Chemistry

Matter

Matter is anything that has mass and occupies space. It exists in three common states — solid (definite shape and volume), liquid (definite volume, no fixed shape) and gas (neither) — with plasma (ionised gas, as in the Sun and stars, lightning, neon signs) and Bose–Einstein condensate as further states.

• Pure substances: elements (one kind of atom: hydrogen, iron) and compounds (two or more elements chemically combined in a fixed ratio: water H₂O, carbon dioxide CO₂).
• Mixtures: homogeneous (solutions, air, brass) and heterogeneous (sand and water, milk). The components of a mixture retain their properties and can be separated by physical methods (filtration, distillation, evaporation, sublimation, chromatography, magnetic separation, sedimentation and decantation, centrifugation).
• A molecule is the smallest particle of a substance capable of independent existence; an atom is the smallest particle of an element that takes part in a chemical reaction.

Laws of chemical combination

• Law of conservation of mass (Lavoisier): mass is neither created nor destroyed in a chemical reaction.
• Law of constant (definite) proportions (Proust): a compound always contains the same elements in the same proportion by mass (water is always 1 part hydrogen to 8 parts oxygen by mass).
• Law of multiple proportions (Dalton) and Gay-Lussac's law of gaseous volumes.
• Avogadro's law: equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules. One mole of any substance contains $6.022\times 10^{23}$ particles (Avogadro's number); one mole of a gas occupies 22.4 litres at STP.

Development of atomic models

Drawbacks of Rutherford's model: it could not explain the stability of the atom (an orbiting electron should radiate energy and spiral into the nucleus). Bohr's model explained the hydrogen spectrum but failed for larger atoms.

Subatomic particles

• The nucleus is about 10⁻¹⁵ m across, about 100,000 times smaller than the atom, but holds almost all the mass.
• Protons and neutrons are nucleons; they are made of quarks. Hydrogen (protium) is the only atom with no neutron.
• An atom is electrically neutral: number of protons = number of electrons.

Atomic number, mass number, isotopes

• Atomic number ($Z$) = number of protons. It identifies the element.
• Mass number ($A$) = protons + neutrons (nucleons). Number of neutrons $=A-Z$.
• Notation: ${}_Z^{A}X$; e.g. ${}_6^{12}\text{C}$: 6 protons, 6 neutrons, 6 electrons.

Uses of isotopes: ¹⁴C for radiocarbon dating; ⁶⁰Co for cancer treatment; ¹³¹I for thyroid disorders; ²³⁵U as nuclear fuel; deuterium (heavy water) as a moderator in reactors; ²⁴Na for detecting blood clots; isotopes as tracers in agriculture and industry.

Why classify elements?

More than 118 elements are known. Classifying them into a table makes it easy to study their properties and to predict those of unknown elements.

History

Defects of Mendeleev's table: the position of hydrogen was uncertain; isotopes had no place; some elements were placed out of order of atomic mass (e.g. cobalt before nickel, tellurium before iodine); the noble gases were not included (discovered later).

The modern periodic table

• Periods: the 7 horizontal rows. Elements in a period have the same number of shells, and properties change gradually across it. The number of elements in periods 1 to 7 is 2, 8, 8, 18, 18, 32, 32 (incomplete).
• Groups: the 18 vertical columns. Elements in a group have the same number of valence electrons and similar chemical properties.
• Blocks by the last electron's orbital: s-block (groups 1–2), p-block (groups 13–18), d-block (groups 3–12, the transition elements), f-block (lanthanides and actinides, placed at the bottom).
• Metals lie to the left and centre, non-metals to the top right, and metalloids (semi-metals: boron, silicon, germanium, arsenic, antimony, tellurium) along the zig-zag line between them.

Important families

Lanthanides (rare earths, atomic numbers 57–71) and actinides (89–103) form the f-block; transuranic elements (beyond uranium) are man-made.

Physical and chemical change

Signs of a chemical reaction: change of colour, evolution of a gas, formation of a precipitate, change of temperature (heat or light given out), change of state.

Chemical equations

A chemical equation represents a reaction with symbols and formulae. Reactants are on the left, products on the right, with an arrow between them.

• Physical states: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution. An arrow ↑ shows a gas escaping; ↓ a precipitate.
• Conditions such as heat, catalyst, light or pressure are written on the arrow.

Balancing

By the law of conservation of mass, the number of atoms of each element must be the same on both sides. Balance by adding coefficients, never by changing the formulae.

Unbalanced: Fe + H₂O → Fe₃O₄ + H₂. Balanced: 3Fe + 4H₂O → Fe₃O₄ + 4H₂.

Types of chemical reactions

1. Combination (synthesis)

Two or more substances combine to form a single product: $\text{A}+\text{B}\to \text{AB}$.

• $\text{CaO}+{\text{H}}_2\text{O}\to {\text{Ca(OH)}}_2$ (quicklime + water → slaked lime, exothermic; used in whitewashing).
• $2\text{Mg}+{\text{O}}_2\to 2\text{MgO}$ (burning magnesium with a dazzling white light).
• $\text{C}+{\text{O}}_2\to {\text{CO}}_2$; ${\text{N}}_2+3{\text{H}}_2\to 2{\text{NH}}_3$ (Haber process).

2. Decomposition

A single compound breaks into two or more simpler substances: $\text{AB}\to \text{A}+\text{B}$.

• Thermal decomposition: ${\text{CaCO}}_3\stackrel{\text{heat}}{\to }\text{CaO}+{\text{CO}}_2$ (limestone → quicklime; used in cement making). Also $2{\text{Pb(NO}}_3{)}_2\to 2\text{PbO}+4{\text{NO}}_2+{\text{O}}_2$ (brown fumes).
• Electrolytic decomposition: $2{\text{H}}_2\text{O}\to 2{\text{H}}_2+{\text{O}}_2$ on passing electricity (hydrogen at the cathode, twice the volume of oxygen at the anode).
• Photolytic decomposition: silver chloride and silver bromide turn grey in sunlight (basis of black-and-white photography): $2\text{AgCl}\stackrel{\text{light}}{\to }2\text{Ag}+{\text{Cl}}_2$.

3. Displacement

A more reactive element displaces a less reactive one from its compound: $\text{A}+\text{BC}\to \text{AC}+\text{B}$.

• $\text{Fe}+{\text{CuSO}}_4\to {\text{FeSO}}_4+\text{Cu}$ (an iron nail in blue copper sulphate solution turns brownish and the solution turns pale green).
• $\text{Zn}+{\text{CuSO}}_4\to {\text{ZnSO}}_4+\text{Cu}$.

4. Double displacement

Two compounds exchange ions: $\text{AB}+\text{CD}\to \text{AD}+\text{CB}$.

• Precipitation: ${\text{Na}}_2{\text{SO}}_4+{\text{BaCl}}_2\to {\text{BaSO}}_4\downarrow +2\text{NaCl}$ (white precipitate).
• Neutralisation: acid + base → salt + water: $\text{HCl}+\text{NaOH}\to \text{NaCl}+{\text{H}}_2\text{O}$.

5. Oxidation and reduction (redox)

• Oxidation: gain of oxygen, or loss of hydrogen, or loss of electrons (OIL).
• Reduction: loss of oxygen, or gain of hydrogen, or gain of electrons (RIG). Remember OIL RIG.
• They occur together in a redox reaction. The substance that gets oxidised is the reducing agent; the one reduced is the oxidising agent.
• $\text{CuO}+{\text{H}}_2\to \text{Cu}+{\text{H}}_2\text{O}$: CuO is reduced (loses oxygen), H₂ is oxidised.
• Examples of oxidising agents: oxygen, potassium permanganate, hydrogen peroxide, potassium dichromate, chlorine. Reducing agents: hydrogen, carbon, carbon monoxide, metals.

Acids

An acid is a substance that gives hydrogen ions (H⁺) (hydronium ions, H₃O⁺) in water. Acids are sour in taste, turn blue litmus red, and react with metals, carbonates and bases.

Properties

• Acid + metal → salt + hydrogen: $\text{Zn}+{\text{H}}_2{\text{SO}}_4\to {\text{ZnSO}}_4+{\text{H}}_2$. (The hydrogen gas burns with a "pop" sound.)
• Acid + metal carbonate/bicarbonate → salt + water + carbon dioxide (brisk effervescence; the gas turns lime water milky).
• Acid + base → salt + water (neutralisation).
• Acids conduct electricity in solution, since they produce ions.
• Concentrated acids are corrosive; dilute an acid by adding the acid slowly to water (never water to acid), since the process gives out a lot of heat.

Common acids

Basicity: the number of replaceable H⁺ ions: HCl (monobasic), H₂SO₄ (dibasic), H₃PO₄ (tribasic).

Bases and alkalis

A base gives hydroxide ions (OH⁻) in water, or accepts H⁺. A base soluble in water is an alkali (NaOH, KOH, Ca(OH)₂, NH₄OH). Bases are bitter, feel soapy, turn red litmus blue, and conduct electricity in solution.

• Base + acid → salt + water.
• Base + ammonium salts liberates ammonia.
• Strong alkalis are corrosive.

Common bases

Antacids (sodium bicarbonate, magnesium hydroxide, aluminium hydroxide) relieve the excess acid in the stomach.

Indicators

An indicator changes colour in acid or alkaline solution.

Olfactory indicators: onion, vanilla and clove oils lose their smell in a base.

The pH scale

The pH is a measure of the hydrogen-ion concentration: $\text{pH}=-{\log }_{10}[{\text{H}}^{+}]$. The scale runs from 0 to 14:

A change of one pH unit corresponds to a tenfold change in $[{\text{H}}^{+}]$. The pH is measured with pH paper or a pH meter.

Acid rain has a pH below 5.6, caused by sulphur dioxide and nitrogen oxides; it damages buildings (the Taj Mahal), plants and aquatic life.

pH in daily life

• Tooth decay starts when the pH in the mouth falls below about 5.5; toothpastes are basic and neutralise the acid.
• Soil pH decides plant growth: acidic soils are treated with lime; alkaline soils with gypsum or organic matter.
• Indigestion is due to excess HCl in the stomach and is relieved by antacids.
• Bee sting (methanoic acid) is treated with baking soda; a wasp sting (alkaline) with vinegar.
• The pH of blood is maintained by buffers.

Metals and non-metals

Chemical properties of metals

• With oxygen: metal + O₂ → metal oxide (basic). $2\text{Mg}+{\text{O}}_2\to 2\text{MgO}$. Sodium and potassium burn in air (stored under kerosene); iron burns as filings; copper turns black (CuO) on heating; aluminium forms a protective oxide layer (anodising thickens it). Amphoteric oxides (Al₂O₃, ZnO) react with both acids and bases.
• With water: Na and K react violently with cold water; Ca less so; Mg with hot water; Al, Zn, Fe with steam; Cu, Ag, Au do not react.
• With dilute acids: metals above hydrogen liberate H₂; copper, silver and gold do not (except with oxidising acids like nitric acid).
• With salt solutions: displacement of a less reactive metal by a more reactive one.

Reactivity series

K > Na > Ca > Mg > Al > Zn > Fe > Pb > (H) > Cu > Hg > Ag > Au

Ionic compounds

Formed when a metal transfers electrons to a non-metal. They are hard, brittle crystalline solids with high melting and boiling points, are soluble in water and conduct electricity in the molten state and in solution (not as solids).

Occurrence and extraction of metals

• Minerals are naturally occurring compounds of metals; an ore is a mineral from which the metal can be profitably extracted. Gangue is the impurity in the ore.

Steps in extraction

1. Concentration of the ore (hand-picking, washing, magnetic separation, froth flotation for sulphide ores).
2. Conversion to oxide: calcination (heating carbonate ores in limited or no air: ZnCO₃ → ZnO + CO₂) or roasting (heating sulphide ores in excess air: 2ZnS + 3O₂ → 2ZnO + 2SO₂).
3. Reduction of the oxide to the metal:
   • Low-reactivity metals (Hg, Cu): heating alone: 2HgS + 3O₂ → 2HgO + 2SO₂; 2HgO → 2Hg + O₂.
   • Medium-reactivity metals (Zn, Fe, Pb): reduction by carbon (coke) or carbon monoxide: ZnO + C → Zn + CO; Fe₂O₃ + 3CO → 2Fe + 3CO₂ (blast furnace). The thermite reaction: Fe₂O₃ + 2Al → Al₂O₃ + 2Fe (used to weld railway tracks).
   • High-reactivity metals (K, Na, Ca, Mg, Al): electrolysis of the molten compound — aluminium by electrolysis of alumina dissolved in molten cryolite (Hall–Héroult process); sodium by electrolysis of molten NaCl (Down's process).
4. Refining: electrolytic refining (impure metal as anode, pure metal as cathode: copper, zinc, silver, gold); distillation (Zn, Hg).

Iron and steel: pig iron (cast iron) has 3–4 % carbon; wrought iron is almost pure; steel has 0.1–1.5 % carbon. Blast furnace flux: limestone; slag: calcium silicate.

Corrosion and alloys

Corrosion (rusting of iron, tarnishing of silver, green coating on copper) and its prevention are described in the Chemical Reactions note.

Alloys are homogeneous mixtures of a metal with other metals or non-metals; they are harder, stronger, more corrosion-resistant and have lower melting points than pure metals.

Non-metals

• Hydrogen, carbon, nitrogen, oxygen, phosphorus, sulphur, the halogens and noble gases.
• With oxygen: form acidic or neutral oxides (CO₂, SO₂, P₂O₅). They accept electrons to form anions; not in the reactivity series.
• Sulphur is used in vulcanisation of rubber, gunpowder and matches; phosphorus in matches (red phosphorus on the box) and fertilisers; nitrogen in fertilisers and ammonia; chlorine for disinfecting water.

Carbon and its compounds

Carbon (valency 4) forms a vast number of compounds because of catenation (the ability to form long chains and rings by bonding with itself) and its tetravalency; it forms strong single, double and triple covalent bonds. It occurs as diamond, graphite, fullerenes (C₆₀, buckminsterfullerene), graphene and carbon nanotubes.

• Covalent compounds have low melting and boiling points, are poor conductors of electricity and are generally insoluble in water.
• Hydrocarbons: compounds of carbon and hydrogen. Saturated (alkanes): single bonds only, general formula ${\text{C}}_n{\text{H}}_{2n+2}$ (methane CH₄, ethane C₂H₆, propane C₃H₈, butane C₄H₁₀). Unsaturated: alkenes (double bond, ${\text{C}}_n{\text{H}}_{2n}$: ethene, propene) and alkynes (triple bond, ${\text{C}}_n{\text{H}}_{2n-2}$: ethyne/acetylene, used in welding). Aromatic: benzene (C₆H₆).
• Homologous series: a family of compounds with the same functional group and general formula, with successive members differing by a –CH₂– unit (14 u).
• Isomers: compounds with the same molecular formula but different structures (butane and isobutane).
• Functional groups: alcohol (–OH), aldehyde (–CHO), ketone (>C=O), carboxylic acid (–COOH), halo (–Cl, –Br), ester (–COO–), amine (–NH₂).

• Combustion: carbon compounds burn in oxygen to give CO₂, water, heat and light; saturated hydrocarbons burn with a clean blue flame; unsaturated ones with a yellow sooty flame (incomplete burning).
• Addition reaction: unsaturated hydrocarbons add hydrogen in the presence of nickel — hydrogenation of vegetable oils to make vanaspati ghee.
• Substitution reaction: alkanes react with chlorine in sunlight.

Ethanol reactions

• Ethanol + sodium → sodium ethoxide + hydrogen.
• Ethanol → ethene (with conc. H₂SO₄ at 443 K, dehydration).
• Ethanol + ethanoic acid → ester (esterification).
• Ethanol → ethanoic acid (oxidation with alkaline KMnO₄ or acidified K₂Cr₂O₇).

Soaps and detergents

Soap is the sodium or potassium salt of a long-chain fatty acid; it is made by the saponification of fats or oils with NaOH. It has a hydrophobic tail and a hydrophilic head, and forms micelles around grease. Soaps form scum in hard water (they react with calcium and magnesium salts); detergents (sodium alkyl sulphonates) do not.

Fuels