The ultimate interactive encyclopedia — organic, functional groups, equilibrium, bonding, nuclear, electrochemistry, and 9 live calculation tools.
Hydrocarbons are organic compounds consisting entirely of Carbon (C) and Hydrogen (H) atoms — nothing else. They form the backbone of fossil fuels, polymers, and countless synthetic materials. Understanding hydrocarbons is the foundation of all organic chemistry.
Contain the maximum possible hydrogen atoms. All C–C bonds are single bonds (σ bonds only) — no π bonds. This makes them chemically stable and unreactive at room temperature.
Formula: CnH2n+2
React only in: combustion (burning) and halogenation (substitution with Cl₂/Br₂ under UV light). Both require breaking the strong C–H bond first.
Contain double (alkenes) or triple (alkynes) C–C bonds. The π bond electrons are above and below the bond axis — they are loosely held and highly reactive. These π electrons are attacked by electrophiles in addition reactions.
Alkene: CnH2n · Alkyne: CnH2n−2
React with: H₂ (hydrogenation), Br₂ (bromine water test), HX (hydrohalogenation), H₂O (hydration to alcohol).
Both alkenes and alkynes cannot exist at n=1. A double or triple bond requires at least two carbon atoms — you need two carbons to share a double bond between them.
n=1: Methane (CH₄) only — no ethene or ethyne equivalent
n=2: Ethene (C₂H₄) and Ethyne (C₂H₂) are the minimum
As the carbon chain gets longer, boiling point increases. More carbons = larger surface area = stronger van der Waals (dispersion) forces between molecules = more energy needed to separate them.
Branched isomers have lower boiling points than straight-chain isomers. Branching makes the molecule more spherical, reducing surface area contact between molecules, so less energy is needed to separate them.
All three are C₅H₁₂ but have very different boiling points.
All hydrocarbons are non-polar — they do not mix with water ("like dissolves like"). They dissolve freely in non-polar organic solvents.
This is why crude oil is separated by fractional distillation — each fraction boils at a different temperature range.
All hydrocarbons burn in oxygen. Complete combustion (excess O₂) gives only CO₂ and H₂O. Incomplete combustion (limited O₂) gives CO and/or carbon soot — both toxic/polluting.
C₄H₁₀ + 6½O₂ → 4CO₂ + 5H₂O (complete)
2C₄H₁₀ + 9O₂ → 8CO + 10H₂O (incomplete)
A quick test to distinguish saturated from unsaturated hydrocarbons:
The degree of unsaturation (DoU) tells you how many double bonds or rings are in a molecule. Calculate it from the molecular formula CnHm:
DoU = (2n + 2 − m) ÷ 2
DoU=0 → alkane · DoU=1 → one ring or one C=C · DoU=2 → one C≡C or two C=C · DoU=4 → benzene ring
Enter any value of n to instantly see the molecular formula, compound name, and series rules for all three hydrocarbon families.
IUPAC rules ensure every molecule has one unique, unambiguous systematic name understood by chemists worldwide. Follow these five steps in order — every time.
Look at the whole molecule and trace the longest continuous chain of carbon atoms — this is your parent chain and gives the base name.
Once you have the parent chain, number its carbons from one end to the other. You must choose the direction that gives the lowest set of locants (position numbers).
Any group attached to the main chain (but not part of it) is a substituent. These are named using the carbon count of the branch + the suffix "-yl".
When two or more different substituents are present, list them in the name in strict alphabetical order based on the substituent name alone.
You have all the pieces. Now assemble them in this exact format:
2,3- not 2 3-2-methyl not 2 methylbut-2-ene or pent-1-yne
My Excellent Puppy Bites People
Happily Helping Old Neighbours Daily
Meth · Eth · Prop · But · Pent · Hex · Hept · Oct · Non · Dec
Isomers share the exact same molecular formula but differ in how their atoms are arranged — either in bonding order or in 3D space. The same atoms, rearranged, can produce dramatically different chemicals with different smells, melting points, and biological effects.
Atoms are bonded in a different sequence. They have the same molecular formula but different structural (displayed) formulas. Three sub-types:
Same bonding sequence, but atoms are arranged differently in 3D space. They cannot be interconverted without breaking bonds. Two types:
A C=C bond consists of a σ (sigma) bond AND a π (pi) bond. The π bond locks the two carbons in place — they cannot rotate freely. This is why groups on either side of a double bond are fixed in space, making cis/trans isomers stable and distinct compounds.
Single bonds (C–C) rotate freely → no geometric isomers possible.
Both conditions must be true:
CH₂=CH₂ → NO (each C has 2×H)
CH₃–CH=CH–CH₃ → YES (each C has H and CH₃)
Cis-but-2-ene and trans-but-2-ene are genuinely different compounds:
A chiral carbon (asymmetric carbon) is a carbon atom bonded to four completely different atoms or groups. It is marked with an asterisk: C*.
Example: CH₃–C*H(OH)–COOH
(Lactic acid — the C* has: H, OH, CH₃, COOH)
The two mirror-image arrangements cannot be superimposed on each other, just like a left and right hand look identical in a mirror but are not the same object.
Many drug molecules contain chiral centres. The two enantiomers can have very different effects in the body because biological receptors are also chiral and only fit one "hand" of the molecule.
Maximum number of optical isomers = 2ⁿ where n = number of chiral centres.
Note: sometimes internal symmetry reduces this number (meso compounds).
Step 1: Start with the straight chain: CH₃–CH₂–CH₂–CH₃ → n-Butane
Step 2: Try branching — take one CH₃ off the chain and attach it to C2: CH₃–CH(CH₃)–CH₃ → 2-Methylpropane (Isobutane)
Step 3: Any more unique structures? No — any other arrangement repeats one of the above.
✅ C₄H₁₀ has exactly 2 structural isomers.
C₅H₁₂ has 3 isomers · C₆H₁₄ has 5 isomers · C₇H₁₆ has 9 isomers — branching possibilities grow rapidly!
C₅H₁₂ alone has 3 structural isomers with different boiling points and reactivities. In pharmaceuticals, the wrong isomer can be the difference between a medicine and a poison — Thalidomide being the famous tragic example. This is why modern drug synthesis must be stereospecific — producing only the correct enantiomer.
Pick a compound type and carbon count to instantly see every structural isomer with IUPAC names and condensed formulas.
Enter a carbon count and instantly generate every possible structural isomer of that alkane — complete with IUPAC names, molecular formulas, and condensed structures. Powered by a recursive carbon skeleton algorithm.
Complete homologous series of the first ten members with systematic IUPAC names.
| n | Prefix | Alkane | Formula | Alkene | Formula | Alkyne | Formula |
|---|
Functional groups are specific atom arrangements within organic molecules that determine their characteristic chemical reactivity — they are the "active sites" of organic chemistry. Knowing them lets you predict how any organic compound will behave.
Hydroxyl group. Polar, forms H-bonds → high boiling points. Reacts with acids to form esters (esterification).
Terminal carbonyl (C=O). Easily oxidised to carboxylic acid. Gives positive Tollens' test (silver mirror).
Internal carbonyl flanked by two C-groups. Cannot be easily oxidised. Negative Tollens' test. Used as solvents.
Both carbonyl and hydroxyl. Weak acid, donates H⁺. Reacts with alcohols (esterification) and bases (neutralisation).
Formed from acid + alcohol (esterification, loses H₂O). Fruity odours, used in flavourings. Hydrolysis reverses to acid + alcohol.
Nitrogen with lone pair. Acts as a base — accepts H⁺. Fishy odour. Essential in amino acids, proteins, and drugs.
Oxygen bridging two C-chains. Relatively unreactive, non-polar. Excellent organic solvents. Named as alkoxy derivatives.
Carbonyl bonded to nitrogen. Very stable — the peptide bond in proteins is an amide bond. Used in nylon and paracetamol.
Halogen attached to a carbon chain. Undergoes nucleophilic substitution (SN1/SN2) and elimination reactions. Used as solvents and intermediates.
Sulfur analogue of alcohol. Very strong odour (responsible for garlic and skunk smell). Oxidised to disulfide bonds in proteins.
Triple bond between carbon and nitrogen. Polar and high-boiling. Hydrolysed to carboxylic acids; reduced to amines. Used in pharmaceuticals.
Cyclic, planar, delocalised π electrons satisfying Hückel's rule (4n+2 π e⁻). Undergoes electrophilic aromatic substitution (EAS) rather than addition.
Sulfur bridging two carbon chains — the sulfur analogue of an ether. Less reactive than thiols. Oxidised to sulfoxides and sulfones. Found in methionine (amino acid).
Two acyl groups joined via an oxygen. Formed by dehydration of two carboxylic acid molecules. More reactive than esters — reacts readily with alcohols, amines, and water.
The most reactive acyl derivative — carbonyl directly bonded to a halogen. Reacts vigorously with water, alcohols, and amines. Used to introduce acyl groups in synthesis.
Carbon's oxidation state increases:
Alkane → Alcohol → Aldehyde → Carboxylic Acid
Each step adds oxygen or removes hydrogen. Ketones cannot be further oxidised without breaking C–C bonds.
Organic reactions transform one functional group into another. Knowing the mechanism — which bonds break and form — lets you design synthesis routes and predict products.
Alkenes/alkynes add atoms across the double/triple bond. The π bond breaks, two new σ bonds form.
Alkanes react with halogens (UV light). A H atom is replaced by a halogen atom via radical chain mechanism.
Removal of a small molecule (H₂O or HX) from adjacent carbons to form a double bond.
A carboxylic acid reacts with an alcohol under acid catalyst to form an ester and water (condensation). Reversible equilibrium.
Oxidising agents (K₂Cr₂O₇/H₂SO₄ or KMnO₄) selectively oxidise alcohols. Primary → Aldehyde → Carboxylic Acid. Secondary → Ketone.
Organic molecules + O₂ → releases energy. Complete combustion gives CO₂ + H₂O. Limited O₂ gives CO or C (soot).
Regioselectivity: where reaction occurs (Markovnikov)
Stereoselectivity: which face is attacked (syn vs anti addition)
Chemoselectivity: which functional group reacts first
Polymers are giant molecules (macromolecules) formed by joining thousands of small monomer units. They dominate modern materials science — from packaging to DNA to spider silk.
Monomers with C=C double bonds open their bonds to join. No atoms are lost — the monomer formula equals the repeat unit formula. Requires initiator (free radical or catalyst).
Two different functional groups react and release a small molecule (usually H₂O or HCl) with each bond formed. Monomers need two reactive groups each (bifunctional).
Benzene's delocalized π system drives a unique preference for substitution over addition. This section covers aromaticity, EAS and NAS mechanisms, directing effects, preparation of benzene, and alkyne chemistry.
Six carbon atoms in a planar ring, each bonded to one hydrogen. The π electrons are fully delocalized across all six carbons — not fixed alternating double bonds.
Bond lengths are all equal at 1.40 Å — intermediate between a single (1.54 Å) and double (1.34 Å) bond.
A cyclic, planar, fully conjugated molecule is aromatic if it has 4n + 2 π electrons (n = 0, 1, 2…). Benzene has 6 π electrons (n=1) → aromatic ✅. This extraordinary stability is why benzene strongly prefers substitution over addition.
H₂SO₄ protonates HNO₃ to generate the nitronium ion (NO₂⁺), the electrophile that attacks the ring.
The Lewis acid catalyst polarises the halogen molecule, generating a powerful electrophile (Cl⁺ equivalent) that can attack the ring.
The electrophile is SO₃. Unlike other EAS reactions, sulfonation is reversible — the sulfonic acid group can be removed with steam.
AlCl₃ generates a carbocation (R⁺) electrophile. Prone to polyalkylation since the product is more reactive than benzene.
Adding across a double bond breaks the conjugated π system. The molecule loses ~150 kJ/mol of resonance stabilization — a huge energy cost. Substitution avoids this by restoring the ring after attack.
Benzene can undergo addition under forcing conditions — e.g., H₂ / Ni catalyst at high pressure and temperature — yielding cyclohexane. But this requires far harsher conditions than a simple alkene.
| Substituent | Type | Directs to | Effect on rate | Example |
|---|---|---|---|---|
| —OH, —OR | Electron-donating | ortho / para | Activates (faster) | Phenol |
| —NH₂, —NHR | Electron-donating | ortho / para | Activates (much faster) | Aniline |
| —CH₃, alkyl | Electron-donating (inductive) | ortho / para | Activates (slightly) | Toluene |
| —NO₂ | Electron-withdrawing | meta | Deactivates (slower) | Nitrobenzene |
| —COOH, —CHO | Electron-withdrawing | meta | Deactivates | Benzoic acid |
| —SO₃H | Electron-withdrawing | meta | Deactivates | Benzenesulfonic acid |
| —X (halogens) | Mixed (–I inductive, +M mesomeric) | ortho / para | Deactivates (ring) but o/p director | Chlorobenzene |
Naphtha fractions undergo dehydrogenation over a platinum-alumina catalyst. Accounts for the majority of global benzene production.
When coal is heated without air, a complex coal tar mixture forms. Fractional distillation yields benzene, toluene, and xylenes. The historic original source of benzene (Faraday, 1825).
The carboxyl group (—COOH) is removed as CO₂ in the presence of soda lime. A classic lab method — benzoic acid is cheap and readily available.
An aryl halide treated with an alkyl halide and sodium metal forms an arene. A modification of the Wurtz reaction involving aryl radicals or carbanions.
Phenol is reduced to benzene by passing its vapour over heated zinc dust. The —OH group is removed and replaced by —H. Zinc is oxidised to ZnO.
Two aryl halides couple using sodium metal, forming a biaryl (biphenyl). Important for making biphenyl used in pharmaceuticals and liquid crystals.
The nucleophile first adds to the ring forming a carbanion intermediate called the Meisenheimer complex. The leaving group then departs, restoring aromaticity. Fluorine is the best leaving group in NAS despite being a poor one in SN2.
The two nitro groups at positions 2 and 4 strongly activate the ring toward nucleophilic attack at C1. Hydroxide displaces chloride via the Meisenheimer complex.
Carbon atoms in alkynes are sp hybridised — two sp orbitals form σ bonds (linear, 180°), leaving two unhybridised p orbitals each that overlap to form two π bonds alongside the σ bond: a triple bond total.
The triple bond is shorter (1.20 Å) and stronger (~839 kJ/mol) than a double bond. The linear geometry means all four atoms H–C≡C–H are collinear.
Alkynes follow the general formula CₙH₂ₙ₋₂. The simplest is ethyne (acetylene), C₂H₂. Terminal alkynes have the triple bond at the end of the chain (—C≡CH); internal alkynes have it in the middle (R—C≡C—R'). Terminal alkynes are weakly acidic due to the sp C—H bond.
The sp C—H bond in terminal alkynes (R–C≡C–H) has more s-character (50%) than sp² (33%) or sp³ (25%). Greater s-character means the electrons are held closer to the nucleus → the bond is more polar → H is more easily lost as H⁺.
pKₐ ≈ 25 (terminal alkyne) vs 44 (alkane) vs 10 (phenol) vs −1 (HCl)
RC≡CH + NaNH₂ → RC≡CNa + NH₃
2RC≡CH + 2Na → 2RC≡CNa + H₂↑
Terminal alkynes react with sodamide (NaNH₂) or sodium metal to form sodium acetylides (alkynide salts). These are strong nucleophiles — useful for extending carbon chains via SN2 on haloalkanes.
Full hydrogenation over Pd/C gives the alkane. Lindlar's catalyst (poisoned Pd) stops at the cis-alkene. Birch reduction (Na/liq. NH₃) gives the trans-alkene instead.
Bromine adds across the triple bond in two steps. First equivalent gives a dibromoalkene (trans, anti addition); second gives the tetrabromoalkane. Decolourisation is a positive test for unsaturation.
Water adds across the triple bond following Markovnikov's rule via an enol intermediate. The enol immediately tautomerises to the more stable ketone. Ethyne gives acetaldehyde since both carbons are equivalent.
Three molecules of ethyne trimerize to form benzene at high temperature. One of the first synthetic routes to benzene — a conceptual link between alkynes and aromatic chemistry.
Hydrocarbon derivatives are organic compounds formed by replacing one or more hydrogen atoms with functional groups. Each class has characteristic reactions governed by the electronic nature of the functional group.
| Class | Functional Group | General Formula | Example | Key Reaction Type |
|---|---|---|---|---|
| Haloalkanes | —X (F, Cl, Br, I) | R—X | CH₃Cl | Nucleophilic substitution |
| Alcohols | —OH | R—OH | CH₃CH₂OH | Oxidation, esterification |
| Ethers | —O— | R—O—R' | (CH₃)₂O | Relatively inert |
| Aldehydes | —CHO | RCHO | HCHO (methanal) | Nucleophilic addition |
| Ketones | C=O | RCOR' | CH₃COCH₃ | Nucleophilic addition |
| Carboxylic Acids | —COOH | RCOOH | CH₃COOH | Esterification, decarboxylation |
| Esters | —COO— | RCOOR' | CH₃COOC₂H₅ | Hydrolysis, transesterification |
| Amines | —NH₂ | RNH₂ | CH₃NH₂ | Nucleophilic, basic reactions |
| Amides | —CONH₂ | RCONH₂ | CH₃CONH₂ | Hydrolysis |
| Nitriles | —C≡N | RCN | CH₃CN | Hydrolysis, reduction |
Alcohols react with hydrogen halides to form haloalkanes. SOCl₂ (thionyl chloride) is preferred for chlorides as it gives clean reaction with only gaseous by-products (SO₂ + HCl).
One-step backside attack — nucleophile attacks from opposite side to leaving group. Results in inversion of configuration (Walden inversion). Rate = k[RX][Nu⁻]. Favoured by primary substrates.
Two-step mechanism: rate-determining ionisation forms a carbocation, then rapid nucleophile attack. Rate = k[RX] only. Gives racemic mixture (planar carbocation attacked from both sides).
A base removes a β-hydrogen while the halide leaves, forming an alkene. Hot alcoholic KOH favours elimination; cold aqueous KOH favours substitution. The more substituted alkene is the major product (Zaitsev's rule).
Primary alcohols oxidise to aldehydes (with limited oxidant) or carboxylic acids (excess oxidant). Secondary alcohols oxidise to ketones and stop. Tertiary alcohols cannot be oxidised without breaking C—C bonds.
Acid-catalysed reversible reaction. Yield improved by using excess of one reagent, removing water, or using a dehydrating agent. The oxygen in the ester comes from the alcohol (shown by isotope labelling).
At high temperature, conc. H₂SO₄ acts as a dehydrating agent, removing water across adjacent C—H and C—OH bonds to give an alkene. At lower temperature (140 °C), intermolecular dehydration gives an ether.
Alcohols react with sodium (less vigorously than water) releasing hydrogen gas and forming a sodium alkoxide (RONa). This shows alcohols are weakly acidic. Reactivity: 1° > 2° > 3° for this reaction.
CN⁻ attacks the electrophilic carbonyl carbon, followed by protonation. Extends the carbon chain by one — useful in synthesis. Aldehydes react faster than ketones due to less steric hindrance.
Hydride (H⁻) is delivered to the carbonyl carbon. Aldehydes give primary alcohols; ketones give secondary alcohols. LiAlH₄ also reduces carboxylic acids and esters; NaBH₄ is more selective.
Aldehydes and ketones react with 2,4-DNPH to form an orange or yellow crystalline hydrazone precipitate. A positive test confirms a C=O group. The melting point of the precipitate can identify the specific compound.
Aldehydes are oxidised by Tollens' reagent (silver mirror) and Fehling's solution (red precipitate). Ketones are not oxidised — key distinction. Fehling's also distinguishes aliphatic from aromatic aldehydes.
React with metals (H₂↑), carbonates (CO₂↑), and bases to form salts. The carboxylate ion (RCOO⁻) is stabilised by resonance — negative charge spread over both oxygens.
Ethanoic acid + ethanol → ethyl ethanoate (fruity smell). Reversible — equilibrium mixture. Yield improved by excess alcohol, removing water, or adding drying agent.
Carboxylic acids are harder to reduce than aldehydes/ketones. LiAlH₄ directly gives the primary alcohol in one step. NaBH₄ alone is insufficient for the —COOH group.
The carboxyl group is removed as CO₂, shortening the carbon chain by one. Classic lab: sodium ethanoate + soda lime → methane. Used industrially to produce simpler hydrocarbons from fatty acid salts.
Acid hydrolysis is the exact reverse of esterification. Equilibrium mixture — excess water pushes it toward hydrolysis. Gives carboxylic acid + alcohol.
Base hydrolysis is irreversible — the carboxylate salt (soap) formed is stable and does not react back with the alcohol. This is the basis of saponification (soap making from fats/oils + NaOH).
Amines are basic due to the lone pair on nitrogen. Alkyl groups (electron-donating) increase basicity compared to NH₃. Aromatic amines (aniline) are much weaker bases — the lone pair is delocalised into the ring.
Nitriles are reduced to primary amines, extending the carbon chain. LiAlH₄ in ether is the lab method; catalytic hydrogenation is used industrially.
Amines (as nucleophiles) attack acyl chlorides to form amides. Very fast, exothermic reaction. Used in pharmaceutical synthesis (forming peptide-like bonds).
Primary aromatic amines form stable diazonium salts at 0–5 °C. These are versatile intermediates — coupling with phenols/amines gives azo dyes; Sandmeyer reactions give halides, —CN, —OH.
The amide bond (—CO—NH—) is the basis of peptide bonds in proteins and nylon. Acyl chlorides react immediately with ammonia — much faster route than heating ammonium salts.
Amides hydrolyse slowly under reflux. Base hydrolysis gives carboxylate salt + ammonia gas — which turns damp red litmus blue (useful identification test).
CN⁻ is an excellent nucleophile — displaces halide via SN2, adding one carbon to the chain. Important synthetic step for making longer-chain carboxylic acids or amines.
Nitriles hydrolyse to amides, then to carboxylic acids. Gives a carboxylic acid with one more carbon than the starting alcohol/halide — useful chain extension in synthesis.
Cyclic hydrocarbons are ring-shaped molecules where carbon atoms form a closed loop. They are divided into alicyclic (non-aromatic) and aromatic classes. Ring size and saturation determine stability, physical properties, and reactivity.
Cyclohexane rapidly interconverts between chair conformations at room temperature. Large substituents (e.g. —C(CH₃)₃) prefer the equatorial position to avoid 1,3-diaxial steric clashes.
| Ring | Formula | Bond Angle | Deviation from sp³ (109.5°) | Stability | Notes |
|---|---|---|---|---|---|
| Cyclopropane (3) | C₃H₆ | 60° | −49.5° | Very unstable | Banana bonds, reacts like alkene |
| Cyclobutane (4) | C₄H₈ | 90° | −19.5° | Unstable | Puckered to reduce torsional strain |
| Cyclopentane (5) | C₅H₁₀ | 108° | −1.5° | Nearly strain-free | Envelope conformation |
| Cyclohexane (6) | C₆H₁₂ | 111° | +1.5° | Strain-free (chair) | Chair conformation; most stable |
| Cycloheptane (7) | C₇H₁₄ | 128° | +18.5° | Moderate strain | Torsional + transannular strain |
| Cyclooctane (8) | C₈H₁₆ | 135° | +25.5° | Medium ring strain | Transannular H–H repulsion |
Cyclopropane's extreme angle strain makes it react like an alkene — the ring opens via electrophilic addition. Bromine water is decolourised. Larger rings do not undergo ring-opening addition.
Cycloalkanes burn in excess oxygen to give CO₂ and H₂O. The enthalpy of combustion per CH₂ unit is slightly higher for strained rings — direct calorimetric evidence for ring strain.
Unstrained cycloalkanes (C5+) undergo free-radical substitution rather than ring-opening addition — just like open-chain alkanes. All H atoms are equivalent in cyclohexane, so only one monosubstituted product forms.
Cycloalkenes undergo catalytic hydrogenation at the double bond to give cycloalkanes — ring remains intact. Cyclohexene → cyclohexane. Heat of hydrogenation is higher for strained rings, reflecting stored ring energy.
| Type | General Formula | Degrees of Unsaturation | Typical Reactions | Example |
|---|---|---|---|---|
| Cycloalkane | CₙH₂ₙ | 1 (ring only) | Substitution (radical), combustion | Cyclohexane C₆H₁₂ |
| Cycloalkene | CₙH₂ₙ₋₂ | 2 (ring + one C=C) | Electrophilic addition, hydrogenation | Cyclohexene C₆H₁₀ |
| Cycloalkadiene | CₙH₂ₙ₋₄ | 3 (ring + two C=C) | Addition, Diels-Alder (conjugated) | 1,3-Cyclohexadiene C₆H₈ |
| Arene (benzene) | CₙH₂ₙ₋₆ | 4 (ring + 3 delocalised) | EAS (substitution preferred) | Benzene C₆H₆ |
| Small ring (strained) | CₙH₂ₙ (n=3,4) | 1 | Ring-opening addition (like alkenes) | Cyclopropane C₃H₆ |
Alcohols contain —OH bonded to a saturated sp³ carbon; phenols have —OH bonded directly to a benzene ring. This structural difference dramatically changes acidity, reactivity, and chemical behaviour.
The hydroxyl group (—OH) is the defining functional group. Oxygen is sp³ hybridised with two lone pairs, making the C—O—H angle ~108.5°.
Hydrogen bonding between —OH groups raises the boiling point dramatically: ethanol bp = 78 °C vs ethane bp = −89 °C (similar MW, no H-bonding).
| Molecular Formula | Structural Formula | IUPAC Name | Common Name | Notes |
|---|---|---|---|---|
| CH₄O | CH₃—OH | Methanol | Wood alcohol | Simplest alcohol, toxic |
| C₂H₆O | CH₃CH₂—OH | Ethanol | Grain alcohol | Beverages, fuel, solvent |
| C₃H₈O | CH₃CH₂CH₂—OH | Propan-1-ol | n-Propanol | —OH on C1 |
| C₃H₈O | CH₃CH(OH)CH₃ | Propan-2-ol | Isopropanol (IPA) | —OH on C2, secondary |
| C₄H₁₀O | (CH₃)₃C—OH | 2-Methylpropan-2-ol | tert-Butanol | Tertiary, branched |
| C₂H₆O₂ | HO—CH₂CH₂—OH | Ethane-1,2-diol | Ethylene glycol | Dihydric, antifreeze |
| C₃H₈O₃ | HOCH₂CH(OH)CH₂OH | Propane-1,2,3-triol | Glycerol | Trihydric, in fats/oils |
Alcohols react with sodium to release hydrogen gas and form sodium alkoxides (RONa). Reactivity: 1° > 2° > 3° (steric hindrance). This confirms the weakly acidic nature of the O—H bond (pKₐ ~16–18).
Primary alcohols oxidise stepwise to aldehydes then carboxylic acids. Distilling the aldehyde removes it before further oxidation. Secondary alcohols stop at ketones. Tertiary alcohols cannot be oxidised without C—C bond cleavage.
At 170 °C: intramolecular dehydration gives alkene. At 140 °C: two alcohol molecules lose water to give an ether. The more substituted (Zaitsev) alkene is the major product from secondary/tertiary alcohols.
The oxygen in the ester comes from the alcohol (proven by ¹⁸O isotope labelling). Yield improved by excess alcohol, removing water, or using a drying agent.
| Alcohol Type | Observation | Time | Reason |
|---|---|---|---|
| Tertiary (3°) | Immediate cloudiness / turbidity | < 1 min | Stable 3° carbocation — fast SN1 |
| Secondary (2°) | Cloudiness within 5 min | ~5 min | Less stable 2° carbocation — slower SN1 |
| Primary (1°) | No turbidity at room temperature | No reaction | 1° doesn't form carbocation easily — only reacts on heating |
The —OH group is bonded directly to the aromatic ring. The oxygen lone pairs delocalise into the π system, weakening the O—H bond and making phenol acidic (pKₐ ≈ 10) — much stronger acid than alcohols (pKₐ ≈ 16).
The ring is activated — electron density increases at the ortho and para positions, directing EAS to those positions.
Phenol reacts with NaOH to form sodium phenoxide (soluble). Alcohols do NOT react with NaOH — this is the key chemical distinction. Phenol is acidic enough to react with carbonates too.
Phenol reacts instantly with bromine water to give 2,4,6-tribromophenol as a white precipitate. All three activated positions are substituted simultaneously. Used as a qualitative test for phenol.
Phenol reacts with dilute nitric acid (no catalyst needed) to give a mixture of ortho- and para-nitrophenol. Concentrated HNO₃/H₂SO₄ gives 2,4,6-trinitrophenol (picric acid). Much milder than for benzene.
Phenols form a characteristic violet or purple complex with iron(III) chloride solution. Alcohols give no colour with FeCl₃. This is a key identification test to distinguish phenols from alcohols.
Like alcohols, phenol reacts with sodium to release H₂ and form sodium phenoxide. The reaction is more vigorous because phenol is a stronger acid. The phenoxide ion is stabilised by resonance with the ring.
Sodium phenoxide reacts with CO₂ under pressure to give sodium salicylate, which on acidification gives salicylic acid — the precursor to aspirin. Classic industrial process in pharma.
| Property | Alcohols (R—OH) | Phenols (Ar—OH) |
|---|---|---|
| —OH attached to | Saturated sp³ carbon | Aromatic ring (sp²) |
| Acidity (pKₐ) | ~16–18 (weak) | ~10 (much stronger) |
| Reaction with NaOH | No reaction | Reacts → sodium phenoxide |
| Reaction with Na₂CO₃ | No reaction | Reacts → CO₂ not evolved |
| FeCl₃ test | No colour | Violet / purple colour |
| Bromine water | No reaction (saturated) | Immediate white ppt (tribromophenol) |
| Oxidation | Oxidised by K₂Cr₂O₇ | Ring resists simple oxidation |
| Esterification | Easy (+ RCOOH / H⁺) | Requires acyl chloride / anhydride |
| Lone pair delocalisation | No (sp³ O) | Yes — into aromatic π system |
Organic synthesis is the art of building target molecules from simpler starting materials by choosing the right sequence of reactions. Understanding retrosynthetic analysis — working backwards from the target — is the key skill.
| Factor | SN1 | SN2 | E1 | E2 |
|---|---|---|---|---|
| Substrate | 3° > 2° | 1° > 2° | 3° > 2° | 3° > 2° > 1° |
| Nucleophile/Base | Weak Nu | Strong Nu | Weak base | Strong base |
| Solvent | Polar protic | Polar aprotic | Polar protic | Polar aprotic / alc. |
| Mechanism | 2 steps (carbocation) | 1 step (concerted) | 2 steps | 1 step (concerted) |
| Stereochemistry | Racemisation | Inversion | Racemisation | Anti periplanar |
| Kinetics | 1st order | 2nd order | 1st order | 2nd order |
| Rearrangement | Yes | No | Yes | No |
When a molecule has two reactive functional groups and you want to react only one, the other must be temporarily protected. Then you react the exposed group, and finally deprotect to reveal the original group.
Core chemistry terms and their meanings — the essential vocabulary for understanding all branches of chemistry.
The smallest unit of an element that retains its chemical properties. Composed of protons, neutrons, and electrons.
Two or more atoms chemically bonded together. Ex: H₂O, CO₂, O₂
A pure substance made of only one kind of atom. Cannot be broken down by chemical means. Ex: Hydrogen (H), Oxygen (O)
A substance formed when two or more different elements combine chemically in a fixed ratio. Ex: NaCl, H₂O, CO₂
A physical combination of two or more substances that are not chemically bonded. Can be separated physically. Ex: Air (N₂, O₂, CO₂…), sand + water
An atom or molecule that carries an electric charge due to gaining or losing electrons. Ex: Na⁺, Cl⁻, SO₄²⁻
A positively charged ion formed by the loss of electrons. Ex: Na⁺ (sodium loses 1e⁻), Ca²⁺
A negatively charged ion formed by the gain of electrons. Ex: Cl⁻ (chlorine gains 1e⁻), O²⁻
A substance that donates protons (H⁺) or produces hydrogen ions in aqueous solution. pH < 7. Ex: HCl, H₂SO₄, CH₃COOH
A substance that accepts protons or produces hydroxide ions (OH⁻) in aqueous solution. pH > 7. Ex: NaOH, KOH, NH₃
An ionic compound formed by the reaction of an acid and a base (neutralisation). Ex: NaCl, CaSO₄, K₂CO₃
A homogeneous mixture where a solute is uniformly dissolved in a solvent. Ex: Salt water, sugar water
The substance that is dissolved in a solution. Present in lesser amount. Ex: Salt in salt water
The substance that dissolves the solute to form a solution. Present in greater amount. Ex: Water (universal solvent)
A process in which substances (reactants) undergo chemical change to form new substances (products). Bonds break and form. Ex: A + B → C
A scale (0–14) that measures the acidity or basicity of a solution. pH = −log[H⁺]. pH 7 = neutral, <7 = acidic, >7 = basic.
A substance that increases the rate of a chemical reaction without being consumed. Lowers activation energy. Ex: MnO₂, enzymes, Pt
A substance that produces ions when dissolved in water, allowing the solution to conduct electricity. Ex: NaCl, HCl, KOH
The reactivity series ranks metals (and hydrogen) from most reactive to least reactive. A more reactive metal can displace a less reactive one from its compound in solution.
A more reactive metal displaces a less reactive one from its salt solution.
K, Na, Li, and Ca react with cold water. Mg reacts with hot water/steam. Metals below Mg react only with steam or dilute acids.
Metals above hydrogen react with dilute acids to produce a salt + hydrogen gas.
Metals below H (Cu, Hg, Ag, Au, Pt) do NOT react with dilute acids.
Common cations (positive ions) and anions (negative ions) organized by charge. The building blocks of ionic compounds.
| Symbol | H | Li | N | Na | S | K | C | Cl | O |
|---|---|---|---|---|---|---|---|---|---|
| Mass | 1 | 7 | 14 | 23 | 32 | 39 | 12 | 35.5 | 16 |
| Symbol | Al | Ca | P | Si | Ag | Fe | Ba | Mg | Cu |
|---|---|---|---|---|---|---|---|---|---|
| Mass | 27 | 40 | 31 | 28 | 107 | 56 | 137 | 24 | 63.5 |
Multi-atom ions that behave as a single charged unit in ionic compounds.
The mole is chemistry's counting unit — a "mole" means 6.02 × 10²³ particles (Avogadro's number). Cover the unknown quantity in each triangle to find the formula to use.
1 mole of rice grains would cover Earth to ~75 metres depth. Yet 1 mole of carbon weighs just 12 grams — that's why the mole bridges atomic-scale counting and lab-scale measurement.
The coefficients in a balanced equation give the mole ratio of each substance. Use them as a bridge to convert between moles of any reactant or product.
N₂ + 3H₂ → 2NH₃n = mass ÷ molar massHow many moles of NH₃ form from 6 mol H₂?
Ratio: 3 mol H₂ : 2 mol NH₃ → 6 mol H₂ × (2/3) = 4 mol NH₃
The limiting reagent is the reactant that runs out first and determines the maximum yield. The other reactant is in excess.
2H₂ + O₂ → 2H₂O. You have 5 mol H₂ and 2 mol O₂.
H₂: 5÷2 = 2.5 | O₂: 2÷1 = 2.0
O₂ is limiting → 2 mol O₂ × 2 = 4 mol H₂O formed.
% Yield = (Actual ÷ Theoretical) × 100
Always ≤ 100%. Losses occur from incomplete reactions, side reactions, or transfer losses.
Simplest whole-number ratio of atoms. e.g. CH₂O (glucose simplified). Found from % composition data.
Actual number of atoms. e.g. C₆H₁₂O₆ (glucose). = Empirical × n, where n = Molar Mass ÷ Empirical Mass.
n = Mᵣ ÷ Empirical MassC₁V₁ = C₂V₂
Moles of solute stay constant when diluting. Multiply initial concentration × volume = final concentration × volume.
At equivalence point: moles acid = moles base (for 1:1 reactions). Use C = n ÷ V to find unknown concentration from titre volume.
1 calorie = energy to raise 1g water by 1°C = 4.18 joules. Always check which unit is expected before solving.
The International System of Units (SI) defines 7 base units from which all other measurement units are derived.
| Quantity | SI Unit | Symbol |
|---|---|---|
| Length | metre | m |
| Mass | kilogram | kg |
| Time | second | s |
| Electric Current | ampere | A |
| Temperature | kelvin | K |
| Amount of Substance | mole | mol |
| Luminous Intensity | candela | cd |
| Quantity | Unit | Symbol |
|---|---|---|
| Force | newton | N |
| Energy | joule | J |
| Power | watt | W |
| Pressure | pascal | Pa |
| Frequency | hertz | Hz |
| Electric Charge | coulomb | C |
| Voltage | volt | V |
| Speed / Velocity | metre/second | m/s |
K = °C + 273.15 | °C = K − 273.15 | °F = (°C × 9/5) + 32
Always use Kelvin in gas law and thermodynamic calculations — never Celsius.
Thermochemistry studies heat changes in chemical reactions. Every reaction either releases or absorbs energy — understanding this predicts spontaneity and drives industrial design.
Release heat to surroundings. ΔH is negative. Products at lower energy than reactants. Examples: combustion, neutralisation, rusting of iron.
Absorb heat from surroundings. ΔH is positive. Products at higher energy. Examples: photosynthesis, dissolving NH₄NO₃, thermal decomposition.
ΔH = Hproducts − Hreactants. Measured at 25°C and 1 atm (standard conditions). We can only measure changes, never absolute enthalpy.
e.g. Boiling pot without a lid
e.g. A sealed container
e.g. A perfect thermos
Nuclear chemistry deals with atomic nuclei transformations — releasing enormous energy via E = mc². Underpins nuclear energy, medical imaging (PET scans), cancer radiation therapy, and radiocarbon dating.
(no. protons × proton mass) + (no. neutrons × neutron mass)Δm = Calculated Mass − Actual Atomic WeightMass Defect × 931 (MeV)N.B.E ÷ Mass Number| Property | α Alpha | β Beta | γ Gamma |
|---|---|---|---|
| Symbol | α | β | γ |
| Nature | ₂He⁴ nucleus | ₋₁e⁰ electron | Electromagnetic |
| Mass | 4× proton | 1/1800 proton | Massless |
| Charge | +VE | −VE | None |
| Penetration | Weak | Medium | Very High |
| Ionization | Very High | High | Low |
| E/B field | Low deflection | Large deflection | Unaffected |
Chemical bonds are the forces that hold atoms together to form molecules and compounds. The type of bond that forms depends on the electronegativity difference between the atoms involved. Understanding bonds explains why water is liquid at room temperature, why metals conduct electricity, and why NaCl dissolves in water but not in oil.
Electron transfer metal → non-metal. ΔEN > 1.7.
Shared electrons between non-metals. ΔEN < 1.7.
Electrons spend more time near the more electronegative atom, creating δ+ and δ− partial charges — giving H₂O its unusual surface tension and boiling point.
Dots = valence electrons. Pairs between atoms = bonds. Octet rule: atoms seek 8 valence electrons (H seeks 2). Lone pairs don't bond directly.
Recognising the type of reaction lets you predict products and balance equations — the core skill of general chemistry.
Two or more substances combine into one product.
One compound breaks into two or more simpler substances.
One element replaces another in a compound. Activity series determines feasibility.
Two compounds exchange partners — yields precipitate or water.
Substance + O₂ → energy. Hydrocarbons → CO₂ + H₂O.
Acid + Base → Salt + Water. H⁺ + OH⁻ → H₂O.
C₂H₆ + O₂ → H₂O + CO₂C₂H₆ + 7/2 O₂ → 3H₂O + 2CO₂2C₂H₆ + 7O₂ → 6H₂O + 4CO₂Follow this order when balancing — it avoids backtracking. Always leave H and O until last.
Metals: 1 Al left → 2 needed (×2). Non-metals: 1 S left → 3 needed (×3 H₂SO₄). Hydrogen: 6H left → 6H right (×3 H₂). Oxygen: 12O each side ✓ → Balanced!
Acids donate protons (H⁺); bases accept protons or donate OH⁻. The pH scale measures acidity logarithmically from 0–14.
Donate H⁺. pH < 7. Taste sour.
Accept H⁺ or donate OH⁻. pH > 7. Feel slippery.
H⁺ + OH⁻ → H₂O. Resulting salt may be acidic, basic, or neutral depending on the relative strengths of original acid and base.
Resist pH change when small amounts of acid/base are added. Weak acid + conjugate base. Blood uses HCO₃⁻/H₂CO₃ to maintain pH ≈ 7.4.
Conjugate acid-base pairs — stronger acids have weaker conjugate bases. Acid strength increases downward ↓ on the left; base strength increases upward ↑ on the right.
Electrochemistry studies the relationship between chemical reactions and electrical energy — including oxidation-reduction and electrochemical cells.
Loses electrons. Oxidation number increases. In galvanic cells: Anode is −.
Zn → Zn²⁺ + 2e⁻Gains electrons. Oxidation number decreases. In galvanic cells: Cathode is +.
Cu²⁺ + 2e⁻ → CuChemical energy → electricity spontaneously (ΔG < 0). Salt bridge maintains neutrality. e.g. Zn-Cu Daniell cell.
Uses electricity to drive non-spontaneous reactions (ΔG > 0). Used in electroplating, refining copper, producing aluminium.
| Property | Oxidation | Reduction |
|---|---|---|
| Electrons | Loss | Gain |
| Oxygen | Gain | Loss |
| Hydrogen | Loss | Gain |
| Oxidation Number | Increases ↑ | Decreases ↓ |
| Electrode (electrolysis) | Anode (+) | Cathode (−) |
| Agent | Reducing agent (gives e⁻) | Oxidising agent (takes e⁻) |
Fe → Fe²⁺ + 2e⁻ (oxidation). O₂ + 4H⁺ + 4e⁻ → 2H₂O (reduction). Overall: 4Fe + 3O₂ → 2Fe₂O₃. Requires both O₂ and H₂O.
Carbon in fuel is oxidised (gains O₂). CH₄ + 2O₂ → CO₂ + H₂O. Carbon oxidation state goes from −4 to +4 — maximum oxidation.
C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + ATP. Glucose (C) is oxidised; oxygen is reduced. Biological redox chain in mitochondria.
Spontaneous redox reactions generate electrical current. Zn anode oxidises (Zn → Zn²⁺ + 2e⁻); MnO₂ cathode is reduced. e⁻ flow through circuit = current.
The electrochemical series ranks metals (and non-metals) by their standard electrode potential (E°) — measured against the standard hydrogen electrode (SHE = 0.00 V). More negative E° = stronger reducing agent. More positive E° = stronger oxidising agent.
| Half-Reaction (Reduction) | E° (V) | Reducing Strength |
|---|---|---|
| Li⁺ + e⁻ → Li | −3.04 | ⬆ Strongest Reducing Agent |
| K⁺ + e⁻ → K | −2.93 | Very strong reducer |
| Ca²⁺ + 2e⁻ → Ca | −2.87 | |
| Na⁺ + e⁻ → Na | −2.71 | |
| Mg²⁺ + 2e⁻ → Mg | −2.37 | |
| Al³⁺ + 3e⁻ → Al | −1.66 | |
| Zn²⁺ + 2e⁻ → Zn | −0.76 | |
| Fe²⁺ + 2e⁻ → Fe | −0.44 | |
| Ni²⁺ + 2e⁻ → Ni | −0.25 | |
| Pb²⁺ + 2e⁻ → Pb | −0.13 | |
| 2H⁺ + 2e⁻ → H₂ | 0.00 | ← Standard (SHE) |
| Cu²⁺ + 2e⁻ → Cu | +0.34 | |
| I₂ + 2e⁻ → 2I⁻ | +0.54 | |
| Fe³⁺ + e⁻ → Fe²⁺ | +0.77 | |
| Ag⁺ + e⁻ → Ag | +0.80 | |
| Br₂ + 2e⁻ → 2Br⁻ | +1.07 | |
| O₂ + 4H⁺ + 4e⁻ → 2H₂O | +1.23 | |
| Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O | +1.33 | |
| Cl₂ + 2e⁻ → 2Cl⁻ | +1.36 | |
| MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O | +1.51 | |
| F₂ + 2e⁻ → 2F⁻ | +2.87 | ⬇ Strongest Oxidising Agent |
If E°cell > 0, the reaction is spontaneous.
If E°cell < 0, the reaction is non-spontaneous.
A metal higher in the series (more negative E°) will displace a metal lower in the series from solution.
The activity series for metals in order of reactivity with water and acids:
The quantitative side of electrochemistry — connecting cell potential to Gibbs free energy, equilibrium constants, and reaction spontaneity through the Nernst equation and Faraday's laws.
All three quantities are linked: ΔG° = −nFE° = −RT ln K. If you know any one of them, you can find the other two. A spontaneous reaction (ΔG° < 0) has E° > 0 and K > 1 — products are favoured at equilibrium.
E = E°. All concentrations are 1 M, all pressures 1 atm — standard conditions. The equation reduces to just E°.
−ln Q > 0, so E > E°. More reactants than products pushes the cell forward — higher driving force.
E = 0. The cell is at equilibrium and can no longer do work. This is a dead battery. E = 0 when Q = K.
Mathematical relationships between pressure (P), volume (V), temperature (T) and moles (n) for ideal gases.
At constant T: P and V are inversely proportional.
At constant P: V and T are directly proportional. T in Kelvin.
At constant V: P and T are directly proportional. T in Kelvin.
Merges Boyle's, Charles's and Gay-Lussac's. n constant.
R = 8.314 J/(mol·K). n = moles. Assumes ideal behaviour.
Equal volumes at same T, P → equal moles. 1 mol = 22.4 L at STP.
A reversible reaction reaches equilibrium when the rate of forward reaction equals the rate of reverse reaction — concentrations become constant but both reactions continue. This is dynamic, not static equilibrium.
When a system at equilibrium is subjected to a stress (change in concentration, pressure, or temperature), the equilibrium shifts in the direction that opposes the stress and partially restores equilibrium.
The ICE (Initial–Change–Equilibrium) table is the standard method for calculating equilibrium concentrations when you know initial conditions and Kc.
| N₂ | 3H₂ | 2NH₃ | |
|---|---|---|---|
| I — Initial | 1.000 M | 3.000 M | 0.000 M |
| C — Change | −x | −3x | +2x |
| E — Equilibrium | 1−x | 3−3x | 2x |
Q has the same form as Kc but uses current (not equilibrium) concentrations. If Q < Kc: reaction shifts forward. If Q > Kc: reaction shifts backward. If Q = Kc: at equilibrium.
Kp uses partial pressures instead of concentrations (for gas phase reactions). Kp = Kc(RT)^Δn where Δn = moles of gaseous products − moles of gaseous reactants.
Solutions are homogeneous mixtures. Concentration measures how much solute is present relative to the solution or solvent.
Kinetics studies how fast reactions occur and what factors control that speed. Understanding rate is essential for industrial chemistry, drug design, and predicting reaction feasibility.
For a reaction to occur, particles must collide with sufficient energy (≥ activation energy) and with the correct orientation. Increasing either factor increases the rate.
Minimum energy needed for a collision to result in reaction. Lower Eₐ = faster reaction. Catalysts work by providing an alternative pathway with lower Eₐ.
Must have: (1) energy ≥ Eₐ and (2) correct geometric orientation. Even high-energy collisions fail if molecules approach at the wrong angle.
Not all particles have the same energy. The distribution curve shows most particles cluster around average energy — only the tail exceeds Eₐ and reacts.
Higher T → particles move faster → more collisions, more exceed Eₐ. A 10°C rise roughly doubles the rate for many reactions.
More particles per unit volume → more frequent collisions → higher rate. Doubles concentration can more than double rate depending on order.
Smaller particle size → greater surface area exposed → more collision sites. Powder reacts faster than a lump of the same mass.
Provides an alternative reaction pathway with lower Eₐ. Not consumed in the reaction. Homogeneous (same phase) or heterogeneous (different phase).
Increased pressure = particles closer together = more collisions per second. Equivalent to increasing concentration for gas-phase reactions.
UV/visible light can provide activation energy for photochemical reactions. e.g. halogen + alkane reactions initiated by UV light.
Sum of all individual orders: m + n. Zero order: rate independent of [A]. First order: rate ∝ [A]. Second order: rate ∝ [A]².
k = Ae^(−Eₐ/RT)
Links rate constant to temperature. A = frequency factor. R = 8.314 J/(mol·K). As T↑, k↑, rate↑.
Time for [reactant] to halve. First order: t½ = ln2/k (constant). Second order: t½ = 1/(k[A]₀) — depends on initial concentration.
Core practical skills used to separate, purify, and identify substances. These techniques underpin every branch of chemistry and are essential exam content.
Separates an insoluble solid from a liquid. Mixture poured through filter paper in a funnel. The solid (residue) stays on the filter; the liquid (filtrate) passes through.
Dissolve solid in hot solvent. Evaporate slowly to form pure crystals. Filter, wash with cold solvent, dry. Used to purify salts.
Separates liquids by boiling point. Heat mixture → vapour rises → condenses in condenser → collected as distillate. Fractional distillation uses a fractionating column for close boiling points.
Separates mixtures based on how far components travel through a stationary phase carried by a mobile phase. Rf = distance moved by spot ÷ distance moved by solvent.
Separates compounds using two immiscible solvents (e.g. water + ethyl acetate). The compound partitions into the solvent it's more soluble in. Used in a separating funnel.
Heats a reaction mixture at its boiling point without losing solvent. Vapour condenses in a vertical condenser and falls back. Used for slow reactions needing prolonged heating.
| Type | Mobile Phase | Stationary Phase | Use |
|---|---|---|---|
| TLC | Solvent (liquid) | Silica on plate | Purity check, identify compounds |
| Paper | Solvent (liquid) | Water on paper | Separate coloured mixtures, amino acids |
| GC | Inert gas (N₂, He) | Liquid on solid | Separate volatile compounds, food/forensics |
| HPLC | Solvent (high pressure) | Solid column | Pharmaceuticals, large biomolecules |
Rf = distance travelled by spot ÷ distance travelled by solvent front
Always between 0 and 1. Same compound always gives the same Rf under the same conditions. Used to identify unknowns by comparison.
Spectroscopy uses the interaction of electromagnetic radiation with matter to identify substances and determine structure. Three techniques dominate chemistry: IR, Mass Spec, and NMR.
Molecules are ionised and fragmented. The detector records mass-to-charge ratio (m/z) of each fragment. Used to find molecular mass and deduce structure.
The rightmost peak = molecular mass of the compound. The M+1 peak indicates presence of ¹³C isotopes. M+2 indicates Cl or Br (characteristic isotope pattern).
Molecule breaks at weakest bonds. Each fragment gives a peak at its m/z. Common fragments: m/z 15 (CH₃⁺), 29 (C₂H₅⁺), 45 (OC₂H₅⁺), 77 (C₆H₅⁺ phenyl).
The tallest peak — most stable/abundant fragment. Used as reference (set to 100%). Does not have to be the molecular ion peak.
Bonds absorb IR at characteristic frequencies (wavenumbers, cm⁻¹). Identifying absorption peaks reveals which functional groups are present.
| Bond | Wavenumber (cm⁻¹) | Functional Group | Notes |
|---|---|---|---|
| O–H (alcohol) | 3200–3550 | Alcohol | Broad, strong |
| O–H (acid) | 2500–3300 | Carboxylic acid | Very broad |
| N–H | 3300–3500 | Amine / Amide | Medium, 1 or 2 peaks |
| C–H | 2850–3100 | Alkane/Alkene | Medium |
| C≡N | 2200–2260 | Nitrile | Strong, sharp |
| C≡C | 2100–2260 | Alkyne | Weak/absent if symmetric |
| C=O | 1630–1820 | Aldehyde/Ketone/Ester/Acid | Strong, sharp |
| C=C | 1620–1680 | Alkene | Medium |
| C–O | 1000–1300 | Alcohol / Ester / Ether | Strong |
¹H NMR detects hydrogen environments in a molecule. Each unique H environment gives a signal. Chemical shift (δ, ppm) indicates the electronic environment of each H.
TMS (tetramethylsilane) = 0 ppm reference. Electron-withdrawing groups shift signals downfield (higher δ). e.g. –CHO ≈ 9–10 ppm, –COOH ≈ 10–12 ppm, –CH₃ ≈ 0.9 ppm.
A signal is split into n+1 peaks by n adjacent H atoms. Doublet = 1 neighbour. Triplet = 2 neighbours. Quartet = 3 neighbours. Singlet = no neighbours.
Area under each peak ∝ number of H atoms in that environment. Ratios identify how many H each signal represents. Combined with splitting → full structure.
① Find molecular formula (from mass spec M⁺) → calculate degrees of unsaturation. ② Count distinct signals = number of unique H environments. ③ Use integration for ratios. ④ Use splitting pattern (n+1) to find neighbours. ⑤ Match δ values to functional group table to confirm structure.
The magnetic behaviour of a substance is determined by how electrons are arranged in its atoms or ions. Unpaired electrons are the key: each unpaired electron acts as a tiny magnet, and their collective effect determines whether a material is attracted to, repelled by, or unaffected by a magnetic field.
All electrons are paired. Paired electrons have opposite spins that cancel out each other's magnetic moment. Result: the substance is weakly repelled by a magnetic field. Most non-metals and main-group compounds with complete subshells are diamagnetic.
Examples: He, N₂, NaCl, H₂O, Cu⁺ (d¹⁰), Zn²⁺ (d¹⁰)
Note: Cu²⁺ is d⁹ (1 unpaired e⁻) → paramagnetic, not diamagnetic.
One or more unpaired electrons are present. Each unpaired electron produces a net magnetic moment. The substance is weakly attracted to a magnetic field but loses magnetism when the field is removed (moments are randomly oriented without a field).
Examples: O₂, Fe³⁺ (5 unpaired d-electrons), Ti³⁺ (1 unpaired), Mn²⁺ (5 unpaired)
A special form of paramagnetism found in Fe, Co, and Ni. In these metals, magnetic domains (regions where all atomic moments align) form spontaneously. Domains can align permanently with an external field, producing a permanent magnet. Disappears above the Curie temperature.
Curie temps: Fe 768°C · Co 1115°C · Ni 358°C
Write the electronic configuration of the ion. Fill d-orbitals using Hund's rule (one electron per orbital before pairing). Count unpaired electrons.
The magnetic moment (μ) of an ion can be estimated using the spin-only formula:
μ = √[n(n+2)] Bohr Magnetons
where n = number of unpaired electrons. Fe³⁺ has n=5 → μ = √(35) ≈ 5.92 BM. This can be measured experimentally to determine the number of unpaired electrons and oxidation state.
Transition metals and their compounds are outstanding catalysts — substances that increase the rate of a reaction without being consumed. Their variable oxidation states and ability to adsorb reactants on their surface make them industrially indispensable.
Two key properties drive catalytic activity:
Catalyst and reactants are in the same phase (usually both dissolved in solution).
Catalyst is in a different phase from reactants (usually a solid catalyst with gas or liquid reactants).
| Process | Catalyst | Conditions | Purpose |
|---|---|---|---|
| Haber Process | Fe (iron) | 450°C, 200 atm | Ammonia synthesis for fertilisers |
| Contact Process | V₂O₅ | 450°C, 1 atm | H₂SO₄ production (SO₂→SO₃) |
| Catalytic Converter | Pt, Pd, Rh | ~600°C | Remove CO, NOₓ, hydrocarbons |
| Hydrogenation | Ni | ~180°C, pressure | Oil → margarine (saturating C=C) |
| Ziegler-Natta | TiCl₄ / Al(C₂H₅)₃ | Low T, low P | Stereospecific polymerisation (HDPE) |
Iron is extracted from its ores (primarily haematite, Fe₂O₃) in a blast furnace — a continuous industrial process that operates at temperatures up to 2000°C and produces molten iron (pig iron) that is then refined into steel.
C + O₂ → CO₂ (combustion of coke, very exothermic)
CO₂ + C → 2CO (CO is the main reducing agent)
Fe₂O₃ + 3CO → 2Fe + 3CO₂ (main reduction)
Fe₃O₄ + 4CO → 3Fe + 4CO₂ (magnetite reduction)
CaCO₃ → CaO + CO₂ (limestone decomposes)
CaO + SiO₂ → CaSiO₃ (slag formation)
Pig iron is refined in a Basic Oxygen Furnace (BOF). High-purity O₂ is blown into molten pig iron to oxidise excess carbon and impurities:
C + O₂ → CO₂ (removes carbon)
Carbon content is controlled to produce different steel grades: mild steel (0.1–0.3% C), medium carbon steel (0.3–0.6% C), high carbon steel (0.6–1.4% C). Alloying elements (Cr, Ni, Mn) are added to make specialty steels.
Pure iron (symbol Fe, atomic number 26) is a lustrous silver-grey metal with remarkable magnetic, mechanical, and chemical properties. It rarely occurs in nature in its pure form — it is almost always found as oxides or carbonate ores.
Rusting is an electrochemical process requiring both oxygen AND water (not just one alone). Iron acts as an anode, atmospheric oxygen is reduced at the cathode, with the moisture acting as the electrolyte:
Anode: Fe → Fe²⁺ + 2e⁻
Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻
Overall: Fe²⁺ + 2OH⁻ → Fe(OH)₂ → oxidised further to Fe₂O₃·H₂O (rust)
Prevention methods: Galvanising (zinc coating), electroplating, painting, oiling, alloying (stainless steel), cathodic protection (sacrificial anode of Mg or Zn).
Iron forms three principal oxides: FeO, Fe₂O₃, and Fe₃O₄. Each has a distinct colour, oxidation state composition, crystal structure, and set of industrial uses. Understanding them is central to iron metallurgy, pigments, and materials science.
| Property | FeO | Fe₂O₃ | Fe₃O₄ |
|---|---|---|---|
| Fe oxidation state(s) | +2 | +3 | +2, +3 |
| Colour | Black | Red/brown | Black |
| Magnetism | Paramagnetic | Weakly paramagnetic | Ferrimagnetic |
| Stability in air | Unstable (oxidises) | Stable | Stable |
An alloy is a metallic material made by mixing two or more elements, where at least one is a metal. Alloys almost always have superior properties to pure metals — greater hardness, corrosion resistance, strength, or specific melting points — making them the backbone of modern engineering.
Pure metals have layers of atoms that can slide past each other easily — this is why they are soft and malleable. Adding differently-sized atoms of another element distorts the regular lattice, making it harder for layers to slide. Result: increased hardness and strength.
The added atoms can be larger (substitutional alloy) or smaller (interstitial alloy) than the host metal atoms.
The most important industrial alloy. Carbon is an interstitial impurity that dramatically changes properties:
Typically: 74% Fe, 18% Cr, 8% Ni. The chromium reacts with O₂ to form a thin, adherent Cr₂O₃ layer that prevents further oxidation — the steel is effectively self-healing and corrosion-resistant.
Duralumin: Al + 4% Cu + Mn + Mg. Age-hardened to be almost as strong as mild steel but only one-third the density. Used in aircraft fuselages and spacecraft.
Ti alloys (Ti-6Al-4V): Titanium + aluminium + vanadium. Strongest-to-weight ratio of any structural metal. Used in jet engines, medical implants, high-performance sporting goods.
70% Cu + 30% Zn. Golden colour. Harder than pure copper, good corrosion resistance, easily machined. Used in musical instruments, plumbing fittings, zippers, electrical terminals, ammunition casings.
~88% Cu + 12% Sn. Historically the first alloy ever made (Bronze Age). Much harder than copper, resistant to seawater corrosion. Used in ship propellers, bearings, bells, sculpture, coins.
Traditional solder: 63% Sn + 37% Pb (mp 183°C — true eutectic). Used in electronics. Lead-free versions (SAC alloys: Sn + Ag + Cu) are now mandated by RoHS regulations. Gold alloys (karat system): 24K = pure gold, 18K = 75% Au + 25% other metals.
Basic radicals (cations) in an unknown salt can be systematically identified through a series of chemical tests. The classical scheme uses flame tests, precipitation with NaOH and NH₃, and confirmatory reactions to narrow down and confirm the identity of the cation present.
Metallic ions impart characteristic colours when introduced into a flame. The heat excites electrons to higher energy levels; when they fall back, they emit visible light.
Add NaOH(aq) to the salt solution in drops then in excess:
Add dilute NH₃ in drops then in excess:
Quantitative analysis determines the amount (mass, concentration, or percentage) of a substance in a sample using precise measurements. The core techniques include gravimetric analysis, volumetric (titrimetric) analysis, and colorimetry.
The analyte is converted into a pure, stable, sparingly soluble precipitate that is filtered, dried, and weighed. The mass of the precipitate is used to calculate the amount of analyte.
Steps: Dissolve sample → precipitate analyte → filter → wash → dry/ignite to constant mass → weigh → calculate.
Example: BaSO₄ gravimetry — add excess BaCl₂ to a sulfate solution: Ba²⁺ + SO₄²⁻ → BaSO₄↓ (white). Weigh BaSO₄ → calculate sulfate content.
Accuracy tip: Digestion (heating near boiling) is used to grow larger crystals → easier to filter, less surface area to adsorb impurities.
A solution of known concentration (standard solution / titrant) is delivered from a burette until the reaction is complete (equivalence point). An indicator signals the end-point.
Key formula: C₁V₁/n₁ = C₂V₂/n₂
Types of titration:
Based on Beer-Lambert Law: the absorbance of a solution is proportional to concentration.
A = εcl
Calibration curve: Prepare standards of known concentration → measure absorbance → plot A vs. c → use to read off unknown concentration.
Used when the analyte reacts slowly, is insoluble, or the end-point is hard to detect directly. An excess of known reagent is added to react completely with the analyte, then the unreacted excess is titrated with a second standard solution.
moles(analyte) = moles(reagent added) − moles(reagent titrated back)
Example: Determining CaCO₃ purity — add excess HCl, then back-titrate with NaOH.
Pressure is the force exerted per unit area. In chemistry, pressure plays a central role in gas behaviour, equilibrium, and industrial processes. Understanding different types of pressure — atmospheric, partial, osmotic, and vapour pressure — is essential for physical chemistry.
The pressure exerted by the weight of the atmosphere. Standard atmosphere (STP): 1 atm = 101,325 Pa = 760 mmHg = 760 torr = 1.01325 bar.
Measured by barometer (mercury column). At sea level: 1 atm. Decreases with altitude — at 10 km (cruising altitude) it is ~0.25 atm.
The total pressure of a mixture of non-reacting gases equals the sum of the partial pressures of each individual gas:
P_total = P₁ + P₂ + P₃ + ...
Partial pressure of gas i: Pᵢ = xᵢ × P_total, where xᵢ is the mole fraction. Used in calculations involving gas collection over water: P(dry gas) = P(total) − P(water vapour).
The pressure exerted by a vapour in equilibrium with its liquid phase at a given temperature. Increases with temperature (more molecules have enough energy to escape). At the boiling point, vapour pressure = atmospheric pressure.
Raoult's Law (ideal solutions): P = x_solvent × P°_solvent. Adding a non-volatile solute lowers vapour pressure → elevates boiling point and depresses freezing point.
The pressure required to prevent osmosis (net flow of solvent through a semipermeable membrane from low to high concentration). Van't Hoff equation:
π = iMRT
From Le Chatelier's principle, increasing pressure on a gas-phase equilibrium shifts it toward the side with fewer moles of gas.
Ionic equilibrium describes reversible processes involving ions in solution — including the dissociation of weak acids and bases, solubility equilibria, buffer systems, and hydrolysis. It extends the concept of chemical equilibrium to ionic species.
A weak acid HA only partially dissociates in water:
HA ⇌ H⁺ + A⁻
Kₐ = [H⁺][A⁻] / [HA]
A weak base B accepts a proton from water:
B + H₂O ⇌ BH⁺ + OH⁻
K_b = [BH⁺][OH⁻] / [B]
A buffer resists changes in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base (or weak base + conjugate acid) in comparable concentrations.
Henderson-Hasselbalch:
pH = pKₐ + log([A⁻]/[HA])
For a sparingly soluble salt MA dissolving:
MA(s) ⇌ M⁺(aq) + A⁻(aq)
Ksp = [M⁺][A⁻]
When a salt dissolves in water, the resulting solution may be acidic, basic, or neutral depending on the nature of the parent acid and base:
For a salt of weak acid and strong base (e.g. CH₃COONa), the degree of hydrolysis (h) is:
h = √(Kw / Kₐ × C)
pH = 7 + ½(pKₐ + log C)
The Common Ion Effect also suppresses hydrolysis — adding the weak acid (CH₃COOH) to CH₃COONa solution reduces the degree of hydrolysis of the acetate ion.
Six live calculators built directly into the encyclopedia — no external apps needed.
Criss-cross method — select ions and get the formula automatically.
Enter a balanced equation and a known quantity to find unknown amounts. Format: coefficients then formulas separated by spaces (reactants + products).
Enter any one value — get all the others instantly.
The periodic table is divided into four blocks based on the subshell being filled with electrons (s, p, d, f). This directly explains each group's chemical properties and reactivity patterns.
Alkali metals (1 valence e⁻) and alkaline earth metals (2 valence e⁻). Highly reactive, low ionisation energy. React vigorously with water.
The most diverse block — contains metals, metalloids, nonmetals, halogens, and noble gases. Properties vary widely across the block.
Transition metals — filling d-subshell. Variable oxidation states, coloured compounds, catalytic activity, form complex ions. Essential in industry and biochemistry.
Inner transition metals. Lanthanides are used in high-tech magnets, phosphors, lasers. Actinides (Z=89–103) are mostly radioactive — uranium and plutonium for nuclear energy.
Electrons fill orbitals in order of increasing energy: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p...
Within a subshell, electrons fill each orbital singly before pairing (maximise spin multiplicity). e.g. N: [↑][↑][↑] across 2p — not [↑↓][↑][ ].
No two electrons in an atom can have the same four quantum numbers. Each orbital holds max 2 electrons with opposite spins (↑↓).
d-d electron transitions absorb visible light. Different oxidation states = different colours. e.g. Cu²⁺ = blue, Fe³⁺ = orange, MnO₄⁻ = purple.
d-electrons can be removed easily. Fe shows +2 and +3, Mn shows +2 to +7, Cu shows +1 and +2. This enables catalysis and redox chemistry.
Transition metals and their ions are excellent catalysts due to variable oxidation states. Fe in Haber process, Pt in catalytic converters, V₂O₅ in Contact process.
Unpaired d-electrons make many transition metals and compounds paramagnetic (attracted to magnetic field). Fe, Co, Ni are ferromagnetic — permanently magnetic.
Form complex ions with ligands (Lewis bases donating lone pairs). e.g. [Cu(H₂O)₄]²⁺, [Fe(CN)₆]³⁻. Coordination number 6 is most common.
Small atomic radii, strong metallic bonding. High melting points, densities, and hardness. W has highest melting point (3422°C). Os is densest metal (22.59 g/cm³).
Written using the Aufbau principle: fill from lowest energy level. Notation: 1s², 2s², 2p⁶ … etc.
| Element | Symbol | Z | Configuration |
|---|---|---|---|
| Hydrogen | H | 1 | 1s¹ |
| Helium | He | 2 | 1s² |
| Lithium | Li | 3 | 1s² 2s¹ |
| Beryllium | Be | 4 | 1s² 2s² |
| Boron | B | 5 | 1s² 2s² 2p¹ |
| Carbon | C | 6 | 1s² 2s² 2p² |
| Nitrogen | N | 7 | 1s² 2s² 2p³ |
| Oxygen | O | 8 | 1s² 2s² 2p⁴ |
| Fluorine | F | 9 | 1s² 2s² 2p⁵ |
| Neon | Ne | 10 | 1s² 2s² 2p⁶ |
| Sodium | Na | 11 | 1s² 2s² 2p⁶ 3s¹ |
| Magnesium | Mg | 12 | 1s² 2s² 2p⁶ 3s² |
| Aluminium | Al | 13 | 1s² 2s² 2p⁶ 3s² 3p¹ |
| Silicon | Si | 14 | 1s² 2s² 2p⁶ 3s² 3p² |
| Phosphorus | P | 15 | 1s² 2s² 2p⁶ 3s² 3p³ |
| Sulfur | S | 16 | 1s² 2s² 2p⁶ 3s² 3p⁴ |
| Chlorine | Cl | 17 | 1s² 2s² 2p⁶ 3s² 3p⁵ |
| Argon | Ar | 18 | 1s² 2s² 2p⁶ 3s² 3p⁶ |
| Potassium | K | 19 | [Ar] 4s¹ |
| Calcium | Ca | 20 | [Ar] 4s² |
| Scandium | Sc | 21 | [Ar] 3d¹ 4s² |
| Titanium | Ti | 22 | [Ar] 3d² 4s² |
| Vanadium | V | 23 | [Ar] 3d³ 4s² |
| Chromium ★ exception | Cr | 24 | [Ar] 3d⁵ 4s¹ |
| Manganese | Mn | 25 | [Ar] 3d⁵ 4s² |
| Iron | Fe | 26 | [Ar] 3d⁶ 4s² |
| Cobalt | Co | 27 | [Ar] 3d⁷ 4s² |
| Nickel | Ni | 28 | [Ar] 3d⁸ 4s² |
| Copper ★ exception | Cu | 29 | [Ar] 3d¹⁰ 4s¹ |
| Zinc | Zn | 30 | [Ar] 3d¹⁰ 4s² |
Cr prefers [Ar] 3d⁵ 4s¹ (half-filled d = extra stability) over [Ar] 3d⁴ 4s². Cu prefers [Ar] 3d¹⁰ 4s¹ (fully-filled d) over [Ar] 3d⁹ 4s². Half-filled and fully-filled d subshells are extra stable due to exchange energy and symmetry.
When transition metals form cations, ns electrons are removed first (before (n−1)d electrons), regardless of which filled last. This is because in cations the 4s is higher energy than 3d.
Transition metal ions do not usually attain a noble gas configuration (unlike main-group ions). This is because the ns orbital empties first into the (n−1)d orbitals when ionising. The resulting partial d-filling explains their colour, variable oxidation states, and magnetic properties.
Enter an unbalanced equation. Use + to separate substances, → or -> for the arrow. Coefficients will be found automatically.
Based on: C₁V₁/n₁ = C₂V₂/n₂. Leave the unknown blank.
Enter percentage composition (must total ~100%). Optionally enter molar mass to get molecular formula.
The complete periodic table of all 118 elements. Click any element for detailed information including electron configuration, uses, and key facts. Color-coded by element type.
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