Chemistry: Elements, Reactions and Compounds

Acids, bases and alkalis

An acid is a substance that produces hydrogen ions, \mathrm{H^+} (present in water as hydronium ions, \mathrm{H_3O^+}), in aqueous solution. A base is a substance that neutralises an acid to form a salt and water, and an alkali is a base that dissolves in water. Every alkali is a base, but not every base is an alkali, because insoluble bases such as copper(II) oxide and iron(III) hydroxide do not dissolve in water.

Alkalis have characteristic properties: they feel soapy to touch, taste bitter and are corrosive in nature.

Acids in nature and daily life. Many acids occur naturally in plants, animals and food. Each of the following should be remembered with its source:

• Methanoic acid (formic acid), \mathrm{HCOOH} — the simplest carboxylic acid. It is the irritant in the stinging hairs of nettle leaves and in ant stings, which causes the burning pain. Example: a nettle sting burns because of methanoic acid.
• Tartaric acid — the main acid of tamarind and grapes. Example: the sourness of tamarind pulp comes mainly from tartaric acid.
• Citric acid — the main acid of citrus fruits such as lemon.
• Ascorbic acid — vitamin C, found in orange juice.
• Acetic acid — the main constituent of vinegar, which is a 5 to 8 per cent solution of acetic acid in water.
• Oxalic acid — is not the acid of tamarind. Do not confuse it with tartaric acid.
• Hydrochloric acid, \mathrm{HCl} — a mineral (inorganic) acid secreted by the parietal cells of the stomach lining, which makes gastric juice strongly acidic and helps digest food and kill microbes.
• Lactic acid and uric acid — organic acids, not mineral acids. Uric acid is a nitrogenous waste product.

Acid rain. Oxides of sulphur and nitrogen, released by burning fossil fuels and by industrial processes, react with water and oxygen in the atmosphere to form sulphuric acid and nitric acid. These acids fall in rain and lower its pH. Minerals, dust particles and soil solutions dissolved in rain do not make it acidic to any significant extent.

pH scale and indicators. The pH scale runs from 0 to 14 at 25 °C. A solution with pH equal to 7 is neutral, a solution with pH less than 7 is acidic, and a solution with pH greater than 7 is basic. Hypotonic describes osmotic pressure, not acidity, so it should not be used for pH.

• Litmus — a natural dye obtained from lichens, which are symbiotic associations of a fungus with an alga or a cyanobacterium.
• Universal indicator — a mixture of several dyes. Its colour changes continuously with the hydrogen ion concentration, so each pH value gives a different colour. Because the colour indicates the approximate pH, it roughly shows how strong a given acid or base is.

Example ordering by increasing pH: lemon juice (about pH 2, strongly acidic) < coffee (about pH 5, weakly acidic) < blood (about pH 7.4, slightly basic) < milk of magnesia (about pH 10.5, basic).

Mild and strong bases. Milk of magnesia is a suspension of magnesium hydroxide in water. Magnesium hydroxide is sparingly soluble and only weakly basic, so it is a mild base. It neutralises excess stomach acid in indigestion. A strong base is highly soluble and ionises completely.

Amphoteric oxides. An amphoteric oxide reacts with both acids and bases to form a salt and water. Aluminium oxide, \mathrm{Al_2O_3}, is the typical example. Magnesium oxide and sodium oxide are basic oxides, and tetraphosphorus decoxide, \mathrm{P_4O_{10}}, is an acidic oxide.

Safe dilution of acids. Diluting a concentrated acid lowers the number of \mathrm{H_3O^+} ions per unit volume. Because mixing releases a large amount of heat, the concentrated acid must be added slowly to water, with stirring. Adding water to concentrated acid can cause local boiling and spattering.

Carbon and its compounds

Allotropes of carbon. Carbon exists in several allotropic forms with very different properties, because the arrangement of its atoms differs.

• Diamond — very hard, because each carbon is bonded tetrahedrally in a three-dimensional network. All its valence electrons are localised in bonds, so it does not conduct electricity.
• Graphite — soft, smooth and slippery, so it is used as a dry lubricant. Its atoms form layers of hexagonal rings, and each carbon is bonded to three neighbours. This leaves one delocalised electron per atom, which moves freely within the layers and makes graphite a good conductor of electricity.
• Fullerenes — such as \mathrm{C_{60}}, which are generally poor conductors in their pure form.
• Graphene — a single layer of graphite, one atom thick, which makes it the thinnest known material. It is almost transparent and highly conducting because of its delocalised \pi electrons. It has zero band gap and behaves as a semimetal, so it is not a wide band-gap semiconductor.

Occurrence. Carbon makes up only a small fraction of the Earth's crust, about 0.025 per cent by mass, mainly in carbonate minerals such as limestone.

Combustion of hydrocarbons. With sufficient oxygen (excess air), hydrocarbons burn completely to carbon dioxide and water and give a clean blue flame. With insufficient oxygen, combustion is incomplete, unburnt carbon particles escape as soot, and the flame becomes yellow and smoky. Fuels with a high carbon to hydrogen ratio are especially prone to this. Unsaturated hydrocarbons have a high carbon to hydrogen ratio and therefore burn with a yellow, sooty flame.

• Naphthalene, \mathrm{C_{10}H_8} — has a high carbon to hydrogen ratio, so its incomplete combustion gives a yellow sooty flame.

Saturated and unsaturated compounds. Saturated hydrocarbons (alkanes and cycloalkanes) contain only single bonds, while unsaturated hydrocarbons (alkenes, alkynes and aromatic compounds) contain at least one double or triple bond. Cyclohexane, \mathrm{C_6H_{12}}, has 6 C–C and 12 C–H single bonds, which gives 18 single bonds and no double or triple bonds.

Hydrogenation. Unsaturated hydrocarbons are reduced to saturated hydrocarbons by adding hydrogen in the presence of a finely divided nickel catalyst. Example: the conversion of vegetable oils to vanaspati ghee uses this process.

Structure of benzene, \mathrm{C_6H_6}. Each carbon atom is sp^2 hybridised and forms sigma bonds to two carbons and one hydrogen, with bond angles of 120^\circ. The remaining p orbitals overlap to form a delocalised \pi electron cloud above and below the planar ring. The C–C bond length lies between typical single and double bond lengths.

Octane number. The octane number measures a fuel's resistance to knocking. Shorter straight-chain alkanes have higher octane numbers than longer ones, and cyclic hydrocarbons such as cyclohexane resist knocking better than straight-chain alkanes of similar size. Example order: butane > cyclohexane > pentane > hexane.

Cetane. Cetane is the common name of n-hexadecane, \mathrm{C_{16}H_{34}}, which is the reference compound for the cetane number scale used to rate diesel fuel.

Methanol toxicity. Methanol is not directly toxic. In the liver it is oxidised to methanal (formaldehyde), which coagulates the protoplasm of cells.

Carboxylic acids. A carboxylic acid contains the carboxyl group, –COOH, attached to an alkyl group. Example: propanoic acid, \mathrm{C_2H_5COOH}. Propanol (\mathrm{C_3H_7OH}) is an alcohol, propane (\mathrm{C_3H_8}) is an alkane, and acetone (\mathrm{CH_3COCH_3}) is a ketone, so none of these is a carboxylic acid.

Metals and non-metals properties

Elements are classified as metals, non-metals and metalloids according to their physical and chemical properties. Metals are generally lustrous, malleable, ductile, sonorous and good conductors of heat and electricity, and their oxides are basic. Non-metals are usually dull, brittle and poor conductors, and their oxides are acidic or neutral. These are general trends, and the exceptions must be learned along with them.

• Lustre — the shiny appearance of a metal surface. Some non-metals are also lustrous in the solid state, and iodine is the standard example.
• Malleability and ductility — malleability is the ability to be beaten into sheets, and ductility is the ability to be drawn into wire. Gold is the most malleable and ductile metal.
• Sonorous — the property of giving a ringing sound when struck. This is why metals are used for bells.
• Thermal conductivity — silver and copper are among the best conductors of heat, while lead and mercury conduct heat poorly.
• Oxides — non-metal oxides such as \mathrm{CO_2} and \mathrm{SO_2} dissolve in water to form acids, so they are acidic oxides.
• Tarnishing — silver blackens in air because it reacts with sulphur compounds to form black silver sulphide, \mathrm{Ag_2S}.
• Storage of reactive metals — sodium is highly reactive with air and moisture, so it is kept immersed in kerosene oil.
• Green patina — copper turns green in moist air containing \mathrm{CO_2} because it forms basic copper carbonate, which is a carbonate of copper in the cupric state.
• Solder — a low-melting alloy of lead and tin used to join metal parts.
• Reactivity with dilute HCl — a metal above hydrogen in the reactivity series displaces hydrogen from dilute acid, and a more vigorous reaction means a more reactive metal. In this series, magnesium is more reactive than aluminium, aluminium is more reactive than zinc, and zinc is more reactive than iron, so the order is \mathrm{Mg > Al > Zn > Fe}.

Sodium and calcium compounds

Sodium and calcium compounds are important in both laboratory work and daily life, and their reactions are frequently tested through their applications.

• Sodium hydrogen carbonate in fire extinguishers — in a soda-acid extinguisher, sodium hydrogen carbonate solution reacts with sulphuric acid to release carbon dioxide, which pushes the liquid out onto the fire. The reaction is 2\mathrm{NaHCO_3 + H_2SO_4 \rightarrow Na_2SO_4 + 2H_2O + 2CO_2}.
• Sodium carbonate in water softening — hard water contains \mathrm{Ca^{2+}} and \mathrm{Mg^{2+}} ions. Sodium carbonate precipitates these ions as insoluble carbonates, for example \mathrm{Ca^{2+} + CO_3^{2-} \rightarrow CaCO_3\downarrow}, so it softens both temporary and permanent hardness.
• Quicklime and slaked lime — quicklime, \mathrm{CaO}, reacts vigorously and exothermically with water to form slaked lime, \mathrm{Ca(OH)_2}. A solution of slaked lime, called lime water, is used for whitewashing walls, whereas quicklime itself is not applied directly. Calcium hydroxide is an inorganic compound.
• Setting of whitewash — slaked lime slowly absorbs carbon dioxide from air to form a thin, hard, shiny layer of calcium carbonate. The equation is \mathrm{Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O}, where \mathrm{CO_2} is a gas and \mathrm{CaCO_3} is a solid.
• Drying agent — anhydrous calcium chloride is placed in a guard tube to remove water vapour from hydrogen chloride gas, because it absorbs moisture without reacting with the gas.

In statement-based questions, the key is to distinguish the form that is actually used (slaked lime solution, not quicklime) and the species that actually forms the coating (calcium carbonate, not calcium oxide).

Redox reactions

A redox reaction is one in which oxidation and reduction occur together. Oxidation is the gain of oxygen, the loss of hydrogen, or the loss of electrons, and reduction is the reverse. A reaction that involves no change in oxidation states is not a redox reaction, even if it forms a precipitate.

• Common oxidation processes — the spoilage of butter, rusting of iron, and the souring of wine are all oxidation processes. In wine, air oxidises ethanol to acetic acid.
• Non-oxidation changes — opening a soda bottle releases dissolved carbon dioxide because pressure falls, which is not an oxidation reaction. Mixing barium chloride and sodium sulphate gives a white precipitate by exchange of ions, with no change in oxidation numbers.
• Thermite reaction — aluminium reduces iron(III) oxide in a strongly exothermic displacement reaction: \mathrm{Fe_2O_3 + 2Al \rightarrow 2Fe + Al_2O_3}. Aluminium is the reducing agent because it is itself oxidised. The heat produces molten iron, which is used to weld railway tracks.
• Anodizing — aluminium is made the anode in an electrolytic cell, so oxygen is liberated at the anode and a thick, adherent oxide film forms on the metal. The film protects against corrosion, which is why anodized aluminium is used in the aircraft industry. Aluminium, titanium and magnesium can be anodized.
• Variable oxidation numbers — transition metals such as iron show more than one oxidation state, commonly +2 and +3. Sodium, lithium and calcium form only one stable oxidation state each.

Chemical reactions and types

A chemical equation must be balanced so that the number of atoms of each element is the same on both sides, in accordance with conservation of mass.

• Combustion — a rapid reaction with oxygen that releases heat and light. Magnesium burns in air to form white magnesium oxide, \mathrm{2Mg + O_2 \rightarrow 2MgO}, which makes this a combustion reaction.
• Balancing a redox-type equation with steam — to balance \mathrm{Fe + H_2O \rightarrow Fe_3O_4 + H_2}, first place three iron atoms on the left to match \mathrm{Fe_3O_4}. Four oxygen atoms need four water molecules, which give eight hydrogen atoms and therefore four \mathrm{H_2} molecules, giving \mathrm{3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2}.
• Physical and chemical change — a physical change alters only the state or form of a substance, while a chemical change produces new substances. Burning a candle contains both, since the wax first melts (physical) and then undergoes combustion to form new substances (chemical).
• Denaturation — heat unfolds the proteins in egg white, such as ovalbumin, and the unfolded molecules coagulate into an opaque white solid. This is thermal denaturation of proteins, not of fats or carbohydrates.
• Types of reactions — in a double displacement (precipitation) reaction, the ions of two aqueous compounds exchange partners. In \mathrm{K_2SO_4 + BaCl_2 \rightarrow 2KCl + BaSO_4\downarrow}, the insoluble barium sulphate precipitates while soluble potassium chloride remains in solution.