarbonate (NaHCO$_3$).

• Sodium Carbonate (Washing Soda, Na$_2$CO$_3$·10H$_2$O): Prepared industrially by the **Solvay process**, which uses NaCl, NH$_3$, and CaCO$_3$. It relies on the low solubility of NaHCO$_3$ which precipitates and is then heated to yield Na$_2$CO$_3$. Na$_2$CO$_3$ is a white solid, soluble in water (forming alkaline solution due to hydrolysis), used in glass, soap, paper, cleaning, and as a lab reagent.
• Sodium Chloride (NaCl): Obtained mainly by evaporation of sea water. Crude salt is purified by dissolving in water and precipitating pure NaCl by saturating with HCl gas (common ion effect). Used as table salt and in the production of Na$_2$O$_2$, NaOH, and Na$_2$CO$_3$.
• Sodium Hydroxide (Caustic Soda, NaOH): Prepared commercially by the **electrolysis of brine solution** in the Castner-Kellner cell. Sodium amalgam formed at the mercury cathode reacts with water to produce NaOH and H$_2$. NaOH is a white, deliquescent solid, a strong base, used in manufacturing soap, paper, textiles, petroleum refining, and as a lab reagent.
• Sodium Hydrogencarbonate (Baking Soda, NaHCO$_3$): Prepared by saturating a solution of Na$_2$CO$_3$ with CO$_2$. Decomposes on heating to produce CO$_2$ (used in baking and fire extinguishers), used as a mild antiseptic and antacid ingredient.

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Biological Importance Of Sodium And Potassium

Sodium and potassium ions are crucial for biological functions:

• **Sodium ions (Na$^+$):** Located primarily outside cells (blood plasma, interstitial fluid). Important for nerve signal transmission, water balance across cell membranes, and transport of sugars and amino acids into cells.
• **Potassium ions (K$^+$):** Most abundant cations inside cells. Activate many enzymes, involved in glucose oxidation to produce ATP. Participate with Na$^+$ in nerve signal transmission.

The **sodium-potassium pump** actively transports these ions across cell membranes against concentration gradients, a process vital for nerve function and consuming significant ATP energy.

Group 2 Elements : Alkaline Earth Metals

The alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) follow alkali metals in the periodic table. They are less reactive than alkali metals but more reactive than Group 13 elements. Beryllium is anomalous and shows diagonal relationship with aluminium.

Property	Beryllium (Be)	Magnesium (Mg)	Calcium (Ca)	Strontium (Sr)	Barium (Ba)	Radium (Ra)
Atomic number	4	12	20	38	56	88
Atomic mass (g mol$^{-1}$)	9.01	24.31	40.08	87.62	137.33	226.03
Electronic configuration	[He] 2s$^2$	[Ne] 3s$^2$	[Ar] 4s$^2$	[Kr] 5s$^2$	[Xe] 6s$^2$	[Rn] 7s$^2$
Ionization enthalpy (I) / kJ mol$^{-1}$	899	737	590	549	503	509
Ionization enthalpy (II) /kJ mol$^{-1}$	1757	1450	1145	1064	965	979
Hydration enthalpy (kJ/mol) (M$^{2+}$)	– 2494	– 1921	–1577	– 1443	– 1305	–
Metallic radius / pm	111	160	197	215	222	–
Ionic radius M$^{2+}$ / pm	31	72	100	118	135	148
m.p. / K	1560	924	1124	1062	1002	973
b.p / K	2745	1363	1767	1655	2078	(1973)
Density / g cm$^{-3}$	1.84	1.74	1.55	2.63	3.59	(5.5)
Standard potential E$^{\tiny\mathbf{0}}$ / V for (M$^{2+}$/ M)	–1.97	–2.36	–2.84	–2.89	– 2.92	–2.92
Occurrence in lithosphere	2*	2.76**	4.6**	384*	390 *	10–6*

*ppm (part per million); ** percentage by weight

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Electronic Configuration

Alkaline earth metals have a valence shell configuration of ns$^2$ (Table 10.2). They readily lose these two electrons to form stable $+2$ ions (M$^{2+}$).

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Atomic And Ionic Radii

Atomic and ionic radii are smaller than corresponding alkali metals due to higher nuclear charge. Radii **increase** down the group.

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Ionization Enthalpies

Ionization enthalpies are low but higher than corresponding alkali metals (due to smaller size and higher effective nuclear charge). First ionization enthalpy > second ionization enthalpy. Both **decrease** down the group.

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Hydration Enthalpies

Hydration enthalpies are higher than alkali metals (due to higher charge and smaller size). They **decrease** down the group with increasing ionic size. Alkaline earth metal compounds are generally more hydrated than alkali metal compounds (e.g., MgCl$_2$·6H$_2$O vs NaCl).

Be$^{2+}$ > Mg$^{2+}$ > Ca$^{2+}$ > Sr$^{2+}$ > Ba$^{2+}$ (decreasing hydration enthalpy)

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Physical Properties

• Appearance: Silvery white, lustrous, relatively soft (harder than alkali metals).
• Density: Higher density than alkali metals (generally increasing down the group).
• Melting and Boiling Points: Higher than corresponding alkali metals due to stronger metallic bonding (two valence electrons). No systematic trend in m.p./b.p.
• Flame Colouration: Ca (brick red), Sr (crimson), Ba (apple green) impart characteristic colours due to electron excitation. Be and Mg do not impart colour as their electrons are too strongly bound.
• Electrical and Thermal Conductivity: Good conductors.

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Chemical Properties

Alkaline earth metals are less reactive than alkali metals. Reactivity **increases** down the group.

• Reactivity towards Air and Water: Be and Mg are inert due to oxide film. Powdered Be burns in air to BeO and Be$_3$N$_2$. Mg burns brightly to MgO and Mg$_3$N$_2$. Heavier elements are more reactive, attacking air and water (even cold) to form hydroxides and H$_2$.

  $\text{M} + \text{2H}_2\text{O} \rightarrow \text{M(OH)}_2 + \text{H}_2$ (M = Ca, Sr, Ba)

• Reactivity towards Halogens: Combine at high temperatures to form halides (MX$_2$). Be halides are covalent; others are ionic.

  $\text{M} + \text{X}_2 \rightarrow \text{MX}_2$

• Reactivity towards Hydrogen: Combine with H$_2$ upon heating to form hydrides (MH$_2$), except Be (BeH$_2$ is made indirectly).

  $\text{M} + \text{H}_2 \rightarrow \text{MH}_2$ (M = Ca, Sr, Ba)

• Reactivity towards Acids: React readily with acids, liberating H$_2$.

  $\text{M} + \text{2HCl} \rightarrow \text{MCl}_2 + \text{H}_2$

• Reducing Nature: Strong reducing agents (due to low ionization enthalpies), but weaker than corresponding alkali metals. Reducing power increases down the group (more negative $E^0$).
• Solutions in Liquid Ammonia: Dissolve in liquid ammonia to give blue-black solutions of ammoniated ions and electrons.

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Uses

Alkaline earth metals and their compounds have various uses:

• **Beryllium:** In alloys (with Cu for springs), for X-ray tube windows.
• **Magnesium:** In light alloys (for aircraft), in flash powders, as an antacid (milk of magnesia), in toothpaste.
• **Calcium:** In metallurgy, as a getter to remove air from vacuum tubes.
• **Barium:** In metallurgy, as a getter. Barium sulfate is used as a contrast agent in X-rays.
• **Radium:** Radium salts used in radiotherapy for cancer treatment.

General Characteristics Of Compounds Of The Alkaline Earth Metals

Alkaline earth metals predominantly form ionic compounds with a $+2$ oxidation state. These compounds are generally less ionic than corresponding alkali metal compounds. Compounds of Be and Mg show more covalent character.

• **Oxides (MO):** Formed by burning metals in oxygen. BeO is amphoteric; others (MgO, CaO, SrO, BaO) are basic, becoming more basic down the group. They react with water to form sparingly soluble hydroxides.
• **Hydroxides (M(OH)$_2$):** Formed by reacting oxides with water. Solubility, thermal stability, and basic character increase down the group. They are weaker bases than alkali metal hydroxides. Be(OH)$_2$ is amphoteric.
• **Halides (MX$_2$):** Formed by reacting metals with halogens or oxides/carbonates with hydrohalic acids. Be halides are covalent and soluble in organic solvents; others are ionic. Halides show decreasing tendency to form hydrates down the group. Fluorides are less soluble than chlorides due to high lattice energies.
• **Salts of Oxoacids (Carbonates, Sulfates, Nitrates):** * Carbonates (MCO$_3$): Insoluble in water. Thermal stability increases down the group (BeCO$_3$ least stable). Decompose to MO and CO$_2$ on heating. * Sulphates (MSO$_4$): White solids, stable to heat. Solubility decreases down the group (BeSO$_4$, MgSO$_4$ soluble; CaSO$_4$, SrSO$_4$, BaSO$_4$ sparingly soluble). Solubility is determined by the balance of hydration enthalpy (high for small ions) and lattice enthalpy. * Nitrates (M(NO$_3$)$_2$): Soluble in water. Form hydrates (tendency decreases down the group). Decompose on heating to MO, NO$_2$, and O$_2$.

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Anomalous Behaviour Of Beryllium

Beryllium, the first element in Group 2, exhibits anomalous behaviour compared to other alkaline earth metals. This is due to its **exceptionally small atomic and ionic size** and **high ionisation enthalpy**, leading to its compounds having significant **covalent character**. Its properties also show a **diagonal relationship** with aluminium in Group 13.

Points of difference from other alkaline earth metals:

• Small size and high ionisation enthalpy lead to covalent compounds and easy hydrolysis.
• Does not exhibit coordination number greater than 4 (only 4 valence orbitals). Others can use d-orbitals for coordination number 6.
• Oxide (BeO) and hydroxide (Be(OH)$_2$) are **amphoteric**. Others are basic.

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Diagonal Relationship Between Beryllium And Aluminium

Be resembles Al due to similar ionic sizes (Be$^{2+}$ 31pm, Al$^{3+}$ 53.5pm) and charge/radius ratios. Similarities:

• Both are not readily attacked by acids due to oxide film formation.
• Their hydroxides are amphoteric and dissolve in excess alkali to form beryllate ([Be(OH)$_4$]$^{2-}$) and aluminate ([Al(OH)$_4$]$^-$ or [Al(H$_2$O)$_2$(OH)$_4$]$^-$) ions.
• Their chlorides (BeCl$_2$, AlCl$_3$) are covalent, soluble in organic solvents, have bridged structures in the vapour phase, and are strong Lewis acids (used as Friedel-Crafts catalysts).
• Both ions form complexes (e.g., BeF$_4^{2-}$, AlF$_6^{3-}$).

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Some Important Compounds Of Calcium

Calcium compounds like calcium oxide (quick lime), calcium hydroxide (slaked lime), calcium carbonate (limestone), calcium sulphate (gypsum, plaster of Paris), and cement are industrially important.

• Calcium Oxide (Quick Lime, CaO): Prepared by heating limestone (CaCO$_3$) in a kiln. White amorphous solid, high melting point. Reacts with water (slaking) to form Ca(OH)$_2$. Basic oxide, reacts with acidic oxides (SiO$_2$, P$_4$O$_{10}$). Used in cement, as an alkali, in sugar purification.
• Calcium Hydroxide (Slaked Lime, Ca(OH)$_2$): Prepared by adding water to CaO. White amorphous powder, sparingly soluble in water. Aqueous solution is lime water, suspension is milk of lime. Reacts with CO$_2$ to form CaCO$_3$ (turns lime water milky). Reacts with chlorine to form bleaching powder. Used in mortar, white wash, glass making, tanning, bleaching powder production, sugar purification.
• Calcium Carbonate (CaCO$_3$): Occurs as limestone, chalk, marble. Prepared by reacting Ca(OH)$_2$ with CO$_2$ or CaCl$_2$ with Na$_2$CO$_3$. White powder, almost insoluble in water. Decomposes on heating to CaO and CO$_2$. Reacts with dilute acids to liberate CO$_2$. Used as building material (marble), for quick lime production, as flux in metallurgy, in paper, toothpaste, cosmetics.
• Calcium Sulphate (Plaster of Paris, CaSO$_4$·½H$_2$O): Hemihydrate of calcium sulfate. Obtained by heating gypsum (CaSO$_4$·2H$_2$O) to 393 K. Sets with water to a hard solid (forms gypsum). Used in building, as plasters, for immobilising fractures, in dentistry, for casts.
• Cement (Portland cement): Important building material. Made by strong heating of limestone and clay to form clinker, then mixed with gypsum (2-3%). Main ingredients are dicalcium silicate, tricalcium silicate, tricalcium aluminate. Setting is due to hydration and rearrangement of constituents. Used in concrete, plastering, construction.

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Biological Importance Of Magnesium And Calcium

Magnesium and calcium ions are vital in biological systems:

• **Magnesium ions (Mg$^{2+}$):** Cofactor for all enzymes that use ATP in phosphate transfer (energy metabolism). Component of chlorophyll (light absorption in plants).
• **Calcium ions (Ca$^{2+}$):** Major component of bones and teeth ($\approx 99\%$ of body calcium). Essential for neuromuscular function, nerve transmission, cell membrane integrity, and blood coagulation. Blood calcium levels are tightly regulated by hormones.

Calcium is continuously exchanged between bone and plasma, highlighting the dynamic nature of bone tissue.

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