Chemical Bonding & Reactions

Look at the noble gases: helium, neon, argon. They never join with anything, and they happily float around as single atoms. Every other atom is restless. Why? Because noble gases have a full outer shell (eight electrons, or two for helium), and a full shell is a state of great stability. All other atoms try to copy this arrangement by gaining, losing or sharing electrons. That urge is the whole reason for chemical bonding. In 1916 Kossel and Lewis put it as the octet rule: atoms combine so that each ends up with eight electrons in its outer shell (hydrogen settles for two, the duet). The electrons in the outer shell are the valence electrons, and only they take part in bonding.

An ionic bond (also called an electrovalent bond) forms when one atom hands over electrons completely to another. The best example is common salt. Sodium (2, 8, 1) has one electron too many for a full shell, so it gives it away and becomes Na⁺ (2, 8). Chlorine (2, 8, 7) is one short, so it takes that electron and becomes Cl⁻ (2, 8, 8). Now the two ions have opposite charges and attract each other strongly. That electrostatic attraction between oppositely charged ions is the ionic bond. Notice the pattern: a metal (few outer electrons, low ionisation energy) hands electrons to a non-metal (almost full shell, high electron affinity). The number of electrons lost or gained is the electrovalency.

Why does a higher charge mean a higher melting point? The strength of an ionic lattice depends on two things: the charges on the ions and how close they sit. NaCl has +1 and -1 ions; CaO has +2 and -2 ions. The pull between two charges is multiplied when each charge doubles, so CaO is held together far more tightly than NaCl. That is why NaCl melts at about 1074 K but CaO needs about 2850 K. Quick guide: larger charge and smaller ions give a stronger lattice and a higher melting point.

Writing formulae of ionic compounds. Use the criss-cross method: write the symbols with their charges, then swap the numbers and drop the signs. Mg²⁺ with PO₄³⁻ gives Mg₃(PO₄)₂; Al³⁺ with SO₄²⁻ gives Al₂(SO₄)₃; Ca²⁺ with OH⁻ gives Ca(OH)₂. Put a polyatomic ion in brackets when you need more than one of it. Some metals have more than one valency: iron gives Fe²⁺ (ferrous) and Fe³⁺ (ferric); copper gives Cu⁺ (cuprous) and Cu²⁺ (cupric). The lower-valency ion takes the ending -ous and the higher takes -ic.

What happens when two non-metals meet? Neither wants to give up electrons, because both are close to a full shell. So they compromise: they share electrons. Each hydrogen atom has one electron and needs one more; two hydrogen atoms share a pair and both feel complete. A shared pair of electrons is a covalent bond. Sharing one pair gives a single bond (H–H, Cl–Cl), two pairs a double bond (O=O, CO₂), three pairs a triple bond (N≡N). The more pairs shared, the shorter and stronger the bond. This is why the nitrogen molecule, with its triple bond, is so hard to break and nitrogen gas is so unreactive.

Polar and non-polar covalent bonds. When two identical atoms share electrons (H–H, Cl–Cl), the pair sits exactly in the middle: a non-polar bond. When the atoms differ (H–Cl), the more electronegative atom pulls the pair towards itself and becomes slightly negative, leaving the other slightly positive: a polar bond. Water is polar and bent; carbon dioxide has polar bonds but is a straight line, so the pulls cancel and the molecule is non-polar. Ionic and covalent are the two ends of a scale: the bigger the difference in electronegativity, the more ionic the bond.

Coordinate (dative) bond. Sometimes both electrons of the shared pair come from the same atom. The atom that gives the pair is the donor and the atom that accepts it is the acceptor; the bond is shown by an arrow (donor → acceptor). Once formed, it behaves like any other covalent bond. The standard example is the ammonium ion: nitrogen in NH₃ has a lone pair, and a hydrogen ion (H⁺, with no electron) accepts it to make NH₄⁺. In the same way water takes up H⁺ to form the hydronium ion H₃O⁺, and carbon monoxide, ozone and many complex ions contain coordinate bonds.

Hydrogen bond. When hydrogen is attached to a very electronegative small atom (fluorine, oxygen or nitrogen), it becomes strongly positive and is attracted to a lone pair on a neighbouring molecule. This weak attraction is the hydrogen bond. It is much weaker than a covalent bond, yet it explains many everyday facts. Water boils at 100 °C although H₂S, which is heavier, is a gas at room temperature, because water molecules cling to one another by hydrogen bonds. Ice floats on water because hydrogen bonds hold water molecules in an open cage with gaps, so ice is less dense than liquid water. The two strands of DNA are also held together by hydrogen bonds between the bases.

Metallic bond. In a metal the outer electrons do not belong to any one atom. They form a shared sea of electrons that flows around a lattice of positive metal ions. This free sea explains why metals conduct heat and electricity, why they shine, and why they can be hammered into sheets (malleable) or drawn into wires (ductile) without breaking: when a layer slips, the sea simply flows along and the bonding continues.

Forces between molecules decide boiling points. Covalent compounds exist as separate molecules. Inside a molecule the atoms are held strongly by covalent bonds, but between molecules only weak forces act. When a covalent substance melts or boils, only these weak intermolecular forces are overcome and the covalent bonds stay intact. That is why covalent compounds have low melting and boiling points. The weakest are the van der Waals (London dispersion) forces, which exist between all molecules, including noble gas atoms; they are the reason noble gases and non-polar gases can be liquefied at all. In the van der Waals gas equation, the constant 'a' corrects for these attractive forces and the constant 'b' for the actual volume of the molecules. In general: ionic bond > hydrogen bond > dipole-dipole attraction > van der Waals forces.

A chemical equation is the shorthand of chemistry: it tells in one line what reacts, what is formed and under what condition. Reactants are written on the left, products on the right, joined by an arrow: Mg + O₂ → MgO is a skeleton equation. It is not yet correct, because there are two oxygen atoms on the left and only one on the right. Atoms are never lost or created in a reaction, so a correct equation must have the same number of each kind of atom on both sides. Such an equation is balanced: 2Mg + O₂ → 2MgO. Physical states are shown in brackets: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution. An upward arrow ↑ shows a gas escaping, a downward arrow ↓ a precipitate, and Δ above the arrow shows heating.

Why must we balance an equation? Because of the law of conservation of mass: atoms are only rearranged, never created or destroyed, so the atoms entering the reaction must equal the atoms leaving it. That is also why the total mass of the products equals the total mass of the reactants. So if 10 g of A reacts completely with 15 g of B, exactly 25 g of products are formed. Remember the one thing you may never do while balancing: change a subscript inside a formula. Changing H₂O to H₂O₂ changes the substance itself, whereas a coefficient in front only changes how much of it you take.

Mole and limiting reagent (basic idea). A mole is the amount of substance that contains 6.022 × 10²³ particles (Avogadro's number). One mole of any gas occupies 22.4 litres at STP (0 °C and 1 atmosphere). The balanced equation gives the ratio in moles. In 2H₂ + O₂ → 2H₂O, two moles of hydrogen need exactly one mole of oxygen. Suppose we mix 5 moles of H₂ with 2 moles of O₂. Two moles of O₂ would need 4 moles of H₂, and we have 5, so hydrogen is in excess and oxygen finishes first. The reactant that is used up first and so decides how much product forms is the limiting reagent; here it is O₂, and the products are 4 moles of water. Stoichiometry is the quantitative study of the amounts of reactants and products in a reaction.

When we watch many reactions, a few patterns repeat again and again. Learning the patterns saves us from memorising hundreds of reactions. In a combination reaction two or more substances join to make a single product (A + B → AB). In a decomposition reaction one substance breaks into two or more simpler ones (AB → A + B), the opposite of combination. In a displacement reaction a more reactive element pushes a less reactive element out of its compound (A + BC → AC + B). In a double displacement reaction two compounds exchange their ions (AB + CD → AD + CB). Displacement reactions are always redox reactions, while double displacement reactions never are.

Decomposition needs energy. Every decomposition reaction takes in energy, which may be heat, light or electricity, and that decides its name. Heating limestone gives quicklime and carbon dioxide (thermal decomposition, used in making cement). Passing electricity through acidified water gives hydrogen at the cathode and oxygen at the anode in the volume ratio 2 : 1 (electrolytic decomposition). Silver chloride, which is white, turns grey in sunlight because silver is formed, and silver bromide behaves in the same way; this photolytic change is the basis of black-and-white photography. Heating lead nitrate gives yellow lead oxide, brown fumes of nitrogen dioxide and oxygen, with a crackling sound. Heating green ferrous sulphate crystals gives brown ferric oxide, with sulphur dioxide and sulphur trioxide, and the smell of burning sulphur.

Displacement follows the reactivity series. Place the metals in order of reactivity: K, Na, Ca, Mg, Al, Zn, Fe, Pb, (H), Cu, Hg, Ag, Au. A metal higher in the series displaces any metal below it from its salt solution. Iron nails dipped in blue copper sulphate solution become coated with a reddish-brown layer of copper, and the blue colour fades to pale green because iron sulphate forms. The nail is the more reactive metal, so it wins. The reverse never happens: copper cannot displace iron from iron sulphate. The thermite reaction is displacement too: aluminium displaces iron from iron oxide and gives so much heat that the iron is produced molten, which is used for welding railway tracks and cracked machine parts.

Double displacement gives a precipitate, a gas or water. These reactions happen in solution, where the ions simply swap partners. They go to completion when one product leaves the solution: as an insoluble solid (a precipitate), as a gas, or as water. Mixing sodium sulphate and barium chloride gives white insoluble barium sulphate. Mixing lead nitrate and potassium iodide gives a bright yellow precipitate of lead iodide (the 'golden rain' demonstration). A neutralisation reaction between an acid and a base is also a double displacement, and so is the reaction of an acid with a metal oxide, which gives a salt and water but no hydrogen. Another name for double displacement is metathesis. Silver nitrate with sodium chloride gives a white curdy precipitate of silver chloride. The reaction potassium bromide + barium iodide → potassium iodide + barium bromide (precipitate) is also of this type.

Disproportionation: one substance both oxidised and reduced. In some reactions the same element goes up in oxidation state in one part of the reaction and down in another. Chlorine in water is the standard example: Cl₂ + H₂O → HCl + HOCl. Here chlorine (0) becomes chloride (-1) in HCl and hypochlorite (+1) in HOCl. Hydrogen peroxide decomposing into water and oxygen is another one. This type is different from the four types above; it is a special kind of redox reaction.

Rusting of an iron gate, a cut apple turning brown, ghee or oil in an old packet smelling bad, a silver chain going black: all of these are oxidation at work. Oxidation and reduction always happen together, and a reaction in which both take place is a redox reaction. There are three ways to define the two terms, and exams use all three. In the oldest view, oxidation is the gain of oxygen (or loss of hydrogen) and reduction is the loss of oxygen (or gain of hydrogen). In the modern electron view, oxidation is the loss of electrons and reduction is the gain of electrons. In terms of oxidation number, oxidation is an increase and reduction a decrease. The substance that gets oxidised is the reducing agent (it gives electrons), and the one that gets reduced is the oxidising agent (it takes electrons).

Oxidation in daily life. When a metal is slowly eaten away by air, moisture or acids, the process is corrosion. The rusting of iron is the best-known example: iron reacts with oxygen and water to form reddish-brown hydrated iron(III) oxide, Fe₂O₃·xH₂O. Silver tarnishes and turns black because it reacts with hydrogen sulphide and sulphur compounds in the air to form silver sulphide, Ag₂S. Copper slowly gets a green coating of basic copper carbonate, as seen on old statues and domes. Aluminium also oxidises, but its thin oxide layer is tough and sticks firmly, so it protects the metal underneath. Rancidity is the oxidation of fats and oils in food: they turn bad in smell and taste. It is slowed by keeping food in airtight containers, in the refrigerator, by adding antioxidants (BHA and BHT) and by flushing packets of chips with nitrogen gas.

Electrochemical cells and electrolysis. A redox reaction can be made to push electrons through a wire, and this is how a battery works. In a galvanic (Daniell) cell with zinc and copper, zinc is oxidised at the anode and copper ions are reduced at the cathode, and electrons flow from the anode to the cathode through the external wire. In an electrolytic cell the opposite happens: an external supply forces a non-spontaneous redox reaction to occur. Electrolysis of brine (concentrated sodium chloride solution) gives sodium hydroxide, chlorine and hydrogen, the three products of the chlor-alkali industry. Faraday's first law says that the mass of a substance deposited at an electrode is directly proportional to the total electric charge passed (charge = current × time).

Every chemical reaction breaks old bonds and makes new ones, and it either gives out energy or takes it in. A reaction that releases heat to the surroundings is exothermic; the surroundings get warm. A reaction that absorbs heat from the surroundings is endothermic; the surroundings get cold. The heat exchanged at constant pressure is the enthalpy change (ΔH). For an exothermic reaction ΔH is negative, because the products have less energy than the reactants. For an endothermic reaction ΔH is positive. Burning of fuels, respiration, neutralisation and the slaking of lime are exothermic. Photosynthesis, thermal decomposition of limestone and dissolving of ammonium nitrate in water are endothermic.

Combustion. Combustion is a reaction in which a substance reacts rapidly with oxygen and gives out heat and usually light. The substance that burns is the fuel. A fuel starts to burn only when it is heated to its ignition temperature. Substances with a very low ignition temperature, such as petrol, alcohol and LPG, catch fire easily and are called inflammable. The products of the complete combustion of a carbon-containing fuel are carbon dioxide and water. When oxygen is limited, incomplete combustion produces carbon monoxide, which is poisonous, and carbon soot. This is why burning a coal angeethi in a closed room is dangerous. A gas stove burns with a blue flame because LPG (propane and butane) gets enough air through the holes of the burner and burns completely. If the air holes are blocked, the flame turns yellow and sooty because the burning is incomplete.

Putting out a fire. Fire needs three things: a fuel, oxygen (air) and heat up to the ignition temperature. Remove any one and the fire stops. Water cools the fuel below its ignition temperature, so it is good for wood and paper. Water should not be used on an oil or petrol fire, because oil floats on water and spreads the fire, and it must never be used on an electrical fire because water conducts electricity. For these fires a carbon dioxide extinguisher is used: CO₂ is heavier than air, so it covers the fire like a blanket and cuts off the oxygen supply. Carbon dioxide is also stored as a liquid under pressure in the cylinder and cools the fire as it expands.

A fuel needs air, and a good fuel needs a good burn. The more completely a fuel burns, the more heat it gives and the cleaner it is. That is why a clean blue flame is better than a smoky yellow one. For the same reason, it is dangerous to burn charcoal or coal in a closed room: the carbon monoxide produced binds to haemoglobin in blood about 200 times more strongly than oxygen does and chokes the body of oxygen.

Some reactions are over in a flash (an explosion), others take years (rusting of an iron beam). The speed of a reaction is its rate: how fast the reactants are used up or products are formed. For particles to react, they must collide, and the collision must be energetic enough. The minimum energy needed is the activation energy, a kind of hill that the reactants must climb before they roll down to become products. Any change that increases the number of effective collisions speeds up the reaction.

Catalysts. A catalyst is a substance that changes the rate of a chemical reaction without being used up. A positive catalyst speeds the reaction up by giving it an alternative route with a lower activation energy. A negative catalyst, called an inhibitor, slows the reaction down. Important features: a catalyst is not consumed (it is chemically unchanged at the end); a small amount is enough; it does not start a reaction that could not otherwise take place; it does not change the position of equilibrium or the equilibrium constant, but only helps the system reach equilibrium sooner; it lowers the activation energy of both the forward and the backward reactions equally; and it is often specific to a reaction. Finely powdered solid catalysts are more effective than lumps because they offer a larger surface area.

Why does a catalyst not change the equilibrium? Imagine a ball that can roll down either side of a hill. A catalyst lowers the hill, so the ball rolls over it faster in both directions. But the heights of the starting valley and the ending valley are unchanged, so the balance between the two is the same. In the same way a catalyst speeds up the forward and backward reactions equally, so the equilibrium constant does not change, only the time to reach it.

Enzymes, promoters and poisons. Enzymes are biological catalysts, almost all of them proteins. They work at body temperature, are highly specific (each enzyme suits one kind of substrate), and are damaged by high temperature or a wrong pH. Digestion depends on them: amylase in saliva breaks starch, pepsin in the stomach breaks proteins, lipase breaks fats. A promoter is a substance that boosts a catalyst's activity (molybdenum for iron in the Haber process). A catalyst poison destroys a catalyst's activity (arsenic poisons platinum in the Contact process, and lead in petrol poisons the catalytic converter, which is why unleaded petrol is used in modern cars). An autocatalyst is a product that catalyses its own reaction.

Reversible reactions and equilibrium. Many reactions can go both ways, and they are written with the double arrow ⇌. When the forward and backward rates become equal, the concentrations stop changing and the system is said to be in dynamic equilibrium: the reactions have not stopped, they just cancel out. Le Chatelier's principle says that if a system at equilibrium is disturbed, it shifts in the direction that opposes the disturbance. For the equilibrium CH₃COOH ⇌ CH₃COO⁻ + H⁺, adding H⁺ ions raises the concentration of a product, so the equilibrium shifts to the left, towards undissociated acetic acid. In the Haber process N₂ + 3H₂ ⇌ 2NH₃, the forward reaction is exothermic and reduces the number of gas molecules, so high pressure and a moderate temperature (around 450 °C, chosen to keep a good rate) with an iron catalyst are used. The Law of Mass Action was stated by Guldberg and Waage in 1864.