Periodicity and Group Trends

Posted on May 5, 2024 in Chemistry

Positive and Negative Ions

Positive ions are smaller as they lose electrons, forming a cation and losing the outer shell with the largest radius. Negative ions are bigger as they gain an electron, forming an anion with an increase in electron-electron repulsion, causing an expansion of the charge cloud.

Metals

Metals react by losing electrons. Metallic character increases down Group 1 and decreases across Period 2.

Carbon and Lead

Carbon in diamond forms a giant covalent structure where each carbon atom is bonded to four others. All four outer electrons are involved in bonding, giving a tetrahedral structure.

Lead (Pb) forms a giant metallic structure with an array of metal ions surrounded by a sea of delocalized electrons. The bonding force is the attraction between positive ions and delocalized electrons. Carbon at the top of Group 4 is a non-metal, while lead at the bottom is a metal. Metal oxides are basic and ionic (blue/pH 13), while non-metal oxides are acidic and covalent (red/pH 1-6).

Amphoteric Behavior

Aluminum (Al³⁺) doesn’t affect pH and is insoluble when added to water. As a base: Al₂O₃ + 3H₂SO₄ [6H⁺] > Al₂(SO₄)₃ [2Al³⁺] + 3H₂O. As an acid: Al₂O₃ + 3H₂O + 2NaOH [2OH^–] > 2NaAl(OH)₄ [2Al(OH)₄^–] (aluminate ion).

Beryllium (Be²⁺) has a high charge density, which polarizes the anion, creating a high degree of covalent character when forming salts. It has high electronegativity for a metal but is insufficient for ionic bonding with a non-metal. As a base: Be(OH)₂ + 2HCl > BeCl₂ + 2H₂O. As an acid: Be(OH)₂ + 2NaOH > Na₂Be(OH)₄.

Chlorides and Water

Ionic chlorides dissolve in water to give a neutral solution: NaCl(s) > Na⁺(aq) + Cl^–(aq). Ions are hydrated, and the lattice is broken. Covalent chlorides hydrolyze with water: AlCl₃(s) > Al³⁺ + 3Cl^–. The aluminum ion has a high charge density, attracting water molecules and forming dative bonds: [Al(H₂O)₆]³⁺ + aq > [Al(H₂O)₅(OH)]²⁺ + H⁺. The hydrated ion is acidic as the high charge density of Al³⁺ attracts electrons from the OH bond, releasing H⁺ to form an acidic solution. Further protonation: [Al(H₂O)₄(OH)₂]⁺ + 3H⁺.

Amphoteric Behavior of Aluminum

If NaOH is added to a solution of Al³⁺, the OH^– will remove H⁺ from the hydrated Al³⁺ ion to form Al(OH)₃: [Al(H₂O)₆]³⁺ + 3OH^– > [Al(H₂O)₃(OH)₃] + 3H₂O (a white precipitate). In excess NaOH: [Al(OH)₄]^– + 3H₂O. This can be reversed by adding acid to the sodium aluminate ion. Al(OH)₃ is an amphoteric hydroxide and reacts with dilute acids to form salts and warm Na/KOH to form an aluminate ion: 2Al + 6HCl > 2AlCl₃ + 3H₂ (ionic); 2Al + 6H⁺ > 2Al³⁺ + 3H₂. Amphoteric oxide Al₂O₃ with acid: Al₂O₃ + 6HCl > 2AlCl₃ + 3H₂O. With base as warm excess NaOH: Al₂O₃ + 2NaOH > 2NaAl(OH)₄.

Amphoteric Behavior of Lead

(Lead nitrate) Pb²⁺ + 2OH^– (NaOH) > Pb(OH)₂ (white precipitate). With excess NaOH: [Pb(OH)₄]^2- (plumbate ion) that redissolves to a colorless solution. Addition of acid to the ion reprecipitates Pb(OH)₂, which dissolves in more acid > Pb²⁺(aq). Amphoteric oxide PbO with acid: PbO + 2HCl > PbCl₂ + H₂O; PbO + 4HCl > PCl₂ + Cl₂ + 2H₂O. With base: PbO + 2OH^– + H₂O > Pb(OH)₄^2-.

Lead Oxide (PbO)

Lead sulfate/chloride are sparingly soluble in cold water, so lead nitrate is a better choice. As a base: PbO + 2HNO₃ > Pb(HNO₃)₂ + H₂O. On heating with NaOH, forming a colorless solution. As an acid: PbO + 2NaOH + H₂O > Na₂Pb(OH)₄. Pb²⁺ with NaOH: Pb²⁺ + 2OH^– > Pb(OH)₂. This dissolves in excess NaOH: [Pb(OH)₄]^2-. Pb²⁺ with Cl^– gives a white precipitate of PbCl₂, which is more soluble in hot water than cold and soluble in concentrated HCl: Pb(NO₃)₂ + 2NaCl > PbCl₂ + 2NaNO₃. Pb²⁺ with I^– gives a bright yellow precipitate of PbI₂, but with excess iodide ions, it redissolves and is soluble in hot water, giving a colorless solution.

Inert Pair Effect

Lower valences (+2) become more stable going down Groups 3, 4, and 5, with an increase in the tendency for the s² pair not to be used in bonding as the energy to form +4 ions is so large. Oxidation states of +4 are where all outer electrons are involved in covalent bonding, e.g., CO₂, SiO₂.

Tin and Lead +2

In Sn and Pb²⁺, two of the four outer shell electrons are not involved in bonding. This is due to the inert pair effect – an increase in the stability of the +2 state with increasing atomic number (down the group). So Pb²⁺ is more stable than C²⁺. This also applies to Groups 3 and 4. Ge²⁺ exists, but +4 is more stable. Sn²⁺ and Sn⁴⁺ have similar stabilities. Pb²⁺ is more stable than Pb⁴⁺. C⁴⁺ is more stable than C²⁺.

Phosphorus Valencies

Phosphorus can have valencies of 3 and 5 as it can promote an electron from the 3s to the 3d orbital, creating five unpaired electrons available for bonding, thus expanding its octet (four electron pairs).

Group 3 Elements

Group 3 elements have the electron configuration ns² np¹, with three electrons in the outer shell. They can form three covalent bonds, leaving an electron-deficient compound. They tend to form dative bonds. They can form three bonds with Group 5 elements that donate a lone pair to form a stable octet with a trigonal planar shape (120°). AlCl₃ forms a dimer, Al₂Cl₆, that dissociates above 200°C. NH₃.BF₃ is tetrahedral (109.5°). Cl₂ + AlCl₃ > AlCl₄^– + Cl⁺ (four bonds, one dative, three covalent).

Ionic Liquids

Ionic liquids are liquid over a wide temperature range. Large organic chlorocompounds (R-Cl) react with AlCl₃, where the organic Cl forms a dative bond to AlCl₃, followed by ionization: R⁺AlCl₄^–.

Advantages of ionic liquids:

• Easier to separate from catalysts and products than conventional VOCs.
• Recyclable.
• Organic products are immiscible in ionic liquids, so they can be separated and tapped off.
• Low vapor pressure, non-volatile, so they don’t evaporate on use.

Boron Nitride (BN)

BN is a giant covalent structure with the same number of electrons as graphite and diamond, making them isoelectronic. BN forms three bonds with an unbounded p-orbital. It has hardness, chemical inertness, a high melting temperature, and semiconductor properties. It is used as a lubricant, wear-resistant coating, and in nanotubes for wire sleeving. Nitrogen is more electronegative than boron due to its lone pair, and it is not possible to form pi bonds between N and B as their energies are too dissimilar.

BN and Carbon

BN and C are isoelectronic and have the same hexagonal layered structure. Both are lubricants with weak van der Waals forces between layers. Carbon is an electrical conductor, but BN is an insulator at room temperature. In carbon, the delocalization of electrons due to unbounded p-orbitals allows the conduction of electricity. Unlike carbon, in BN, each N has a full unbounded p-orbital, whereas B has an empty unbounded p-orbital. Nitrogen is more electronegative than boron, so electron density is not spread evenly. BN is non-conducting due to N being more electronegative. Carbon has one electron per atom for delocalization, B has none, and N has a lone pair that would have to split.

Hexagonal BN: An insulator but oxidizes at high temperatures. Chemically inert, it can be used in reaction vessels. BN nanotubes packed with fullerene spheres and exposed to an intense beam of electrons create an insulated conducting wire. The h-BN forms an insulating sleeve around a conducting carbon nanotube wire (wire sleeving).

Cubic BN: The second hardest material with high thermal conductivity and chemical inertness. It is wear-resistant and used as cutting tools and supports for catalysts. It is used in mounting high-power electronic components due to its high thermal conductivity for heat dissipation.

Carbon Monoxide (CO)

CO is a strong reducing agent and readily loses electrons to form C⁺⁴. It is used in reducing metal ores: CuO + CO > Cu + CO₂.

Lead Dioxide (PbO₂)

PbO₂ is a good oxidizing agent. With heating with concentrated HCl: PbO₂ + 4HCl > PbCl₂ + Cl₂ + H₂O.

Properties of Carbon Dioxide (CO₂)

CO₂ is a colorless gas readily soluble in water, forming an acidic solution: CO₂ + H₂O > H₂CO₃. When passed into excess NaOH: CO₂ + 2NaOH > Na₂CO₃ + H₂O. But if CO₂ is in excess: CO₂ + NaOH > NaHCO₃.

Limewater Test

Limewater is a saturated solution of Ca(OH)₂ that is slightly soluble in water. On passing CO₂, a white precipitate of insoluble CaCO₃ forms: CO₂ + Ca(OH)₂ > CaCO₃ + H₂O. But with excess CO₂, a colorless solution of Ca²⁺ and HCO₃^– ions forms: 2CO₂ + Ca(OH)₂ > Ca(HCO₃)₂. Decomposition of HCO₃^– > CaCO₃ + H₂O (stalactites).

Chlorine with Sodium Hydroxide

With cold dilute NaOH: Cl₂ + NaOH > NaCl + NaClO + H₂O; Cl₂ + 2OH^– > Cl^– + ClO^– + H₂O. This is disproportionation as Cl₂ is both oxidized and reduced. With warm concentrated NaOH: 3Cl₂ + NaOH > NaClO₃ + 5NaCl + 3H₂O; 3Cl₂ + 6OH^– > ClO₃^– + 5Cl^– + 3H₂O (also disproportionation).

Chlorine Bubbled in Water

Cl₂ + H₂O > HClO + HCl. The chlorate ion acts as a disinfectant and bleach, destroying microorganisms. It is used to sterilize water supplies. Sodium chlorate is a weed killer and is a powerful oxidizing agent, potentially explosive with organic material.

Concentrated H₂SO₄ on Sodium Halides

NaCl + conc H₂SO₄ > white misty fumes of HCl. HCl isn’t a strong enough reducing agent to affect concentrated H₂SO₄, so no Cl₂ is produced.

NaBr > white mist fumes of HBr + orange-brown fumes of Br₂ & colorless gas of SO₂ as HBr is a strong enough reducing agent to react with H₂SO₄, reducing S from +6 in H₂SO₄ to +4 in SO₂.

NaI > on warming gives purple fumes of I₂, white misty fumes of HI, colorless SO₂ gas, yellow solid of S, and a bad egg smell of H₂S.

F₂ > pale yellow gas; Cl₂ > greenish-yellow gas; Br₂ > red-brown liquid; I₂ > lustrous grey-black solid.

Halogens

Fluorine: Added to drinking water and toothpaste to prevent dental caries. PTFE (polytetrafluoroethylene) is resistant to chemical action and is used in the surface of frying pans.

Chlorine: Bleaching agent, weed killers, antiseptic (TCP).

Iodine: KI used in the treatment of goiter.

Ionization Energies

Group 3 elements have lower ionization energies than Group 2 elements as their outer electrons are now in p-orbitals, which have higher energy and are easier to remove despite the increase in nuclear charge. Group 2 outer electrons are in lower energy s-orbitals with stable shells.

Group 6 elements have lower ionization energies than Group 5 elements as they have paired electrons in their p-orbitals, which repel, making them easier to remove. Group 5 elements have more stable half-filled p-orbitals with no electron pair repulsion.