Inorganic Chemistry

The arrangement of elements in the periodic table according to block classification (s, p, d, f blocks) reflects the electron configuration and chemical properties. Elements are grouped by the subshell being filled: s-block elements (Groups 1-2) fill s orbitals, p-block elements (Groups 13-18) fill p orbitals, d-block transition metals fill d orbitals, and f-block lanthanides/actinides fill f orbi

Inorganic Chemistry

Inorganic Chemistry

Atomic radius decreases across a period from left to right (e.g., Na → Cl). This occurs because the nuclear charge increases while electrons are added to the same shell. Although more electrons are added, they are all in the same main shell, so the effective nuclear charge pulling on the outermost electrons increases significantly, drawing them closer to the nucleus and reducing the atomic radius.

Inorganic Chemistry

Inorganic Chemistry

First ionisation energy generally increases across a period (e.g., Na → Ar). This is because as atomic radius decreases and nuclear charge increases, the valence electrons are held more strongly by the nucleus. However, there are two notable exceptions: a small decrease from Group 2 to Group 13, and a small decrease from Group 15 to Group 16, due to electron pairing effects in p orbitals. First i

Inorganic Chemistry

Inorganic Chemistry

Melting points increase from Na to Si, then decrease sharply from P to Ar. Sodium to silicon show increasing melting points due to increasingly strong metallic/covalent bonding. Phosphorus, sulfur, chlorine, and argon show much lower melting points because they are molecular substances held together by weak intermolecular forces, not strong metallic or covalent bonds. Period 3 demonstrates a tran

Inorganic Chemistry

Inorganic Chemistry

Period 3 oxides show a transition from basic to amphoteric to acidic character. Na₂O and MgO are basic solid ionic oxides; Al₂O₃ is amphoteric (both basic and acidic properties); SiO₂ is acidic and covalent; P₂O₅ and Cl₂O₇ are acidic covalent oxides. This progression reflects the increasing electronegativity of the elements and the transition from ionic to covalent bonding. The nature of oxides c

Inorganic Chemistry

Inorganic Chemistry

Electronegativity increases across a period (left to right) and decreases down a group. This reflects how tightly atoms pull on electron density in a covalent bond. Elements with high electronegativity (like F and Cl) are better at attracting electron density, while elements with low electronegativity (like Na and K) are poor at attracting electrons. Electronegativity trends directly reflect atom

Inorganic Chemistry

Inorganic Chemistry

Group 2 elements (Be, Mg, Ca, Sr, Ba) are called alkaline earth metals (except beryllium which has different properties). They have the electron configuration [noble gas]s² and form +2 cations in ionic compounds. They are highly reactive, especially the lower members, and their reactivity increases down the group. They produce basic hydroxides and oxides. Group 2 elements are all metals with two

Inorganic Chemistry

Inorganic Chemistry

Reactivity of Group 2 metals increases down the group (Be << Mg < Ca < Sr < Ba). This is because atomic radius increases and first ionisation energy decreases down the group, making it progressively easier to lose the two valence 2s electrons and form M²⁺ cations. Barium is highly reactive while beryllium is relatively unreactive. As you move down Group 2, two factors combine to increase reactivi

Inorganic Chemistry

Inorganic Chemistry

Group 2 elements react vigorously with oxygen to form white or cream-colored basic oxides (MgO, CaO, SrO, BaO). The product is a metal oxide with the formula M₂O₂ or MO depending on conditions. These reactions are vigorous and exothermic, releasing significant heat energy. Calcium and strontium can form peroxides (M₂O₂ or MO₂) if burned in excess oxygen. Group 2 metals are reducing agents that re

Inorganic Chemistry

Inorganic Chemistry

Solubility of Group 2 hydroxides M(OH)₂ increases down the group: Mg(OH)₂ is only slightly soluble, Ca(OH)₂ is sparingly soluble, Sr(OH)₂ and Ba(OH)₂ are more soluble. This trend is unusual because lattice energy decreases down the group (usually decreases solubility), but hydration energy of smaller cations (like Mg²⁺) decreases faster, making dissolution unfavorable. Ba(OH)₂ is soluble enough to

Inorganic Chemistry

Inorganic Chemistry

Solubility of Group 2 sulfates M(SO₄) decreases down the group (opposite to hydroxides): MgSO₄ is very soluble, CaSO₄ is sparingly soluble (forms scale), SrSO₄ is insoluble, BaSO₄ is extremely insoluble (Ksp ≈ 1.1 × 10⁻¹⁰). This trend occurs because lattice energy increases faster than hydration energy decreases, making dissolution progressively less favorable. BaSO₄'s low solubility is used analy

Inorganic Chemistry

Inorganic Chemistry

Group 2 carbonates decompose on heating to form the metal oxide and carbon dioxide (MCO₃ → MO + CO₂). Thermal stability increases down the group—BeCO₃ is very unstable, while BaCO₃ is very stable and requires extremely high temperatures to decompose. The temperature required increases from ~200°C for BeCO₃ to >1000°C for BaCO₃. Thermal decomposition of Group 2 carbonates demonstrates the importan

Inorganic Chemistry

Inorganic Chemistry

Group 2 compounds have numerous industrial and practical applications. Calcium oxide (quicklime, CaO) is used in cement production and steel manufacture. Calcium carbonate is used as a building material and to reduce acidity. Barium sulfate is used as an X-ray contrast medium and in drilling fluids. Magnesium compounds are used in medicine and as buffers. The applications of Group 2 compounds ref

Inorganic Chemistry

Inorganic Chemistry

Halogens (Group 17: F, Cl, Br, I) are non-metallic elements with the electron configuration [noble gas]ns²np⁵, possessing seven valence electrons. They exist as diatomic molecules (X₂) in elemental form and readily gain one electron to form -1 halide ions (X⁻). Halogens are highly reactive due to their high electronegativity and strong tendency to complete their valence octet. The halogens are on

Inorganic Chemistry

Inorganic Chemistry

Halogens show a trend in physical properties down the group. Fluorine is a pale yellow gas, chlorine is a yellow-green gas, bromine is a brown liquid/vapor, and iodine is a dark purple/gray solid. Boiling points increase (F₂ -188°C, Cl₂ -35°C, Br₂ 59°C, I₂ 184°C), and relative atomic mass increases. These changes reflect increasing intermolecular forces (van der Waals) as the molecules become larg

Inorganic Chemistry

Inorganic Chemistry

Electronegativity decreases down the group (F 3.98 > Cl 3.16 > Br 2.96 > I 2.66), and reactivity also decreases (F₂ > Cl₂ > Br₂ > I₂). Fluorine is by far the most reactive halogen and the most electronegative element. The decrease in reactivity is due to increasing atomic radius and the greater difficulty in breaking the X-X bond and accepting electrons as you move down the group. Halogen reactiv

Inorganic Chemistry

Inorganic Chemistry

A more reactive (more electronegative) halogen displaces a less reactive halogen from its salt. For example, Cl₂ displaces Br⁻ from bromide salts, and Br₂ displaces I⁻ from iodide salts. These are redox reactions where the more reactive halogen is reduced (gains electrons) and the halide ion is oxidized (loses electrons). No displacement occurs if a more reactive halide is added to a less reactive

Inorganic Chemistry

Inorganic Chemistry

Halide ions react differently with concentrated sulfuric acid depending on their reactivity. Chloride ions are oxidized to Cl₂ gas, but HCl gas is produced first (making white fumes). Bromide ions are oxidized to Br₂ (brown solution) because H₂SO₄ is a strong enough oxidizing agent. Iodide ions are oxidized to I₂ (purple solution), and some iodide may be further oxidized to I₂ or even sulfur may b

Inorganic Chemistry

Inorganic Chemistry

Silver nitrate solution forms characteristic colored precipitates with halide ions: white precipitate with Cl⁻ (AgCl), cream/pale yellow with Br⁻ (AgBr), and yellow with I⁻ (AgI). The precipitate appearance allows identification of which halide is present. The test can be confirmed by adding dilute ammonia solution: AgCl and AgBr dissolve (forming complex ions), but AgI does not dissolve. This is

Inorganic Chemistry

Inorganic Chemistry

Disproportionation is when an element is simultaneously oxidized and reduced in the same reaction. Chlorine disproportionates in water to form hypochlorous acid (HClO) and hydrochloric acid (HCl): Cl₂ + H₂O ⇌ HCl + HClO. In the reaction, some Cl atoms are reduced to -1 (HCl) while others are oxidized to +1 (HClO). This equilibrium can be shifted by temperature and alkali. Disproportionation is a

Inorganic Chemistry

Inorganic Chemistry

Chlorine has numerous important applications: disinfection of water supplies (killing bacteria and viruses), production of bleach (hypochlorite), manufacture of chlorinated compounds (PVC, pesticides), and as an oxidizing agent in industry. Chlorates (especially sodium chlorate, NaClO₃) are used as oxidizing agents in explosives, match heads, and bleaching agents. Hypochlorite (ClO⁻) is the active

Inorganic Chemistry

Inorganic Chemistry

Period 3 elements react with oxygen to form oxides. Sodium and magnesium burn vigorously to form basic oxides (Na₂O, MgO). Aluminum burns to form Al₂O₃. Silicon reacts with oxygen at high temperatures to form SiO₂. Phosphorus burns to form P₄O₁₀ (or P₄O₆ depending on oxygen availability). Sulfur burns in oxygen to form SO₂. Chlorine does not burn in oxygen but can form some oxides under specific c

Inorganic Chemistry

Inorganic Chemistry

Period 3 oxides show different structures reflecting the bonding type: Na₂O and MgO are ionic compounds with giant ionic lattices; Al₂O₃ is primarily ionic but shows some covalent character; SiO₂ is a giant covalent network structure; P₂O₅ and Cl₂O₇ are molecular structures. These structural differences explain the dramatic differences in physical properties (melting point, solubility) and chemica

Inorganic Chemistry

Inorganic Chemistry

Period 3 oxides show a transition from basic (left) to acidic (right): Na₂O and MgO are basic (react with acid, form salt + water), Al₂O₃ is amphoteric (reacts with both acids and bases), SiO₂ is acidic (reacts with base to form salt), P₂O₅, SO₃, Cl₂O₇ are increasingly acidic. This reflects decreasing metallic character across the period: metals form basic oxides; nonmetals form acidic oxides. Amp

Inorganic Chemistry

Inorganic Chemistry

The oxidation state of elements in their Period 3 oxides increases across the period: Na is +1 in Na₂O, Mg is +2 in MgO, Al is +3 in Al₂O₃, Si is +4 in SiO₂, P is +5 in P₂O₅, S is +6 in SO₃, Cl is +7 in Cl₂O₇. This reflects the increasing number of valence electrons and the increasing tendency to lose electrons to achieve a noble gas configuration or form stable covalent bonds with oxygen. The ox

Inorganic Chemistry

Inorganic Chemistry

Period 3 oxides react differently with water depending on their acid-base character. Basic oxides (Na₂O, MgO) react to form hydroxides. Amphoteric Al₂O₃ doesn't react significantly with water but reacts with acids and alkalis. Silicon dioxide doesn't react with water. Acidic oxides (P₂O₅, SO₃) react vigorously with water to form oxyacids (H₃PO₄, H₂SO₄). The reaction of oxides with water depends o

Inorganic Chemistry

Inorganic Chemistry

Transition metals (or transition elements) are defined as elements that form at least one stable ion with a partially filled d orbital (d¹-d⁹ configuration). This definition includes the d-block elements (Sc to Zn in the 3d series, Y to Cd in the 4d series, etc.). A few elements like Cu and Cr are exceptions where the s orbital is involved to achieve stability, but they still form ions with partia

Inorganic Chemistry

Inorganic Chemistry

Transition metals can exist in multiple oxidation states (e.g., Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺, Mn²⁺/Mn³⁺/Mn⁷⁺) because electrons can be removed from both the d and s orbitals. The energy difference between d and s orbitals is small, allowing varying numbers of electrons to be ionized depending on reaction conditions. This flexibility enables redox reactions and makes transition metals excellent catalysts.

Inorganic Chemistry

Inorganic Chemistry

A complex ion consists of a central metal ion surrounded by a group of molecules or ions called ligands bonded through coordinate bonds (dative covalent bonds). A ligand is a molecule or anion that can donate an electron pair to a metal ion to form a coordinate bond. Common ligands include H₂O, NH₃, CN⁻, Cl⁻, and CO. The coordination number is the total number of coordinate bonds formed (usually 2

Inorganic Chemistry

Inorganic Chemistry

In ligand substitution reactions, one ligand is replaced by another ligand in a complex ion. The stability of the new complex compared to the old one determines whether substitution is favorable. For example, ammonia preferentially substitutes water in aqua complexes: [Cu(H₂O)₄]²⁺ + 4NH₃ → [Cu(NH₃)₄]²⁺ + 4H₂O. This occurs because [Cu(NH₃)₄]²⁺ is more stable than [Cu(H₂O)₄]²⁺. Ligand substitution

Inorganic Chemistry

Inorganic Chemistry

The coordination number is the total number of coordinate bonds formed between a central metal ion and surrounding ligands. The most common coordination numbers are 2, 4, and 6. Coordination number 6 is most common for transition metals, particularly with small ligands like H₂O, NH₃, and CN⁻. Coordination number 4 is typical for Cu²⁺ with larger ligands and for some tetrahedral complexes. Coordin

Inorganic Chemistry

Inorganic Chemistry

The shapes of complex ions are determined by their coordination number and the nature of the ligands. Coordination number 6 complexes are octahedral (e.g., [Fe(H₂O)₆]²⁺, [Cu(NH₃)₄]²⁺ actually square planar but derived from octahedral). Coordination number 4 complexes can be tetrahedral (e.g., [CuCl₄]²⁻) or square planar (e.g., [Cu(NH₃)₄]²⁺). Coordination number 2 complexes are linear (e.g., [Ag(NH

Inorganic Chemistry

Inorganic Chemistry

Transition metal ions and complexes are often colored because d electrons can absorb visible light and undergo electronic transitions between d orbitals with different energies. The energy of the absorbed light (which determines the color observed) depends on the crystal field splitting—the difference in energy between d orbitals caused by the electric field of surrounding ligands. Different ligan

Inorganic Chemistry

Inorganic Chemistry

Transition metals are excellent catalysts because they can change oxidation state easily, forming intermediate complexes that provide lower activation energy pathways for reactions. Homogeneous catalysts (dissolved in the reaction medium) like Fe³⁺ and Cr₂O₇²⁻ work through redox cycles. Heterogeneous catalysts (solid phase) like Fe in the Haber process and Pt in catalytic converters provide surfac

Inorganic Chemistry

Inorganic Chemistry

When sodium hydroxide or ammonia solution is added to solutions of transition metal ions, colored metal hydroxides precipitate. The color of the precipitate is characteristic: Fe²⁺ → green/white Fe(OH)₂, Fe³⁺ → brown Fe(OH)₃, Cu²⁺ → blue Cu(OH)₂, etc. These precipitates can be dissolved by adding excess acid (all hydroxides dissolve in strong acid) or excess ammonia (if the metal forms soluble amm

Inorganic Chemistry

Inorganic Chemistry

A ligand is a molecule or ion that donates an electron pair to form a coordinate (dative covalent) bond with a central metal ion. Ligands act as Lewis bases (electron pair donors). Common ligands include water (H₂O), ammonia (NH₃), chloride (Cl⁻), cyanide (CN⁻), and carbon monoxide (CO). Ligands can be monodentate (donating one electron pair, like NH₃ or Cl⁻) or polydentate (donating multiple elec

Inorganic Chemistry

Inorganic Chemistry

Metal aqua ions [M(H₂O)₆]ⁿ⁺ can act as weak acids because the positive charge on the metal ion polarizes the O-H bonds of the coordinated water molecules, making them more readily ionizable. The acidity increases with increasing charge on the metal ion and decreasing size (increasing charge density). Fe³⁺ aqua ions are more acidic than Fe²⁺. Highly charged small metal ions produce acidic solutions

Inorganic Chemistry

Inorganic Chemistry

Hydrolysis is the reaction of a metal aqua ion with water, resulting in loss of H⁺ and formation of a hydroxo complex or metal hydroxide. For example: [Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅(OH)]²⁺ + H₃O⁺ (or H⁺). Further hydrolysis produces [Al(H₂O)₄(OH)₂]⁺ and eventually Al(OH)₃ precipitate. This is an important source of acidity in solutions of metal salts, particularly for highly charged metal ions. H

Inorganic Chemistry

Inorganic Chemistry

Adding NaOH solution to metal aqua ions precipitates metal hydroxides. The precipitate color identifies the metal ion. Further addition of excess NaOH causes some metal hydroxides to dissolve, forming soluble hydroxo complexes or soluble aluminate anions. This behavior helps distinguish between different metal ions. Fe(OH)₃ is insoluble in excess NaOH (amphoteric but forms insoluble precipitate).

Inorganic Chemistry

Inorganic Chemistry

Adding ammonia solution to metal aqua ions first produces a metal hydroxide precipitate (because NH₃ is weakly basic), but further addition of excess ammonia dissolves the precipitate by forming soluble ammonia complexes [M(NH₃)ₙ]ᵐ⁺. This is a characteristic reaction that distinguishes ammonia from NaOH: NaOH gives precipitates that may or may not dissolve in excess NaOH, while ammonia often disso

Inorganic Chemistry

Inorganic Chemistry

Amphoteric hydroxides are compounds that can act as both bases (reacting with acids) and acids (reacting with bases). Aluminum hydroxide Al(OH)₃ and zinc hydroxide Zn(OH)₂ are the classic examples. They react with both strong acids and strong bases: Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O (base behavior) and Al(OH)₃ + NaOH → NaAlO₂ + 2H₂O (acid behavior). This amphoteric nature reflects their position on th

Inorganic Chemistry

Inorganic Chemistry

Qualitative analysis uses characteristic test tube reactions to identify ions in solution. Key tests include: silver nitrate for halides (white/yellow/cream precipitates), barium chloride for sulfates (white BaSO₄ precipitate), NaOH for metal ions (colored hydroxide precipitates), dilute HNO₃ with AgNO₃ for carbonate (white precipitate that dissolves in HNO₃), and flame tests for certain metals (N

Inorganic Chemistry