Chemical Reactions and Equations

Class 7 identified rusting and burning as chemical changes because new substances appear. Class 9 explained that every substance is built from atoms that bond together, often by sharing electrons, and proved that mass is conserved because no atom is ever created or destroyed. Put these ideas together and a chemical reaction is simply this: the atoms present in the starting substances break their old bonds and form new bonds, rearranging themselves into different combinations. Milk turning sour, a nail rusting in humid air, grapes fermenting, food cooking, food digesting, even the act of breathing, are all chemical reactions changing the identity of a starting substance into something new.

Clean a 3-4 cm strip of magnesium ribbon with sandpaper to remove its dull oxide coating, then hold it with tongs, well away from your eyes, and burn it over a flame, catching the ash in a watch-glass. The ribbon burns with a dazzling white flame and leaves behind a white powder utterly unlike the shiny grey ribbon you started with: magnesium oxide, formed as magnesium combines with oxygen in the air. A brand new substance appearing, with properties nothing like the original, is the clearest possible sign that a chemical reaction, not just a physical change, has taken place.

Magnesium's dazzling flame is not the only kind of evidence. Mix a colourless solution of lead nitrate with a colourless solution of potassium iodide, and instantly a bright yellow solid clouds the liquid: a sudden colour change, and a new solid appearing from two clear liquids, is exactly the kind of clue that reveals a reaction. Separately, drop a few zinc granules into dilute hydrochloric or sulfuric acid, and bubbles stream off the metal immediately; touch the container, and it feels noticeably warmer. Taken together, these activities show that any of four signs, a change in state, a change in colour, a gas being evolved, or a change in temperature, is good evidence that a chemical reaction, not just mixing, has taken place.

Describing the magnesium reaction in a full sentence every time gets tedious, so chemists shorten it. A word equation writes it as magnesium + oxygen leads to magnesium oxide, with the starting substances, called reactants, on the left and the new substance, the product, on the right, joined by an arrow showing the direction of change. Using chemical formulae instead of names makes it shorter still: Mg + O2 leads to MgO. Written this simply, though, it is only a skeletal equation, a first attempt that has not yet been checked to confirm the atom counts actually match on both sides.

Since atoms are only rearranged, never created or destroyed, a correct equation needs exactly the same number of atoms of each element on both sides. Draw a box around each formula (a reminder never to change what's inside a box while balancing), then count: Mg + O2 leads to MgO has 2 oxygen atoms on the left but only 1 on the right. Placing a coefficient, a number in front of a formula, fixes this without ever touching a formula itself: 2MgO gives 2 oxygen atoms to match, but now also 2 magnesium atoms, so 2Mg is needed on the left too. The balanced equation is 2Mg + O2 leads to 2MgO. A fully informative equation goes one step further, adding a state symbol after each formula to show what physical state it's in: (s) for solid, (l) for liquid, (g) for gas, (aq) for dissolved in water, as in 2Mg(s) + O2(g) leads to 2MgO(s). Sometimes the conditions needed for a reaction, like heat, a catalyst, or sunlight, are written above or below the arrow instead.

Add water slowly to calcium oxide (quicklime) in a beaker, and the mixture doesn't just get wet, it grows hot enough to feel through the glass almost immediately, as calcium oxide reacts vigorously with water to produce calcium hydroxide (slaked lime) and a large amount of heat: CaO(s) + H2O(l) leads to Ca(OH)2(aq) + heat. Two or more reactants forming a single product this way is called a combination reaction. Slaked lime is exactly what's used for whitewashing walls, and over the next two or three days, it slowly reacts with carbon dioxide in the air to form a thin, shiny layer of calcium carbonate, Ca(OH)2(aq) + CO2(g) leads to CaCO3(s) + H2O(l), the same compound, interestingly, that marble is made of. Other combination reactions include carbon burning in oxygen to form carbon dioxide, and hydrogen burning in oxygen to form water.

The heat given off when calcium oxide meets water is not a side detail, it has its own name. A reaction that releases heat as it forms its products is called an exothermic reaction. Burning natural gas is exothermic, and so, remarkably, is something happening inside you right now: respiration, where glucose from digested food combines with oxygen in your body's cells to release the energy that keeps you alive, C6H12O6(aq) + 6O2(aq) leads to 6CO2(aq) + 6H2O(l) + energy. Composting vegetable matter is exothermic too, which is why a compost heap can feel warm from the inside.

Heat pale green ferrous sulfate crystals in a dry boiling tube, and the green colour fades as the crystals lose water, then break down further, releasing the sharp smell of sulfur dioxide and sulfur trioxide gas and leaving reddish-brown ferric oxide behind: 2FeSO4(s) leads to Fe2O3(s) + SO2(g) + SO3(g) on heating. A single reactant breaking down into two or more simpler products like this is called a decomposition reaction, and one done using heat specifically is a thermal decomposition. The same thing happens when limestone (calcium carbonate) is heated to make quicklime for the cement industry, CaCO3(s) leads to CaO(s) + CO2(g), and when lead nitrate is heated in a boiling tube, releasing visible brown fumes of nitrogen dioxide gas.

Pass an electric current through water made slightly conductive with a little dilute sulfuric acid, using a simple two-electrode setup, and bubbles of gas collect at each electrode, filling two upturned test tubes, one with exactly twice the volume of gas of the other. Test each gas with a burning candle: the smaller volume (oxygen) makes the flame burn brighter, while the larger volume (hydrogen) itself burns with a pop. Electrolysis, decomposing water using electricity, splits it back into the same hydrogen and oxygen it was made from, in exactly a 2-to-1 ratio by volume. Sunlight, not just heat or electricity, can also drive a decomposition: white silver chloride left in sunlight slowly turns grey, decomposing into silver metal and chlorine gas, 2AgCl(s) leads to 2Ag(s) + Cl2(g), which is precisely the chemistry that once made black-and-white photography possible.

Decomposition reactions like these need an input of energy, heat for ferrous sulfate and limestone, electricity for water, light for silver chloride, to break the reactant's bonds apart. A reaction that absorbs energy like this, rather than releasing it, is called an endothermic reaction: essentially the opposite of the exothermic reactions from earlier in this chapter.

Dip a clean iron nail into blue copper sulfate solution and leave it for twenty minutes: the nail comes out coated brownish-red, and the solution's blue colour has visibly faded. Iron has displaced copper from copper sulfate solution, Fe(s) + CuSO4(aq) leads to FeSO4(aq) + Cu(s), forming pale green iron sulfate solution and depositing reddish copper metal onto the nail. This is called a displacement reaction, and it only works because iron is more reactive than copper; zinc and lead, being more reactive than copper too, displace it from its compounds in exactly the same way, Zn(s) + CuSO4(aq) leads to ZnSO4(aq) + Cu(s).

Mix a solution of sodium sulfate with a solution of barium chloride, and a white solid that won't dissolve appears at once: barium sulfate, Na2SO4(aq) + BaCl2(aq) leads to BaSO4(s) + 2NaCl(aq), left behind as a precipitate while sodium chloride stays dissolved in the liquid. Here, instead of one element displacing another, two compounds exchange partners entirely, sodium pairing with chloride and barium pairing with sulfate, which is why this is called a double displacement reaction, and any reaction producing an insoluble solid like this is specifically a precipitation reaction. Mixing sodium chloride solution with silver nitrate solution works the same way, instantly forming a white precipitate of silver chloride.

Heat copper powder in a china dish, and its shiny surface turns black as it reacts with oxygen from the air to form black copper oxide, 2Cu + O2 leads to 2CuO, on heating. Now pass hydrogen gas over this same heated black powder, and the coating turns back to brown copper metal, as the reverse happens: CuO + H2 leads to Cu + H2O, on heating. A substance that gains oxygen during a reaction is said to be oxidised, copper was oxidised the first time; a substance that loses oxygen is said to be reduced, copper oxide was reduced the second time, as hydrogen took its oxygen away and was itself oxidised in the process. Since one reactant is always oxidised while another is reduced in the very same reaction, this pairing is called an oxidation-reduction, or redox, reaction. Look back at the very first activity in this chapter: burning magnesium ribbon is magnesium gaining oxygen to form magnesium oxide, so magnesium itself was oxidised the whole time.

Oxidation is not confined to a lab. A shiny iron object left out in humid air slowly grows a reddish-brown coating, corrosion, the same process Class 7 called rusting; silver tarnishes black and copper grows a green coating for the same underlying reason, a metal being slowly attacked by moisture, acids or air. Corrosion of iron alone costs enormous amounts of money every year replacing damaged railings, car bodies, bridges and ships. Fats and oils oxidise too: left out for too long, they turn rancid, their smell and taste spoiling noticeably. That's exactly why antioxidants are added to packaged fatty foods, why airtight containers help slow the process down, and why bags of chips are often flushed with unreactive nitrogen gas instead of ordinary air before sealing.

Five classes of the same question, asked a little more precisely each time: Class 6 watched water quietly change state; Class 7 learned to tell a physical change from a chemical one, and met combustion and rusting; Class 8 explained both kinds of change in terms of particles; Class 9 learned to separate mixed particles, looked inside a single particle to find atoms and bonds, and proved mass is conserved; and Class 10 has now shown exactly how those atoms regroup during a chemical change, classified the patterns that regrouping falls into, and given a name, oxidation and reduction, to the give-and-take of oxygen and hydrogen underneath it all. What started as 'the puddle disappeared' has become 'atoms have rearranged into a new combination, one substance was oxidised while another was reduced, and here is the balanced equation that proves not one atom was lost.'