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

Organic Chemistry

A systematic naming system for organic compounds based on functional groups, carbon chain length, and substituent positions. The IUPAC name consists of: prefix (number of carbons), parent name (functional group), and suffix (functional group type). Example: 2-methylpropanoic acid has a 3-carbon chain with a methyl branch at position 2. IUPAC nomenclature ensures unambiguous identification and enab

Organic Chemistry

Organic Chemistry

A series of organic compounds with the same functional group, differing by successive CH₂ units, showing gradual changes in physical properties but similar chemical properties. Example: alkanes (methane, ethane, propane, butane) differ by CH₂, have similar C-H bonding, but boiling points increase progressively. Homologous series demonstrate periodic trends in organic chemistry. Homologous series

Organic Chemistry

Organic Chemistry

Specific groups of atoms within molecules responsible for characteristic chemical reactions. Common functional groups: −OH (hydroxyl, alcohols/phenols), −COOH (carboxyl, carboxylic acids), C=O (carbonyl, aldehydes/ketones), −NH₂ (amino), −OR (ether), C=C (alkene), C≡C (alkyne). Functional groups determine reactivity; compounds with same functional group undergo similar reactions regardless of the

Organic Chemistry

Organic Chemistry

Isomers with the same molecular formula but different structural arrangements of atoms (different connectivity). Types: chain isomerism (different carbon skeleton, e.g., butane vs. isobutane), position isomerism (same functional group at different position, e.g., 1-propanol vs. 2-propanol), and functional group isomerism (different functional groups, e.g., ethanol vs. methoxyethane, CH₃CH₂OH vs. C

Organic Chemistry

Organic Chemistry

Isomers with the same molecular formula and same atomic connectivity but different 3D spatial arrangement. E-Z isomerism occurs with restricted rotation around C=C (double bond cannot rotate). Each carbon of the double bond is bonded to two different groups; priority is assigned by atomic number (Cahn-Ingold-Prelog rules). Z (zusammen, together) has higher-priority groups on same side; E (entgegen

Organic Chemistry

Organic Chemistry

A general formula describes the relationship between carbon and hydrogen atoms (and other elements) in a homologous series. Examples: alkanes CₙH₂ₙ₊₂, cycloalkanes CₙH₂ₙ, alkenes CₙH₂ₙ, alkynes CₙH₂ₙ₋₂, carboxylic acids CₙH₂ₙO₂. The general formula enables predicting molecular composition from carbon count (e.g., C₅ alkane is C₅H₁₂) without knowing structural details. Different functional groups h

Organic Chemistry

Organic Chemistry

Alkanes (CₙH₂ₙ₊₂) are nonpolar hydrocarbons with only C-C and C-H single bonds. Properties: boiling point increases with molar mass (van der Waals forces increase), immiscible in water (nonpolar, hydrophobic), viscosity increases with chain length, density increases slightly. Alkanes are relatively unreactive at room temperature (strong C-C and C-H bonds), but undergo combustion and free-radical s

Organic Chemistry

Organic Chemistry

Combustion is a redox reaction where alkane CₙH₂ₙ₊₂ burns in excess oxygen: CₙH₂ₙ₊₂ + (3n+1)/2 O₂ → nCO₂ + (n+1)H₂O. Combustion is exothermic (ΔH < 0), providing energy for heating and vehicle engines. Complete combustion produces only CO₂ and H₂O. Incomplete combustion (limited O₂) produces CO and other products. Bond enthalpy calculation: strong bonds form (C=O in CO₂, O-H in H₂O) releasing more

Organic Chemistry

Organic Chemistry

A reaction mechanism where a free radical attacks an alkane C-H bond, removing H and replacing it with another atom/group. The mechanism has three stages: initiation (homolytic bond cleavage creating radicals), propagation (radical attacks molecule, generating new radical), termination (radicals combine, reaction stops). Example: CH₄ + Cl₂ → CH₃Cl + HCl (photochemically initiated). Free radicals a

Organic Chemistry

Organic Chemistry

Stages of free-radical reactions: Initiation—heat or light causes homolytic cleavage of X-X bond (e.g., Cl-Cl → 2Cl•), creating radicals. Propagation—radical attacks alkane (Cl• + CH₄ → •CH₃ + HCl), producing products and new radicals that repeat the cycle (chain reaction, single photon produces many products). Termination—two radicals combine (Cl• + •CH₃ → CH₃Cl), removing reactive intermediates

Organic Chemistry

Organic Chemistry

Photochemical reaction: CH₄ + Cl₂ → CH₃Cl + HCl (with light, hv). Mechanism: initiation (Cl-Cl → 2Cl•), propagation (Cl• + CH₄ → •CH₃ + HCl, then •CH₃ + Cl₂ → CH₃Cl + Cl•), termination. Further chlorination produces CH₂Cl₂, CHCl₃, CCl₄. The reaction is exothermic and dangerous (explosion risk if mixture contains Cl₂ and hydrocarbon in explosive proportions). Chlorination shows poor selectivity—all

Organic Chemistry

Organic Chemistry

A reaction where a nucleophile (electron-rich, seeking positive charge) replaces a leaving group on a carbon atom. Mechanism: SN1 (unimolecular) forms carbocation intermediate (slow step, rate = k[RX]); SN2 (bimolecular) proceeds via direct displacement with inversion of stereochemistry (rate = k[RX][Nu]). Conditions: SN1 favored by polar solvents, weak nucleophiles, good leaving groups (Br, I), a

Organic Chemistry

Organic Chemistry

SN1 (substitution, nucleophilic, unimolecular): rate = k[RX], two-step—first, alkyl halide dissociates slowly (RX → R⁺ + X⁻, carbocation forms), then nucleophile attacks fast (R⁺ + Nu⁻ → RNu). Carbocation is planar (sp²), allowing attack from either face (racemization). SN2 (bimolecular): rate = k[RX][Nu], one-step, nucleophile attacks from back (opposite to leaving group), displacing X⁻. Configur

Organic Chemistry

Organic Chemistry

A reaction where a small molecule (usually H₂O or HX) is removed from an alkyl halide or alcohol, forming an alkene with a C=C double bond. Mechanism E1 (unimolecular elimination): alkyl halide dissociates forming carbocation (slow), then H⁺ is removed by base (fast). E2 (bimolecular): base attacks H while X⁻ leaves (concerted, one step). Conditions: E1 favored by poor bases, polar solvents, stabl

Organic Chemistry

Organic Chemistry

Halogenoalkanes (RX, where X = Cl, Br, I) undergo nucleophilic substitution and elimination. Reactivity order for nucleophilic substitution: RI > RBr > RCl > RF (C-I bond is weakest, easiest to break). Reactivity for SN2: primary > secondary > tertiary (steric hindrance increases). For SN1: tertiary > secondary > primary (carbocation stability). Elimination favored over substitution at high temper

Organic Chemistry

Organic Chemistry

Chlorofluorocarbons (CFCs) like CFC-12 (CF₂Cl₂) were widely used as refrigerants and propellants. Released into atmosphere, they rise to stratosphere. UV light (hv) breaks C-Cl bond: CF₂Cl₂ + hv → CF₂Cl• + Cl•. Cl• radicals catalytically destroy ozone: Cl• + O₃ → ClO• + O₂, then ClO• + O → Cl• + O₂. One Cl atom can destroy ~100,000 O₃ molecules before being deactivated. Ozone depletion increases U

Organic Chemistry

Organic Chemistry

A reaction where an electrophile (electron-seeking, δ+) attacks the π electrons of an alkene C=C, adding across the double bond to form a saturated product. General mechanism: alkene π bond attacks electrophile (forming carbocation, slow), then nucleophile attacks carbocation (fast). Example: CH₂=CH₂ + H₂SO₄ → CH₃CH₂OSO₃H. Markownikoff's rule predicts which position the electrophile adds (to the m

Organic Chemistry

Organic Chemistry

HBr adds to alkene: C=C + HBr → CHBr-CH₂. Mechanism: π electrons attack H⁺ (forming more stable carbocation), then Br⁻ attacks from either face (carbocation is planar). Markownikoff's rule: H adds to the carbon with more hydrogen neighbors; Br adds to the more substituted carbon. Example: 2-methylpropene + HBr → 2-bromo-2-methylpropane (tert-butyl bromide, stable tertiary carbocation), not 2-bromo

Organic Chemistry

Organic Chemistry

When an unsymmetrical alkene adds an HX, the H adds to the carbon with more hydrogen neighbors (and X adds to the more substituted carbon). The rule predicts product: propene + HBr → 2-bromopropane (not 1-bromopropane). Explanation: reaction proceeds via the more stable carbocation intermediate. 2-bromopropane intermediate (secondary carbocation) is more stable than 1-bromopropane intermediate (pr

Organic Chemistry

Organic Chemistry

Bromine water (Br₂ in CCl₄, orange-brown color) decolorizes when added to an alkene or alkyne, indicating C=C or C≡C. The Br₂ adds across the double bond: C=C + Br₂ → CBr-CBr, removing Br₂ color and producing a clear, colorless (or pale yellow) solution. Saturated alkanes don't decolorize bromine water (no reaction). Aromatic rings require catalyst (FeBr₃) to react with Br₂, so don't readily decol

Organic Chemistry

Organic Chemistry

Polymerization where unsaturated monomers (with C=C or C≡C) sequentially add to a growing chain, forming a polymer without releasing byproducts. Free-radical mechanism: initiator generates •C, which attacks alkene (C=C + •R → •R-C-C), then chain grows as •R-C-C attacks another alkene (propagation). Example: n(C₂H₄) → (−C₂H₄−)ₙ (polyethylene). Markownikoff's rule applies in some cases. Radical poly

Organic Chemistry

Organic Chemistry

Alcohols are classified by the carbon bearing the −OH group: primary (RCH₂OH) has one R group, secondary (R₂CHOH) has two R groups, tertiary (R₃COH) has three R groups. Classification determines reactivity: primary alcohols oxidize to aldehydes then carboxylic acids; secondary oxidize to ketones only; tertiary resist oxidation (no H on the −OH carbon). Primary alcohols undergo SN2 nucleophilic sub

Organic Chemistry

Organic Chemistry

Primary alcohols RCH₂OH oxidize to aldehydes RCHO (1 stage) then to carboxylic acids RCOOH (2 stages) using acidified potassium dichromate (orange turns green, Cr₂O₇²⁻ → Cr³⁺). Secondary alcohols R₂CHOH oxidize to ketones R₂C=O (1 stage) but not further (no H on C-OH). Tertiary alcohols don't oxidize (no H on C-OH). Gentle conditions (PCC pyridinium chlorochromate, or controlled stoichiometry) giv

Organic Chemistry

Organic Chemistry

Removal of water from an alcohol by heating with concentrated H₂SO₄ (acid catalyst and dehydrating agent) produces an alkene: R₂CHOH → R₂C=C + H₂O. Mechanism E1: alcohol protonates (forming R₂C⁺OH₂), water leaves (forming carbocation), then base (H₂O or HSO₄⁻) removes adjacent H⁺. Reactivity: tertiary > secondary > primary (carbocation stability). Markownikoff's rule predicts major alkene product

Organic Chemistry

Organic Chemistry

Condensation of carboxylic acid RCOOH and alcohol R'OH, forming ester RCOOR' and releasing H₂O. Reaction: RCOOH + R'OH ⇌ RCOOR' + H₂O (acid-catalyzed, reversible). Mechanism: acid protonates C=O, alcohol attacks (nucleophilic acyl substitution), eliminates H₂O. Equilibrium lies left; to drive toward ester, use excess alcohol, heat, or remove water (Dean-Stark apparatus). Esters are used in polyest

Organic Chemistry

Organic Chemistry

Anaerobic (without oxygen) metabolic process where glucose is converted to ethanol and CO₂ by yeast or bacteria: C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂. Pathway: glycolysis (glucose → pyruvate), pyruvate → acetaldehyde → ethanol (NAD⁺ regeneration for continued glycolysis). Used historically for alcohol production (beer, wine), now also for biofuels. Yeast (Saccharomyces cerevisiae) is most common fermenter. Te

Organic Chemistry

Organic Chemistry

Qualitative tests identify functional groups: alcohols—Na reacts, fizz and heat (no fizz with ether); carbonyl (aldehydes/ketones)—2,4-DNP (yellow/orange precipitate), Fehling's (aldehydes give brick-red Cu₂O, ketones don't react); carboxylic acids—Na gives fizz, litmus turns red; esters—hydrolyze with NaOH, smell pleasant fruity odor; halogenoalkanes—AgNO₃ gives precipitate (Ag-X insoluble). Infr

Organic Chemistry

Organic Chemistry

IR spectroscopy measures absorption of infrared radiation by molecules, revealing functional groups. Characteristic absorptions: C-H stretch ~3000 cm⁻¹, O-H (alcohol) ~3300−3500 cm⁻¹ (broad, hydrogen bonding), O-H (carboxylic acid) ~2500−3300 cm⁻¹ (very broad), C=O ~1700 cm⁻¹ (carbonyl, exact frequency indicates aldehyde/ketone/carboxylic acid/ester), C=C ~1600−1680 cm⁻¹ (alkene, weak). Fingerprin

Organic Chemistry

Organic Chemistry

Mass spectrometry (MS) measures mass-to-charge ratio (m/z) of ions from organic molecules, revealing molecular mass and structure. Electron impact ionization (EI) removes electron, forming molecular ion M⁺ (peak at m/z = M). Molecule then fragments: bonds break, losing neutral species or forming cations. Fragmentation pattern (peaks at different m/z) reflects molecular structure. Common fragments:

Organic Chemistry

Organic Chemistry

In mass spectrometry, fragmentation patterns show peaks at m/z values where bonds break. The molecular ion M⁺ loses neutral molecules or rearranges. Common losses: loss of 18 (H₂O from alcohols), 17 (OH), 44 (CO₂ from carboxylic acids), 29 (CHO), or 43 (CH₃CO from methyl ketones). Peaks at characteristic m/z identify functional groups: carboxylic acids typically show m/z 45 (CHO₂⁺, McLafferty rear

Organic Chemistry

Organic Chemistry

A molecule is chiral if it has a stereocenter (usually a carbon with four different groups) and cannot be superimposed on its mirror image. The two forms are enantiomers. Most organic molecules with a chiral center exist as R and S isomers (Cahn-Ingold-Prelog rules assign priorities by atomic number). Chirality is important in pharmaceuticals: the drug (R)-isomer may be active; the (S)-isomer inac

Organic Chemistry

Organic Chemistry

Enantiomers are mirror-image stereoisomers that are non-superimposable. Chiral molecules with one stereocenter exist as a pair of enantiomers (R and S). Enantiomers have identical physical properties (mp, bp, density, solubility) and identical chemical properties with achiral reagents, but different optical activity (one is dextrorotatory (+), one levorotatory (−)). They react at different rates w

Organic Chemistry

Organic Chemistry

A racemic mixture (or racemate) contains equal amounts (50:50) of R and S enantiomers of a chiral compound. The mixture is optically inactive (equal (+) and (−) rotations cancel). Most synthetic routes produce racemates because they create chiral centers non-selectively. Separating enantiomers requires chiral chromatography or crystallization with chiral resolving agents. In nature, biosynthetic e

Organic Chemistry

Organic Chemistry

Chiral compounds rotate plane-polarized light, exhibiting optical activity. The rotation angle (α) is measured in degrees, with (+) or (−) indicating dextrorotatory or levorotatory. Specific rotation [α] = α / (c × l), where c is concentration (g/100mL) and l is path length (dm). Each enantiomer has opposite specific rotation (R is +, S is −, approximately). Racemic mixtures don't rotate light (eq

Organic Chemistry

Organic Chemistry

The C=O carbonyl (δ+ carbon) is attacked by nucleophiles (−, electron-rich), adding across the C=O to form C−OH intermediate. Mechanism: nucleophile attacks carbon (π bond breaks), forming negatively charged intermediate, then protonation gives product. Examples: RCN + HCN → RC(OH)CN (cyanohydrin, H added in workup), RCHO + NH₂-NH₂ → RCH=N-NH₂ (hydrazone), RCHO + H₂O → RCH(OH)₂ (geminal diol). Ald

Organic Chemistry

Organic Chemistry

Aldehydes and ketones are reduced to primary and secondary alcohols respectively: RCHO + 2H → RCH₂OH (aldehyde → primary alcohol), RCOR' + 2H → RCHOHR' (ketone → secondary alcohol). Reducing agents: NaBH₄ (mild, reduces aldehydes/ketones but not carboxylic acids), LiAlH₄ (strong, reduces aldehydes/ketones, esters, carboxylic acids), H₂ + catalyst (Ni, Pt). Mechanism involves hydride (H⁻) attacking

Organic Chemistry

Organic Chemistry

Aldehydes are oxidized to carboxylic acids: RCHO + [O] → RCOOH using oxidizing agents like acidified potassium dichromate (orange turns green), permanganate, or Tollens reagent. The key difference from alcohols: aldehydes have one more H on the carbonyl carbon, which is oxidized to −OH (carboxylic acid). Ketones cannot be oxidized under these conditions (no H on carbonyl carbon). Mild oxidizing ag

Organic Chemistry

Organic Chemistry

Tollens reagent is ammoniacal silver nitrate (Ag⁺ in ammonia solution); Fehlings is alkaline copper(II) sulfate. Both are oxidizing agents used as qualitative tests: aldehydes give positive test, ketones don't. With Tollens: aldehyde reduces Ag⁺ to Ag (metal mirror forms on test tube—'silver mirror test'). RCHO + 2Ag(NH₃)₂⁺ + 3OH⁻ → RCOO⁻ + 2Ag↓ + 4NH₃ + 2H₂O. With Fehlings: aldehyde reduces Cu²⁺

Organic Chemistry

Organic Chemistry

2,4-Dinitrophenylhydrazine (2,4-DNPH) is a reagent that reacts with aldehydes and ketones (not alcohols, amines). Forms yellow/orange precipitate (2,4-DNP derivative): R₂C=O + H₂N-NH-C₆H₃(NO₂)₂ → R₂C=N-NH-C₆H₃(NO₂)₂ + H₂O. The precipitate can be filtered, recrystallized, and melting point determined (unique for each carbonyl compound). This is a classic qualitative test and identification method:

Organic Chemistry

Organic Chemistry

Carboxylic acids (R−COOH) are weakly acidic (Ka ~ 10⁻⁵ for aliphatic acids), hydrogen bond extensively (−COOH forms dimers and polymers), have high boiling points (hydrogen bonding), and are soluble in water (polar −COOH, hydrogen bonding). Reactivity: undergo nucleophilic acyl substitution (esterification with alcohols, forming amides with amines), dimerization (intermolecular hydrogen bonding),

Organic Chemistry

Organic Chemistry

Acyl chlorides (R−COCl, acid chlorides) are highly reactive compounds formed by treating carboxylic acids with SOCl₂ or PCl₃. Reactivity: readily undergo nucleophilic acyl substitution (with alcohols → esters, with amines → amides, with water → carboxylic acids). Mechanism: nucleophile attacks the δ+ carbonyl carbon; −Cl is an excellent leaving group (Cl⁻ is very stable). The high reactivity stems

Organic Chemistry

Organic Chemistry

Acid anhydrides (R−CO−O−CO−R, or (RCO)₂O) are formed by dehydration of carboxylic acids or reaction of acyl chloride with carboxylate salt. Reactivity: undergo nucleophilic acyl substitution (with alcohols → esters, with amines → amides, with water → carboxylic acids), less reactive than acyl chlorides but more reactive than esters. Mechanism: nucleophile attacks carbonyl carbon; either RCO₂⁻ (car

Organic Chemistry

Organic Chemistry

Organic compounds with the functional group −COOR (or −OCOR), formed by esterification (condensation of a carboxylic acid R−COOH and an alcohol R'−OH). Esters are generally unreactive toward nucleophiles at room temperature (unlike acid halides or anhydrides). Esters have pleasant odors (e.g., ethyl ethanoate smells fruity) and are used as flavorings, solvents, and in polymer synthesis. Esters ar

Organic Chemistry

Organic Chemistry

Fischer esterification is the reaction of carboxylic acid and alcohol in acid catalyst to form ester: R−COOH + R'−OH ⇌ R−COO−R' + H₂O (H₂SO₄ or HCl catalyst). Mechanism: acid protonates C=O (activating), alcohol attacks nucleophilically, proton transfers, and H₂O leaves. The reaction is reversible; equilibrium favors products when excess alcohol is used or water is removed. Ester formation is nucl

Organic Chemistry

Organic Chemistry

Esters hydrolyze (break apart) in acid or base to regenerate carboxylic acid and alcohol. Acid hydrolysis (reversible): RCOOR' + H₂O ⇌ RCOOH + R'OH (H⁺ catalyst, equilibrium favors reactants). Base hydrolysis (saponification, irreversible): RCOOR' + NaOH → RCOONa + R'OH (OH⁻ nucleophile attacks, salt formed). Mechanism: nucleophile (H₂O or OH⁻) attacks carbonyl carbon (nucleophilic acyl substituti

Organic Chemistry

Organic Chemistry

Benzene (C₆H₆) has alternating C−C and C=C bonds (Kekule structure), but experimental evidence (constant bond lengths ~1.39 Å, not 1.48 Å for C−C or 1.34 Å for C=C) reveals delocalized π electrons. Modern understanding: π electrons delocalize over all 6 carbons (resonance structures). Representation: hexagon with circle inside (indicates delocalization). The delocalization stabilizes benzene (reso

Organic Chemistry

Organic Chemistry

In aromatic electrophilic substitution, benzene ring (π electrons, nucleophilic) attacks an electrophile (electron-poor, δ+), replacing a hydrogen. General mechanism: electrophile attacks ring (forming carbocation intermediate), then H⁺ is removed, restoring aromaticity. Example: Br₂ + FeBr₃ → Br⁺ (electrophile), attacks benzene, forms bromobenzene. Nitration: HNO₃ + H₂SO₄ → NO₂⁺ (electrophile), p

Organic Chemistry

Organic Chemistry

Nitration: benzene + HNO₃ (with conc. H₂SO₄ catalyst) → nitrobenzene (C₆H₅NO₂) + H₂O. Mechanism: HNO₃ + H₂SO₄ → NO₂⁺ (nitronium ion, electrophile) + HSO₄⁻. NO₂⁺ attacks benzene π electrons (electrophilic aromatic substitution), forming carbocation intermediate. H⁺ removed by HSO₄⁻ (restoring aromaticity). Nitrobenzene is pale yellow, used as solvent and intermediate for reducing to aniline (C₆H₅NH

Organic Chemistry

Organic Chemistry

Friedel-Crafts alkylation: benzene + RCl + AlCl₃ → alkylbenzene + HCl. Mechanism: RCl + AlCl₃ → R⁺ (carbocation, electrophile) + AlCl₄⁻. R⁺ attacks benzene π electrons (electrophilic aromatic substitution), forming carbocation intermediate. H⁺ removed (restoring aromaticity). Example: benzene + CH₃Cl + AlCl₃ → toluene (C₆H₅CH₃). The alkyl group (−R) is electron-donating, activates ring (makes furt

Organic Chemistry

Organic Chemistry

Friedel-Crafts acylation: benzene + RCOCl + AlCl₃ → benzophenone/phenyl alkyl ketone + HCl. Mechanism: RCOCl + AlCl₃ → RCO⁺ (acylium ion, electrophile, resonance-stabilized). RCO⁺ attacks benzene π electrons, forming carbocation intermediate. H⁺ removed (restoring aromaticity). Example: benzene + CH₃COCl + AlCl₃ → acetophenone (C₆H₅COCH₃, methyl phenyl ketone). Unlike alkylation, acylation doesn't

Organic Chemistry

Organic Chemistry

Amines are classified by carbon-nitrogen bonds: primary (RNH₂, one carbon on N), secondary (R₂NH, two carbons on N), tertiary (R₃N, three carbons on N). Affects reactivity: primary amines are most nucleophilic (lone pair unhindered), secondary intermediate, tertiary least (steric hindrance). Primary amines can be oxidized (to imines → aldehydes); secondary can't be oxidized easily (no H on N). Bas

Organic Chemistry

Organic Chemistry

Amines are prepared by: (1) nucleophilic substitution—halogenoalkane + NX (where X = Br, I, or CN with KCN → nitrile → amine via LiAlH₄ reduction), (2) reduction—nitrile RCN + LiAlH₄ → RCH₂NH₂ (primary amine); amide RCON R' + LiAlH₄ → amine; carbonyl + NH₃ + reducing agent (imine reduction), (3) arene + HNO₃/H₂SO₄ → nitrobenzene → reduction (Sn/HCl or Fe/AcOH) → aniline. Nucleophilic substitution

Organic Chemistry

Organic Chemistry

Amines are weak bases due to lone pair on nitrogen. Primary amines RNH₂ (most basic, pKb ~3−4), secondary (pKb ~3−5), tertiary (pKb ~3−4). Aromatic amines (aniline pKb ~9, much weaker) because lone pair delocalizes into aromatic ring (unavailable for protonation). Inductive effects: electron-donating alkyl groups increase basicity (primary < secondary because steric hindrance), electron-withdrawin

Organic Chemistry

Organic Chemistry

Amines are nucleophiles (lone pair attacks electrophiles). Key reactions: (1) nucleophilic substitution—RNH₂ + RX → RNH−R + HX (primary amine displaces halide), (2) acylation—RNH₂ + RCOCl → RCONHR + HCl (amide formation, used for protection), (3) with nitrous acid—RNH₂ + HNO₂ → unstable diazonium (decomposes, gives alcohol or carbocation), (4) with carbonyls—RNH₂ + R'CHO → RN=CHR' + H₂O (imine for

Organic Chemistry

Organic Chemistry

Addition polymers form when monomers with C=C bonds add to each other sequentially, without losing byproducts. Mechanism: free-radical initiator creates •C, attacks alkene (C=C + •R → •R−C−C), chain grows as •R−C−C attacks more alkenes. Examples: polyethylene (−CH₂−CH₂−)ₙ from ethene, polypropylene from propene, polystyrene from styrene, PVC from chloroethene, PTFE from tetrafluoroethene. Markowni

Organic Chemistry

Organic Chemistry

Condensation polymers form when monomers join by eliminating small molecules (H₂O, MeOH, HCl). Require difunctional monomers: dicarboxylic acid + diol → polyester (−CO−O−), dicarboxylic acid + diamine → polyamide (−CO−NH−). Reaction is reversible (equilibrium), so high conversion requires removing water or using activated monomers (acid chlorides, anhydrides). Polyesters: HOOC−R−COOH + HO−R'−OH →

Organic Chemistry

Organic Chemistry

Polyesters are condensation polymers with repeating −CO−O− ester linkages, formed from dicarboxylic acids and diols. Reaction: HOOC−R−COOH + HO−R'−OH → [−CO−R−CO−O−R'−O−]ₙ + nH₂O (esterification, acid-catalyzed, reversible). Key examples: PET (polyethylene terephthalate, from terephthalic acid + ethylene glycol)—bottles, fibers; unsaturated polyester (with C=C in backbone)—composite resins. Proper

Organic Chemistry

Organic Chemistry

Polyamides (nylons) are condensation polymers with repeating −CO−NH− amide linkages. Reaction: HOOC−(CH₂)ₘ−COOH + H₂N−(CH₂)ₙ−NH₂ → [−CO−(CH₂)ₘ−CO−NH−(CH₂)ₙ−NH−]ₓ + xH₂O. Named by carbon counts: nylon-6,6 (adipic acid m=4 + hexamethylene diamine n=6). Properties: strong, durable, flexible, excellent for textiles and engineering. Strength from hydrogen bonding between −CO and −NH on adjacent chains

Organic Chemistry

Organic Chemistry

Amino acids have general structure: H₂N−CHR−COOH (amino group, chiral center, carboxylic acid). R is the side chain, varying for each amino acid (glycine R = H, alanine R = CH₃, etc.). Chiral center: all amino acids except glycine are chiral (have R and S enantiomers); naturally occurring amino acids are L-enantiomers (by convention, not absolute configuration). In solution pH-dependent: at low pH

Organic Chemistry

Organic Chemistry

Zwitterions (dipolar ions) have both positive and negative charges on the same molecule, characteristic of amino acids. At the isoelectric point (pI), amino acids exist as zwitterionic H₃N⁺−CHR−COO⁻ (protonated amino, deprotonated carboxyl). At low pH, both groups protonated (H₃N⁺−CHR−COOH); at high pH, both deprotonated (H₂N−CHR−COO⁻). The pI is the pH where the amino acid is zwitterionic and ele

Organic Chemistry

Organic Chemistry

Peptide bonds (−CO−NH−) form between amino acids during protein synthesis: carboxyl (−COOH) of one amino acid reacts with amino (−NH₂) of another, releasing water (condensation). Reaction: R₁CH(NH₂)COOH + H₂N−CHR₂COOH → R₁CH(NH₂)CO−NH−CHR₂COOH + H₂O. The bond is planar (resonance: C=O ↔ C-O⁻, restricted rotation around C-N). Polypeptides are chains of amino acids linked by peptide bonds (backbone

Organic Chemistry

Organic Chemistry

Protein structure has four levels: primary (amino acid sequence, determined by peptide bonds), secondary (α-helices, β-sheets stabilized by backbone hydrogen bonds between −CO and −NH), tertiary (3D shape from side chain interactions: H-bonds, disulfide bonds S-S, hydrophobic interactions), and quaternary (multiple polypeptide chains assembled together, like hemoglobin 4 subunits). Primary determi

Organic Chemistry

Organic Chemistry

DNA nucleotides consist of three components: a deoxyribose sugar (5-carbon), a phosphate group (−PO₄²⁻), and a nitrogenous base (purine: adenine A, guanine G; pyrimidine: cytosine C, thymine T). Nucleotides link via phosphodiester bonds (phosphate oxygen bonds to 3'−OH of sugar to 5'−OH of next sugar), forming the DNA backbone (−sugar−phosphate−sugar−phosphate− repeating). Bases attach via glycosi

Organic Chemistry

Organic Chemistry

Multi-step synthesis combines multiple reactions to convert a starting material into a desired product through intermediates. Strategy: identify functional group transformations needed, choose compatible reactions, consider regiochemistry and stereochemistry, and plan protecting group use if needed. Example: converting ethylbenzene to phenylacetic acid requires: (1) oxidize methyl to carboxylic ac

Organic Chemistry

Organic Chemistry

Retrosynthesis works backward from a target molecule, breaking bonds and proposing precursors for each step. Strategy: identify key functional groups and bonds to break, suggest synthetic equivalents (umpolung for disconnections), and repeat until reaching available starting materials. Example: to synthesize acetophenone C₆H₅COCH₃, disconnect C-CO bond: retro-Friedel-Crafts suggests benzene + CH₃C

Organic Chemistry

Organic Chemistry

Synthetic routes are sequences of reactions chosen to convert starting material to product. Considerations: (1) available starting materials, (2) required transformations (oxidation, reduction, substitution, elimination, addition), (3) functional group compatibility (protecting groups if needed), (4) regio/stereochemistry control, (5) yield and cost of reagents. Example: ethylbenzene → benzoic aci

Organic Chemistry

Organic Chemistry

Protecting groups are functional groups temporarily attached to prevent reactivity of sensitive groups during synthesis. Strategy: (1) protect sensitive group (e.g., −OH as −O−TMS), (2) perform desired transformation on another group, (3) remove protecting group (−O−TMS + F⁻ → −OH). Common: −OH protected as −O−Ac (acetyl), −O−Bn (benzyl), −O−TMS (trimethylsilyl); −NH₂ protected as −N−Boc (tert-but

Organic Chemistry

Organic Chemistry

¹³C NMR shows carbon chemical shifts (δ, ppm relative to TMS at 0 ppm), revealing carbon environments in molecules. Chemical shift ranges: alkyl C (0−50 ppm), C-O or C-N (50−100 ppm), unsaturated C (100−150 ppm), aromatic C (120−150 ppm), carbonyl C (150−220 ppm). Each carbon atom produces a peak (or splits into multiplets via ¹H-¹³C coupling if protons attached). DEPT (Distortionless Enhancement

Organic Chemistry

Organic Chemistry

¹H NMR shows proton chemical shifts (δ, ppm relative to TMS at 0 ppm), revealing proton environments and their numbers in molecules. Chemical shifts: alkyl (0−3 ppm), α to electron-withdrawing (3−5 ppm), aromatic (7−8 ppm), aldehyde (9−10 ppm), carboxylic acid/phenol (10−13 ppm). Integrations (peak areas) indicate relative numbers of protons. Spin-spin coupling (J-values, Hz) to nearby protons spl

Organic Chemistry

Organic Chemistry

Chemical shift (δ) measures the difference in resonance frequency between a nucleus and a reference standard (TMS, tetramethylsilane, arbitrarily set to 0 ppm), expressed in parts per million (ppm). Calculated as: δ = (ν_sample − ν_reference) / ν_reference × 10⁶. Chemical shift indicates electronic environment: nuclei shielded by electrons resonate at lower frequency (upfield, smaller δ), deshield

Organic Chemistry

Organic Chemistry

Spin-spin coupling (J-coupling) causes NMR signals to split into multiplets due to interaction with neighboring magnetic nuclei (usually ³J for vicinal protons, three bonds away). Coupling constant J (in Hz) indicates coupling strength (independent of spectrometer field strength). ³J(HH) ~ 6−8 Hz for vicinal coupling (e.g., −CH−CH−), ²J for geminal (two bonds, −CH₂−, ~12−15 Hz), ⁴J for long-range

Organic Chemistry

Organic Chemistry

Integration (peak area in ¹H NMR) indicates the number of protons in an environment (proportional to number of equivalent protons). Example: ethanol CH₃CH₂OH has 3H (methyl CH₃, integral = 3), 2H (methylene CH₂, integral = 2), 1H (−OH, integral = 1). Equivalent protons (same chemical environment due to symmetry or rapid rotation) produce single peak. Example: cyclohexane all 12 H are equivalent (r

Organic Chemistry

Organic Chemistry

A chromatographic technique where a thin layer of adsorbent (usually silica gel or alumina) on a plate is used to separate organic compounds. A mixture is applied to the baseline, and a mobile phase (solvent) moves up the plate. Different compounds move at different rates based on their affinity for the stationary phase. In thin-layer chromatography (TLC), a glass or plastic plate is coated with

Organic Chemistry

Organic Chemistry

Gas chromatography (GC) separates volatile organic compounds into individual components before detection. Apparatus: sample injector (evaporates sample), heated column (packed or capillary with stationary phase, e.g., silica), carrier gas flow (mobile phase, usually He or N₂), and detector (FID flame ionization, measures organic compounds; ECD electron capture, detects halogenated compounds). Sepa

Organic Chemistry

Organic Chemistry

The retardation factor in chromatography, defined as the ratio of the distance traveled by a compound to the distance traveled by the solvent front. Rf = (distance traveled by compound) / (distance traveled by solvent front). Rf values are characteristic of compounds under specific conditions and range from 0 to 1. The Rf value is a quantitative measure of a compound's movement in chromatography.

Organic Chemistry

Organic Chemistry

Carbocation stability depends on alkyl substitution: tertiary (R₃C⁺) > secondary (R₂CH⁺) > primary (RCH₂⁺) > methyl (CH₃⁺). Stabilization from electron-donating alkyl groups (hyperconjugation and inductive effects). Alkyl groups donate electron density via σ bonds (hyperconjugation: C-H bonds overlap with empty p orbital, stabilizing C⁺). Resonance stabilization: allylic (CH₂=CH−CH₂⁺ ↔ CH₂−CH=CH⁺)

Organic Chemistry

Organic Chemistry

Benzene (C₆H₆) is the simplest aromatic hydrocarbon, a planar six-membered ring of carbon atoms with alternating σ and π bonds. The π electrons (6 total) are delocalized over all six carbons, stabilizing the ring (resonance energy ~150 kJ/mol). Experimental evidence: all C−C bond lengths are identical (~1.39 Å, between single ~1.48 Å and double ~1.34 Å), confirming delocalization. Kekule structure

Organic Chemistry

Organic Chemistry

Aromaticity describes the stability and reactivity of aromatic compounds like benzene. Requirements (Hückel's rule): (1) monocyclic, (2) planar, (3) fully conjugated π system, (4) 4n+2 π electrons (n = 0, 1, 2...), so 2, 6, 10, 14... electrons. Benzene (6 π electrons, n=1) is aromatic. Cyclopentadienyl anion (C₅H₅⁻, 6 π electrons) is aromatic (stable). Cyclobutadiene (4 π electrons, 4n with n=1) i

Organic Chemistry

Organic Chemistry

Nitriles (RCN, also called isocyanides or isonitriles for R−NC) contain a C≡N triple bond (two π bonds and one σ bond). Formed by: nucleophilic substitution (RX + KCN → RCN + X⁻), dehydration of primary amides (RCONH₂ → RCN), or oxidation of primary amines. Nitriles are polar (δ+ carbon, δ− nitrogen), making them good electrophiles for nucleophilic addition. Reactions: reduce to primary amines (RC

Organic Chemistry

Organic Chemistry

Disulfide bonds (S−S, disulfide bridges) form between two cysteine residues via oxidation of their sulfhydryl (−SH) groups: 2 R−SH → R−S−S−R + 2H⁺ + 2e⁻. Disulfide bonds are covalent, strong (100 kJ/mol), and crucial for stabilizing protein tertiary and quaternary structure. Found in extracellular proteins (harsh conditions favor disulfide formation) and intracellular proteins (reducing environmen

Organic Chemistry

Organic Chemistry

Tertiary structure is the 3D folding of a single polypeptide chain, determined by interactions between side chains (R groups): hydrogen bonds between backbone and side chains, electrostatic interactions (salt bridges) between charged residues, disulfide bonds between cysteines, hydrophobic interactions (nonpolar residues cluster in protein core), and van der Waals forces. The folding is driven by

Organic Chemistry

Organic Chemistry

Phosphodiester bonds link nucleotides in DNA/RNA backbone: the 3'−OH of one sugar's ribose forms an ester bond with phosphate, which esterifies the 5'−OH of the next sugar. Reaction: sugar1-3'−OH + HO−PO₃²⁻ + 5'−OH−sugar2 → sugar1-3'−O−PO₃−O−5'−sugar2 (with net release of water). The phosphodiester linkage is negatively charged (PO₃²⁻ at physiological pH), making DNA/RNA polyanionic (repulsion req

Organic Chemistry

Organic Chemistry

Nucleophilic acyl substitution is a reaction where a nucleophile attacks the δ+ carbonyl carbon of an acyl compound (RCOX), displacing the leaving group X⁻ and forming a new acyl product (RCONY). General mechanism: nucleophile attacks C=O (π bond breaks, σ bond to X becomes weaker), tetrahedral intermediate forms, then X⁻ leaves (restoration of C=O double bond). Reactivity order (most to least rea

Organic Chemistry

Organic Chemistry

Amides (RCONR'₂, primary RCONH₂, secondary RCONHR', tertiary RCONR'₂) contain C(=O)−N bonds. Formed by: nucleophilic acyl substitution of acyl chlorides, anhydrides, or esters with amines (RCOCl + R'NH₂ → RCONHR' + HCl), or condensation of carboxylic acids with amines (requires activation or dehydrating agent). Amides are relatively unreactive (N lone pair is delocalized into C=O via resonance, re

Organic Chemistry

Organic Chemistry

Carbocation rearrangement occurs when a carbocation is unstable; adjacent C-H or C-C bonds shift to form a more stable carbocation. Mechanisms: (1) hydride shift (1,2-hydride), (2) methyl/alkyl shift (1,2-alkyl shift). Example: 1,2-dimethylpropyl cation (primary) is unstable; methyl shifts from adjacent carbon to form tertiary carbocation (more stable). Rearrangement: CH₃−CH(CH₃)⁺−CH₃ → CH₃−C(CH₃)

Organic Chemistry

Organic Chemistry

Decarboxylation is loss of CO₂ (carbon dioxide) from a molecule, typically carboxylic acids (RCOOH → RH + CO₂) or related compounds. Reactions: (1) thermal decarboxylation (heating carboxylic acids loses CO₂, especially β-keto acids, β-dicarboxylic acids), (2) photochemical (α-amino acids under UV), (3) enzymatic (carboxylases, particularly important in biochemistry—pyruvate decarboxylase converts

Organic Chemistry

Physical Chemistry

The basic building blocks of all matter. Protons (positive charge, found in nucleus), neutrons (neutral, found in nucleus), and electrons (negative charge, orbiting nucleus). Protons and neutrons have roughly equal mass (~1 amu), while electrons are much lighter (~1/2000 amu). The number of protons defines the element; the number of electrons determines the overall charge of an atom or ion. All a

Physical Chemistry

Physical Chemistry

The mass number (A) is the total number of protons and neutrons in a nucleus. Isotopes are atoms of the same element (same proton number) with different numbers of neutrons, and therefore different mass numbers. Isotopes have identical chemical properties because they have the same electron configuration, but different physical properties (e.g., density, melting point). Relative atomic mass is the

Physical Chemistry

Physical Chemistry

The arrangement of electrons in an atom, described using shells (energy levels) and subshells (s, p, d, f orbitals). Electrons fill orbitals in order of increasing energy, following the Aufbau principle. For example, nitrogen (N, atomic number 7) has configuration 1s² 2s² 2p³. Configuration determines an element's chemical properties and bonding behaviour. Electrons occupy orbitals in shells arou

Physical Chemistry

Physical Chemistry

The energy required to remove one mole of electrons from one mole of gaseous atoms or ions. First ionisation energy (IE₁) removes the first electron; second ionisation energy (IE₂) removes the second, etc. Measured in kJ mol⁻¹. Ionisation energies increase across a period (increased nuclear charge) and decrease down a group (increased atomic radius, shielding effect). Ionisation energy reflects h

Physical Chemistry

Physical Chemistry

An analytical technique that measures the mass-to-charge ratio of ions by timing how long they take to travel a fixed distance. Atoms/molecules are ionised (usually by electron impact), accelerated through an electric field, then drift along a field-free region. Lighter ions reach the detector faster than heavier ones. Time of flight correlates with mass, producing a mass spectrum showing relative

Physical Chemistry

Physical Chemistry

A region of space where there is a high probability of finding an electron. Orbitals are described by quantum numbers and have characteristic shapes: s orbitals are spherical, p orbitals are dumbbell-shaped, d orbitals are more complex. Each orbital can hold a maximum of 2 electrons (with opposite spins). Orbitals are grouped into subshells (s, p, d, f) and shells (n=1, 2, 3...). Quantum mechanic

Physical Chemistry

Physical Chemistry

The weighted average mass of all isotopes of an element, relative to one-twelfth of the mass of carbon-12 (which is assigned exactly 12). Denoted Ar. For example, chlorine has Ar = 35.5 because it's a mixture of Cl-35 (75.8%) and Cl-37 (24.2%). Relative atomic masses are found on the periodic table and are dimensionless (no units, though sometimes 'u' for atomic mass units is used). Relative atom

Physical Chemistry

Physical Chemistry

The sum of the relative atomic masses of all atoms in a molecule. Denoted Mr. For example, H₂O has Mr = (2 × 1) + 16 = 18. Relative molecular mass is dimensionless but numerically equals the molar mass in g mol⁻¹. Used to convert between mass and number of moles, and to calculate percentage composition. Relative molecular mass is simply the sum of Ar values weighted by atom counts. For glucose C₆

Physical Chemistry

Physical Chemistry

The mole is the SI unit of amount of substance. One mole contains 6.02 × 10²³ particles (atoms, molecules, ions, electrons, etc.)—Avogadro's constant (NA). The number of moles in a sample is calculated as: n = mass/Mr (for compounds) or n = mass/Ar (for elements). One mole of any gas at RTP occupies 24 dm³ (or 24,000 cm³). The mole connects the atomic scale to the laboratory scale. One mole of ca

Physical Chemistry

Physical Chemistry

The equation PV = nRT relates pressure (P in Pa), volume (V in m³), number of moles (n), and absolute temperature (T in K). R is the gas constant (8.31 J K⁻¹ mol⁻¹). Rearranged forms: P = ρRT/M (where ρ is density, M is molar mass), or used to find moles: n = PV/RT. Ideal gases obey this equation perfectly; real gases approximate it except at very high pressures or low temperatures. The ideal gas

Physical Chemistry

Physical Chemistry

The empirical formula shows the simplest whole-number ratio of atoms in a compound. The molecular formula shows the actual number of atoms. For example, ethene (C₂H₄) and benzene (C₆H₁₂) both have empirical formula CH₂, but different molecular formulas. Finding empirical formula requires: convert mass percentages (or masses) to moles, divide by the smallest value to find whole-number ratios. Findi

Physical Chemistry

Physical Chemistry

A chemical equation with equal numbers of each type of atom on both sides (conserving mass). Coefficients show the mole ratio of reactants to products. For example, 2Na + Cl₂ → 2NaCl shows that 2 moles of sodium react with 1 mole of chlorine to produce 2 moles of sodium chloride. Balancing is essential for stoichiometric calculations. A balanced equation shows equal numbers of each atom type on b

Physical Chemistry

Physical Chemistry

The ratio of actual yield to theoretical yield, expressed as a percentage: % yield = (actual mass / theoretical mass) × 100%. Theoretical yield assumes complete reaction with no losses; actual yield is what's obtained in the lab. Percentage yield is typically less than 100% due to incomplete reactions, side reactions, or product loss during isolation and purification. Theoretical yield is calcula

Physical Chemistry

Physical Chemistry

A measure of how efficiently a reaction uses raw materials, calculated as: atom economy (%) = (Mr of desired product / Σ Mr of all reactants) × 100%. High atom economy (close to 100%) means minimal waste; low atom economy means significant by-products. Unlike percentage yield, atom economy depends on the reaction stoichiometry, not lab performance. Important in green chemistry and industrial proce

Physical Chemistry

Physical Chemistry

The amount of solute per unit volume of solution, typically expressed in mol dm⁻³ (or M, molar). Concentration (c) = n/V, where n is moles of solute and V is volume of solution in dm³. Standard solutions are prepared by dissolving a known mass of solute in distilled water, transferring to a volumetric flask, and diluting to the mark. Dilution formula: c₁V₁ = c₂V₂. Concentration is crucial for rea

Physical Chemistry

Physical Chemistry

The electrostatic attraction between oppositely charged ions. Forms when electrons are transferred from a metal (low ionisation energy) to a nonmetal (high electron affinity). Ionic compounds contain discrete cations and anions arranged in crystal lattices. Common in compounds like NaCl, MgO, and CaCO₃. Ionic compounds conduct electricity when molten or dissolved (mobile ions). Ionic bonds result

Physical Chemistry

Physical Chemistry

The sharing of a pair of electrons between two atoms. Single bonds share 2 electrons; double bonds share 4; triple bonds share 6. Covalent bonds are strong (typically 150-1000 kJ mol⁻¹) and form between nonmetals or between nonmetals and metalloids. Covalent compounds exist as discrete molecules (simple covalent) or giant networks (covalent lattices). Covalent bonds form when two atoms share elec

Physical Chemistry

Physical Chemistry

A covalent bond in which both electrons come from the same atom (the donor), which has a lone pair, while the other atom (acceptor) provides the empty orbital. Also called a coordinate bond. Denoted by an arrow A→B. Once formed, dative bonds are indistinguishable from ordinary covalent bonds. Common in complexes, adducts, and compounds with boron or aluminium. A dative bond (coordinate covalent b

Physical Chemistry

Physical Chemistry

The electrostatic attraction between delocalized electrons and metal cations arranged in a lattice. Metal atoms lose valence electrons to form a 'sea' of mobile electrons that delocalizes across the entire structure. Metallic bonding explains metals' properties: electrical conductivity (mobile electrons), thermal conductivity, malleability, ductility, and lustre (light absorption and re-emission b

Physical Chemistry

Physical Chemistry

The three-dimensional arrangement of atoms, ions, or molecules in a repeating lattice pattern. Crystal structures include ionic (e.g., NaCl with octahedral coordination), covalent (e.g., diamond with tetrahedral bonding), molecular (e.g., ice with hydrogen bonding), and metallic (e.g., copper with metal cations in electron sea). Properties depend on structure: ionic crystals are hard and brittle;

Physical Chemistry

Physical Chemistry

Molecular geometry predicted by Valence Shell Electron Pair Repulsion (VSEPR) theory: electron pairs (bonding and lone pairs) around a central atom repel each other, arranging themselves to maximize distance. Electron geometry (including lone pairs) differs from molecular geometry (only atoms). Common shapes: tetrahedral (CH₄), trigonal planar (BF₃), linear (CO₂), bent (H₂O), trigonal pyramidal (N

Physical Chemistry

Physical Chemistry

Electronegativity is an atom's ability to attract electrons in a covalent bond (Pauling scale, 0.7-4.0). Bond polarity arises from unequal electron sharing when atoms have different electronegativities. In polar covalent bonds (e.g., HCl), electrons shift toward the more electronegative atom, creating a dipole (δ+ and δ-). Polar molecules have a permanent dipole moment; nonpolar molecules have dip

Physical Chemistry

Physical Chemistry

Weak attractions between separate molecules (or atoms): London dispersion forces (induced dipoles, present in all substances), permanent dipole-dipole forces (in polar molecules), and hydrogen bonding (O-H, N-H, F-H groups with lone pairs). Much weaker than covalent/ionic bonds (typically 0.5-50 kJ mol⁻¹ vs. 100-500 kJ mol⁻¹), but critical for melting/boiling points, solubility, and viscosity. Lo

Physical Chemistry

Physical Chemistry

The heat energy change at constant pressure during a reaction, denoted ΔH (kJ mol⁻¹). Exothermic reactions (ΔH < 0) release heat; endothermic reactions (ΔH > 0) absorb heat. Standard enthalpy change (ΔH°) refers to conditions where all substances are in standard states (298 K, 100 kPa, 1 M solutions). Measured using calorimetry or calculated from bond enthalpies and Hess's law. Enthalpy change ΔH

Physical Chemistry

Physical Chemistry

Experimental measurement of heat energy change. Simple calorimeter: insulated container holding the reaction, thermometer to measure temperature change, heat generated/absorbed by reaction changes the temperature of the liquid (usually water). Heat (q) = mcΔT, where m is mass (g), c is specific heat capacity (J g⁻¹ K⁻¹), ΔT is temperature change. For water, c = 4.18 J g⁻¹ K⁻¹. Calorimetry measure

Physical Chemistry

Physical Chemistry

The enthalpy change of a reaction is independent of the route taken, depending only on initial and final states. If a target reaction is the sum of known reactions, its ΔH is the sum of the individual ΔH values. Hess's law cycles (Hess cycles) manipulate thermochemical equations to find unknown enthalpy changes without direct measurement. Hess's Law states that the enthalpy change of a reaction i

Physical Chemistry

Physical Chemistry

The energy required to break one mole of a covalent bond in gaseous atoms (bond dissociation energy), typically 150-1000 kJ mol⁻¹. Bond enthalpy is always positive (breaking requires energy). Enthalpy change for a reaction: ΔH = Σ(bond enthalpies broken) - Σ(bond enthalpies formed). Average bond enthalpies account for multiple environments (C-H in different molecules) by averaging. Bond enthalpy

Physical Chemistry

Physical Chemistry

The enthalpy change (ΔHc) when one mole of a substance completely burns in excess oxygen under standard conditions (298 K, 100 kPa), measured in kJ mol⁻¹. Combustion is always exothermic, so ΔHc is negative. For example, the standard enthalpy of combustion of methane is ΔHc = -890 kJ mol⁻¹ for the reaction CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l). Standard enthalpy values allow calculation of reaction e

Physical Chemistry

Physical Chemistry

The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions (298 K, 100 kPa). Denoted ΔH°f. By definition, ΔH°f of elements in their standard states = 0. Used to calculate reaction enthalpy: ΔH°rxn = Σ(ΔH°f products) - Σ(ΔH°f reactants). Standard enthalpy of formation (ΔH°f) is ΔH for forming one mole of compound from elements in

Physical Chemistry

Physical Chemistry

A model explaining reaction mechanisms: a reaction occurs when reactant molecules collide with sufficient kinetic energy (exceeding the activation energy, Ea) and appropriate orientation. Not all collisions produce reaction—only those with energy ≥ Ea in the correct geometry react. Reaction rate increases with temperature (more energetic collisions), concentration (more frequent collisions), and u

Physical Chemistry

Physical Chemistry

The minimum kinetic energy that colliding molecules must have to react (Ea, in kJ mol⁻¹). Energy diagram shows reactants at one level, products at another (lower for exothermic, higher for endothermic), with an energy barrier (activation energy) between them. Ea is independent of whether the reaction is exothermic or endothermic; it depends only on the reaction mechanism. Activation energy is the

Physical Chemistry

Physical Chemistry

A probability distribution showing the fraction of molecules at each kinetic energy (or velocity) in a gas/liquid at a given temperature. The distribution is skewed: most molecules have near-average energy, a tail extends to very high energies. The area under the curve above Ea equals the fraction of molecules with E ≥ Ea (able to react). As temperature increases, the curve flattens and shifts rig

Physical Chemistry

Physical Chemistry

Reaction rate increases with temperature. Typically, a 10°C increase roughly doubles the rate (rate ≈ 2 × rate at T-10). Two effects: collision frequency increases slightly (~2% per °C), but the fraction of collisions with E ≥ Ea increases exponentially. The exponential factor dominates, explaining the roughly-doubling rule. Quantified by the Arrhenius equation. Temperature strongly increases rea

Physical Chemistry

Physical Chemistry

Increasing reactant concentration increases the rate of reaction (in the absence of saturating conditions). The relationship is quantified by the rate equation: rate = k[A]ⁿ[B]ᵐ, where n and m are orders of reaction (often 0, 1, or 2, determined experimentally). Doubling concentration of a first-order reactant doubles the rate; doubling a second-order reactant increases rate by 4x. Concentration

Physical Chemistry

Physical Chemistry

A substance that increases reaction rate by lowering the activation energy, without being consumed in the reaction. A catalyst speeds up both forward and reverse reactions equally (increases equilibrium approach rate, doesn't shift equilibrium position). Homogeneous catalysts (same phase as reactants, e.g., I⁻ in H₂O₂ decomposition) work by forming intermediates; heterogeneous catalysts (different

Physical Chemistry

Physical Chemistry

When a system at equilibrium is disturbed (by changing concentration, pressure, or temperature), the equilibrium shifts to counteract the disturbance and re-establish equilibrium. Increasing a reactant's concentration shifts equilibrium right (increases products); increasing pressure shifts toward the side with fewer moles of gas; increasing temperature shifts toward the endothermic direction. Le

Physical Chemistry

Physical Chemistry

A state where the forward and reverse reactions occur simultaneously at equal rates, so concentrations of reactants and products remain constant (but not necessarily equal). Equilibrium is dynamic: both reactions continue, but net composition doesn't change. Reached when forward rate = reverse rate (after sufficient time). Represented by ↔ (double arrow). In a reversible reaction aA + bB ⇌ cC + d

Physical Chemistry

Physical Chemistry

A constant (at a given temperature) expressing the ratio of product to reactant concentrations at equilibrium. For aA + bB ↔ cC + dD, Kc = [C]^c[D]^d / [A]^a[B]^b. Kc units depend on stoichiometry (often mol⁻¹ dm³ or unitless). Large Kc (>1) favors products; small Kc (<1) favors reactants. Kc changes with temperature but not with concentration, pressure, or catalysts. Kc is derived from the equil

Physical Chemistry

Physical Chemistry

Temperature changes alter Kc according to the reaction's enthalpy. For exothermic reactions (ΔH° < 0), increasing T decreases Kc (equilibrium shifts left, favoring reactants). For endothermic reactions (ΔH° > 0), increasing T increases Kc (equilibrium shifts right, favoring products). Quantified by van 't Hoff equation: ln(K₂/K₁) = -(ΔH°/R)(1/T₂ - 1/T₁). Temperature directly affects Kc (equilibri

Physical Chemistry

Physical Chemistry

Large-scale production optimizes equilibrium position using Le Chatelier's principle. Haber process (N₂ + 3H₂ ↔ 2NH₃): high pressure favors products (fewer moles); moderate T (~450°C) balances yield (higher T increases rate but decreases Kc for exothermic reaction). Contact process (2SO₂ + O₂ ↔ 2SO₃, exothermic): high pressure, moderate T (~450°C), catalyst increases rate without affecting equilib

Physical Chemistry

Physical Chemistry

A number assigned to an element in a compound that represents electrons lost, gained, or shared. Rules: (1) elements in standard state = 0; (2) monoatomic ions = charge; (3) O = -2 (except peroxides -1); (4) H = +1 (except hydrides -1); (5) Group 1 = +1, Group 2 = +2; (6) sum = zero (neutral) or charge (ions). Used to identify redox reactions (oxidation states change). Oxidation states track elec

Physical Chemistry

Physical Chemistry

Reactions involving transfer of electrons between species. Oxidation is loss of electrons or increase in oxidation state; reduction is gain of electrons or decrease in oxidation state. Oxidizing agents accept electrons (are reduced); reducing agents donate electrons (are oxidized). Redox reactions power batteries, fuel cells, and many industrial processes. Half-equations separate oxidation and red

Physical Chemistry

Physical Chemistry

Separate equations for oxidation and reduction components of a redox reaction. The oxidation half-equation shows species losing electrons; the reduction half-equation shows species gaining electrons. Half-equations must be balanced for atoms and charge. Multiplying half-equations by integers ensures equal electron transfer, then adding gives the overall balanced redox equation. Half-equations clar

Physical Chemistry

Physical Chemistry

An oxidising agent gains electrons (is reduced); a reducing agent loses electrons (is oxidized). Identified by oxidation state change: if a species' oxidation state increases, it's oxidized (by the oxidising agent); if it decreases, it's reduced (by the reducing agent). Example: in Mg + Cl₂ → MgCl₂, Mg is the reducing agent (0 → +2), Cl₂ is the oxidising agent (+0 → -1). The terms 'oxidising agen

Physical Chemistry

Physical Chemistry

A thermochemical cycle used to calculate lattice enthalpy of an ionic compound. Depicts the formation of a solid ionic compound from its elements via two routes: (1) direct formation (ΔH°f), (2) via gaseous ions. Both routes have the same overall enthalpy change (Hess's law), allowing calculation of unmeasurable lattice enthalpy. Born-Haber cycle is a thermochemical cycle combining measured entha

Physical Chemistry

Physical Chemistry

The enthalpy change required to convert one mole of solid ionic compound into gaseous ions: MX(s) → M⁺(g) + X⁻(g), ΔH(lattice). Always positive (energy required to break the ionic lattice). Measured via Born-Haber cycles (cannot be directly calorimetered). Depends on ionic charge and size (higher charge, smaller ions → larger lattice enthalpy). Lattice enthalpy (enthalpy of lattice formation) is

Physical Chemistry

Physical Chemistry

The enthalpy change when one mole of gaseous ions dissolves in water to form an aqueous ion: M⁺(g) + aq → M⁺(aq), ΔH(hydration). Always negative (energy released when water molecules surround ions). Related to lattice enthalpy: solubility depends on competition between lattice enthalpy (positive, unfavourable) and hydration enthalpy (negative, favourable). When an ionic solid dissolves, the latti

Physical Chemistry

Physical Chemistry

The change in disorder or randomness in a system, denoted ΔS (units: J K⁻¹ mol⁻¹). Positive ΔS indicates increased disorder (favored); negative ΔS indicates decreased disorder (unfavored). Entropy increases when solids dissolve, liquids evaporate, or gases expand. Temperature affects entropy: higher temperature increases molecular motion and disorder. Entropy change is crucial for predicting spont

Physical Chemistry

Physical Chemistry

The thermodynamic quantity determining spontaneity of reactions: ΔG = ΔH − TΔS (units: kJ mol⁻¹). Negative ΔG indicates spontaneity (reaction occurs); positive ΔG indicates non-spontaneity (reaction does not occur under standard conditions). ΔG = 0 at equilibrium. Temperature, enthalpy, and entropy all influence whether a reaction is spontaneous. Gibbs free energy allows prediction of reaction fea

Physical Chemistry

Physical Chemistry

Thermodynamic feasibility determined by ΔG < 0 (spontaneous, will proceed forward). Kinetic feasibility determined by activation energy and reaction rate (some feasible reactions are too slow to observe on reasonable timescales). A reaction can be thermodynamically favorable (ΔG < 0) but kinetically slow (high Ea); conversely, a reaction can be thermodynamically unfavorable (ΔG > 0) but proceeding

Physical Chemistry

Physical Chemistry

An equation expressing reaction rate as a function of reactant concentrations raised to powers (orders). General form: rate = k[A]ⁿ[B]ᵐ, where k is rate constant, [A] and [B] are concentrations, and n and m are orders. Orders are determined experimentally, not from stoichiometry. Overall order (n + m) indicates how many concentration changes double the rate. Rate equations explain why some reactio

Physical Chemistry

Physical Chemistry

The power to which a reactant's concentration is raised in the rate equation. For rate = k[A]ⁿ[B]ᵐ, n is the order with respect to A, m is the order with respect to B. Orders are zero (rate independent of concentration), first (rate doubles when concentration doubles), or second (rate quadruples when concentration doubles). Overall order is n + m. Orders must be determined experimentally, not from

Physical Chemistry

Physical Chemistry

The constant k in the rate equation rate = k[A]ⁿ[B]ᵐ. Units depend on overall reaction order: for zero-order reactions, k has units mol dm⁻³ s⁻¹; for first-order, s⁻¹; for second-order, mol⁻¹ dm³ s⁻¹. Rate constant increases with temperature following the Arrhenius equation k = Ae^(-Ea/RT). A catalyst lowers activation energy and increases k without being consumed. k determines reaction speed at a

Physical Chemistry

Physical Chemistry

The slowest step in a multi-step reaction mechanism. The overall reaction rate equals the rate of the rate-determining step. Elementary steps faster than the rate-determining step don't contribute to rate-limiting. The rate-determining step often involves the reactants (or intermediates from earlier steps), explaining why the experimental rate equation often resembles the rate-determining step's s

Physical Chemistry

Physical Chemistry

An equation describing the temperature-dependence of the rate constant: k = Ae^(-Ea/RT), where A is the pre-exponential factor, Ea is activation energy (J mol⁻¹), R is the gas constant (8.31 J K⁻¹ mol⁻¹), T is absolute temperature. Linear form: ln(k) = ln(A) - (Ea/R)(1/T). Used to calculate k at different temperatures or to find Ea from rate constant data. Arrhenius equation k = Ae^(−Ea/RT) quant

Physical Chemistry

Physical Chemistry

In a mixture of ideal gases, the partial pressure of a gas is the pressure that gas would exert if it alone occupied the container at the same temperature. Dalton's Law: total pressure = sum of partial pressures. For gas i in a mixture, Pi = xi × Ptotal (where xi is mole fraction). Partial pressures explain gas behavior in mixtures: calculating equilibrium constants for gases uses partial pressure

Physical Chemistry

Physical Chemistry

The mole fraction (xi) of component i in a mixture is ni / ntotal (moles of i divided by total moles). Mole fractions are dimensionless, ranging 0 to 1, and sum to 1 for all components. Mole fractions relate to partial pressure: Pi = xi × Ptotal (for ideal gases). Mole fractions are useful for non-ideal solutions and gas mixtures where concentrations aren't relevant. Mole fraction is a compositio

Physical Chemistry

Physical Chemistry

For gases in equilibrium, Kp expresses the ratio of partial pressures (in Pa) at equilibrium. For aA(g) + bB(g) ↔ cC(g) + dD(g): Kp = (Pc^c × Pd^d) / (Pa^a × Pb^b) (units depend on stoichiometry, often Pa^Δn). Like Kc (concentration-based), Kp depends only on temperature. Related by: Kp = Kc(RT)^Δn, where Δn = moles products - moles reactants (gaseous only). Kp uses partial pressures instead of c

Physical Chemistry

Physical Chemistry

Changes in pressure, volume, or temperature affect Kp (equilibrium constant in terms of partial pressures). Temperature changes Kp: for endothermic reactions, increasing T increases Kp (shifts equilibrium right); for exothermic reactions, increasing T decreases Kp (shifts left). Pressure and volume changes don't alter Kp value itself, but do shift equilibrium position (Le Chatelier). Catalysts don

Physical Chemistry

Physical Chemistry

The potential (in volts) of a half-cell under standard conditions (298 K, 100 kPa, 1 M solutions) relative to the standard hydrogen electrode (SHE). Denoted E°. Reduction potentials (for reduction reactions) are tabulated. Positive E° indicates a good oxidising agent (reduction is favourable); negative E° indicates a good reducing agent (oxidation is favourable). Electrode potential measures the

Physical Chemistry

Physical Chemistry

A reference electrode with assigned E° = 0 V by definition. Consists of Pt electrode in contact with H₂(g, 1 atm, ~1 bar) and H⁺(aq, 1 M) at 298 K. The half-reaction: 2H⁺ + 2e⁻ ↔ H₂. All other electrode potentials are measured relative to SHE by constructing a galvanic cell with the SHE as one half-cell. Standard hydrogen electrode (SHE, 2H⁺(aq) + 2e⁻ → H₂(g) at 298 K, 100 kPa, [H⁺] = 1 M) is def

Physical Chemistry

Physical Chemistry

A device in which redox reactions occur at two electrodes (anode and cathode) separated by an electrolyte. In a galvanic (voltaic) cell, spontaneous redox reactions produce electrical current (ΔG < 0, E° > 0). In an electrolytic cell, applied external voltage drives non-spontaneous redox reactions. Cell notation: anode | electrolyte | cathode; double line | | represents a salt bridge (in galvanic

Physical Chemistry

Physical Chemistry

The electromotive force (voltage) of a galvanic cell: E°cell = E°(cathode) - E°(anode) under standard conditions. Positive E°cell indicates a spontaneous reaction (ΔG < 0). Related to ΔG° by: ΔG° = -nFE°, where n is moles of electrons, F is Faraday's constant (96,500 C mol⁻¹). Non-standard conditions: Ecell = E°cell - (RT/nF) ln(Q) (Nernst equation). EMF is the driving force for electron flow in

Physical Chemistry

Physical Chemistry

Electrochemical cells generating electricity from continuous supply of fuel (e.g., H₂) and oxidant (e.g., O₂). More efficient than combustion (not limited by Carnot efficiency). The hydrogen fuel cell: anode: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻ (in alkaline), cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻. Overall: 2H₂ + O₂ → 2H₂O. Clean (product is water), quiet, efficient (~60% theoretical max). Fuel cells are galvanic c

Physical Chemistry

Physical Chemistry

An acid is a proton (H⁺) donor; a base is a proton acceptor. An acid-base reaction is proton transfer: HA + B → A⁻ + HB⁺. The conjugate base of an acid (A⁻) is formed by removing a proton; the conjugate acid of a base (HB⁺) is formed by adding a proton. A conjugate acid-base pair differs by one proton. Bronsted-Lowry theory extends acid-base chemistry beyond aqueous solutions and includes reactio

Physical Chemistry

Physical Chemistry

Strong acids (HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄) completely dissociate in water: HA → H⁺ + A⁻ (100% ionisation). Weak acids (e.g., CH₃COOH, HCN) partially dissociate: HA ↔ H⁺ + A⁻, equilibrium favours reactants. Strength is quantified by Ka (acid dissociation constant). Strong acids have Ka >> 1; weak acids have Ka << 1. Strong acids dissociate completely, so [H⁺] from a strong acid = initial conc

Physical Chemistry

Physical Chemistry

pH = -log₁₀[H⁺], where [H⁺] is hydrogen ion concentration in mol dm⁻³. Neutral solutions (at 298 K): pH = 7. Acidic: pH < 7 ([H⁺] > 10⁻⁷). Basic: pH > 7 ([H⁺] < 10⁻⁷). Related to pOH by: pH + pOH = 14 (at 298 K). Converting between pH and [H⁺]: [H⁺] = 10^(-pH). pH = −log[H⁺] expresses hydrogen ion concentration on logarithmic scale (manageable range 0−14 instead of 10⁻¹⁴ to 10⁻¹). [H⁺] from pH: [

Physical Chemistry

Physical Chemistry

The water ion product Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C (units: mol² dm⁻⁶). Kw is the equilibrium constant for water dissociation: H₂O ⇌ H⁺ + OH⁻. In pure water, [H⁺] = [OH⁻] = 10⁻⁷ M (neutral, pH = 7). Adding acid increases [H⁺], shifting equilibrium left and decreasing [OH⁻] proportionally (product Kw remains constant). Kw enables calculating [OH⁻] from [H⁺]: [OH⁻] = Kw / [H⁺]. Kw increases w

Physical Chemistry

Physical Chemistry

For a weak acid HA ↔ H⁺ + A⁻, Ka = [H⁺][A⁻]/[HA] (mol dm⁻³). Larger Ka indicates stronger acid (more dissociation). Tabulated values at 298 K. Related to strength: strong acids have Ka >> 1 (often>10³); weak acids have Ka < 1. pKa = -log Ka; lower pKa = stronger acid. Ka = [H⁺][A⁻] / [HA] quantifies weak acid strength (HA ⇌ H⁺ + A⁻). Large Ka (weak acid is strong, more dissociation): HCl (very la

Physical Chemistry

Physical Chemistry

Solutions that resist pH changes when small amounts of acid or base are added. Composed of a weak acid and its conjugate base (e.g., CH₃COOH + CH₃COONa) or a weak base and its conjugate acid (e.g., NH₃ + NH₄Cl). Buffer capacity depends on the concentrations of the weak acid and conjugate base. pH is given by Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]). Buffers work by neutralizing a

Physical Chemistry

Physical Chemistry

A pH curve (titration curve) plots pH vs. volume of titrant during an acid-base titration. Shapes differ based on acid/base strengths: strong acid/strong base (sharp vertical at equivalence point), weak acid/strong base (curved, steep portion before equivalence point, buffer region), weak base/strong acid (curved, buffer region before equivalence point). Indicators are weak acids or bases that cha

Physical Chemistry

Physical Chemistry

Quantitative analysis using acid-base titrations. At equivalence point: moles of acid = moles of base (accounting for stoichiometry). For strong acid/strong base: use concentration × volume (n = cV) directly. For weak acids/bases, use the same stoichiometric principle, but pH at equivalence point differs. Standardisation: a standard solution (precisely known concentration) is used as the titrant.

Physical Chemistry