Atomic Foundations of Matter

Chapter 8 explained that atoms with fewer than 8 electrons in their outermost shell (or fewer than 2, if that shell is the very first one) are unstable, and become stable by sharing electrons with another atom, or by transferring electrons away or accepting them outright. But that raises a concrete, testable question: when atoms rearrange themselves this way, whether by simply dissolving into a solution or by properly reacting to form something new, does any matter actually get created or destroyed in the process, or is every single atom still fully accounted for at the end? And once two elements do combine, is it in any ratio at all, or always the same one?

Zero a digital weighing balance with an empty beaker, pour in 50 mL of water, and note the reading. Add a spatula of common salt and note the reading again: it increases by exactly the mass of the salt added, as expected. Now swirl the beaker until the salt completely dissolves, and check the balance one final time. The reading does not change at all: the mass of the salt solution exactly equals the mass of the water and salt that went into it, showing that dissolving, a physical change, does not create or destroy any mass, it only mixes it more thoroughly. The same holds if you weigh a sheet of paper, tear it into pieces, and weigh all the pieces together again.

Does the same hold for a genuine chemical change? Place vinegar in a flask and baking soda loosely in a balloon, both together on a weighing balance, and note the reading. Tip the baking soda into the vinegar: it fizzes vigorously, releasing carbon dioxide gas, exactly the same reaction that happens when baking soda meets any acid. Weigh the flask again once the fizzing stops, and the reading has dropped. At first, this looks like proof that a chemical change destroys mass. But look again at what changed about the experiment itself between the first weighing and the second: the flask was left open to the air throughout.

Repeat the exact same reaction, but this time tie the balloon's neck tightly over the mouth of the flask before tipping the baking soda in, so the whole system is sealed. The same vigorous fizzing happens, and this time the balloon visibly inflates as carbon dioxide gas is produced, unable to escape into the room. Weigh the sealed flask-and-balloon again: the final reading now exactly matches the initial one. The mass had not actually vanished the first time either; it had simply escaped into the air as gas before it could be weighed. This principle, that mass can neither be created nor destroyed in a chemical reaction, was proposed in 1789 by the French chemist Antoine Lavoisier, often called the father of modern chemistry, and is known as the Law of Conservation of Mass.

Weigh a flask of sodium sulfate solution and a flask of barium chloride solution separately, then pour one into the other. A white solid, barium sulfate, forms at once, but weigh the combined flasks again and the total is unchanged from the sum of the two starting weights. Even a reaction that produces a solid out of two clear liquids conserves mass exactly, as long as nothing is allowed to escape. This holds numerically too: if 4.0 g of calcium carbonate reacts completely with 2.92 g of hydrochloric acid in a closed container, and the products are measured afterward as 1.76 g of carbon dioxide, 0.72 g of water and 4.44 g of calcium chloride, the reactants' total (4.0 + 2.92 = 6.92 g) exactly equals the products' total (1.76 + 0.72 + 4.44 = 6.92 g).

Purify and analyse water from a river, a borewell, and the ocean, and every single sample breaks down into hydrogen and oxygen in exactly the same mass ratio: 1 part hydrogen to 8 parts oxygen, always, regardless of source. Soon after Lavoisier, the French chemist Joseph Proust proposed that this is true of every compound: the elements in it always combine in a fixed ratio by mass, no matter how or where the compound was formed. This is called the Law of Constant Proportions (or the Law of Definite Proportions, or Proust's Law). Sodium chloride, wherever it comes from, is always sodium and chlorine in a 23 to 35.5 mass ratio; so if 46 g of sodium reacts completely, it needs exactly (35.5 / 23) x 46 = 71 g of chlorine, never more or less.

Long before Proust gave it a name, ancient civilisations already understood this law in practice. Cinnabar, a brilliant red mineral known in India as hingula, was mined for its colour and, on heating, yields two elements: mercury and sulfur, in a fixed mass percentage of about 86.22% mercury to 13.78% sulfur, every single time. Many of these same civilisations also discovered the reverse: grinding mercury and sulfur together in this exact ratio could recreate cinnabar itself, though the toxicity of both elements kept this from ever becoming a widespread practice.

Both of these laws, mass never appearing or disappearing, and elements always combining in the same fixed ratio, needed an explanation, and John Dalton provided one in 1808 with a short list of postulates: all matter is made of tiny particles called atoms; atoms cannot be created, destroyed, or divided in a chemical reaction; atoms of a given element are identical to each other in mass and properties; atoms of different elements differ in mass and properties; atoms combine in ratios of small whole numbers to form compounds; and the relative number and kinds of atoms in a given compound are always constant. If atoms are truly indivisible and merely rearrange during a reaction, conservation of mass follows immediately, since nothing was ever created or destroyed, only reshuffled. And if a compound is always the same fixed combination of atoms, constant proportions follows too, since the same recipe of atoms always weighs out to the same ratio.

A hydrogen atom has only 1 electron in its single shell and needs 1 more to be stable; two hydrogen atoms can each share their single electron with the other, forming a shared pair that holds both nuclei together as a hydrogen molecule, H2, drawn as H-H. Chlorine, with 7 valence electrons, needs only 1 more, so two chlorine atoms share one electron each the same way to form Cl2 (Cl-Cl). This sharing of an electron pair between atoms is called a covalent bond, and a bond formed from just one shared pair is a single bond.

Oxygen has 6 valence electrons and needs 2 more, not just 1, to complete its octet. A single shared pair can't supply two electrons, so two oxygen atoms instead share two electrons each, forming two shared pairs between them at once: a double bond, drawn as O=O. This same idea (sharing one pair for a single bond, or two pairs for a double bond, depending on how many electrons are needed) is exactly how a hydrogen atom and a chlorine atom combine too: each needs only 1 electron, so they share a single pair to form hydrogen chloride, H-Cl, a covalent compound built from two different elements rather than two atoms of the same one.

Oxygen needs 2 electrons to complete its octet, but a single hydrogen atom only has 1 electron to offer. The two needs are matched instead by two separate hydrogen atoms, each sharing one electron with the same oxygen atom, giving oxygen its two extra electrons in total: a water molecule, H2O, held together by two separate single bonds, one to each hydrogen. This is exactly the reasoning behind naming any covalent compound: identify how many electrons each atom needs, then work out how many atoms of each must combine to supply exactly that many.

Covalent compounds are named using a prefix system that states exactly how many atoms of each element are present: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), and so on, with the first element keeping its ordinary name and the second element's name changed to end in -ide. Mono- is normally dropped for the first element, but kept for the second: CO is carbon monoxide, not monocarbon monoxide, while CO2 is carbon dioxide. CS2 is carbon disulfide (two sulfurs), PCl3 is phosphorus trichloride (three chlorines), SF6 is sulfur hexafluoride (six fluorines), and N2O4 is dinitrogen tetroxide. When hydrogen is the first element, no prefix is used at all, however many atoms are present: H2S is simply hydrogen sulfide, never dihydrogen sulfide. A handful of very familiar covalent compounds are known only by their common names instead: H2O, which the prefix system would call hydrogen monoxide, is universally just water, and NH3, technically nitrogen trihydride, is universally ammonia.

Metals like sodium behave completely differently from hydrogen or chlorine. Sodium's valence shell holds only 1 electron, and rather than sharing it, sodium simply loses it entirely to another atom, becoming a positively charged sodium ion (a cation), Na+, with 11 protons but only 10 electrons left. Chlorine, needing just 1 more electron to complete its own octet, accepts that same electron and becomes a negatively charged chloride ion (an anion), Cl-, with 17 protons and 18 electrons. Once formed, these oppositely charged ions attract each other through simple electrostatic force, exactly as opposite charges always do, and that attraction itself is called an ionic bond. Cations and anions together are simply called ions; metals typically form cations by losing electrons, and non-metals typically form anions by gaining them.

A single Na+ ion and a single Cl- ion don't pair off into an isolated 'molecule' the way covalent atoms do. Instead, countless sodium and chloride ions arrange themselves into a repeating three-dimensional pattern called a crystal, where every single Na+ ion is surrounded by six Cl- ions, and every Cl- ion is surrounded by six Na+ ions in turn, held together throughout by the same electrostatic attraction. This regular, repeating arrangement is called a crystal lattice, and it's the reason ionic compounds are described using a formula unit (the simplest whole-number ratio of ions, like NaCl) rather than a molecule.

An ionic compound is named cation first, then anion, with simple anions ending in -ide: sodium chloride, calcium oxide, magnesium sulfide. Not every ion is a single atom, though. A polyatomic ion is a small group of atoms bonded together that carries one single overall charge, and its exact composition is a fixed fact to know, not something worked out by valency, the same way an element's symbol is a fixed fact you look up on the periodic table. Hydroxide (OH-) is always exactly one oxygen and one hydrogen, together carrying a 1- charge; sulfate (SO4 2-) is always one sulfur and four oxygens, together carrying a 2- charge; carbonate is CO3 2-; nitrate is NO3-; ammonium is NH4+, one of the few positive polyatomic ions. Common monatomic ions have their own fixed charges too: Group-1 metals (sodium, lithium, potassium, silver) are always 1+; calcium, barium, magnesium, zinc are always 2+; aluminium is always 3+; the halogens (fluoride, chloride, bromide, iodide) are always 1-; oxide and sulfide are always 2-. A few metals, like iron and copper, can form more than one ion (ferrous Fe2+ or ferric Fe3+; cuprous Cu+ or cupric Cu2+), so both possibilities need to be checked.

Covalent formulas can be written quickly using each element's valency, in almost the same way ionic formulas use ion charges. Write the symbols of the elements, write each one's valency underneath, then criss-cross those valency numbers across as subscripts (dropping a subscript of 1). Hydrogen (valency 1) and chlorine (valency 1) criss-cross to plain HCl. Hydrogen (valency 1) and sulfur (valency 2) criss-cross to H2S, hydrogen sulfide. Carbon (valency 4) and chlorine (valency 1) criss-cross to CCl4, carbon tetrachloride.

Ionic formulas follow the same crossing-over idea, but using each ion's charge: write the cation first, then the anion, write their charge numbers underneath (ignoring the plus or minus sign), then criss-cross those numbers across as the other ion's subscript, simplifying afterward if both numbers share a common factor. Calcium (2+) and oxide (2-) criss-cross to Ca2O2, which simplifies down to plain CaO, since 2 and 2 share a common factor. But this simplifying step doesn't always happen: aluminium (3+) and oxide (2-) criss-cross to Al2O3, and since 2 and 3 share no common factor, that stays exactly as it is, the real formula, not a simplified one. Aluminium hydroxide, similarly, is Al(OH)3 (brackets are needed whenever more than one polyatomic ion is present), and aluminium sulfate, criss-crossing 3 and 2, is Al2(SO4)3.

The same criss-cross method explains acids too, since an acid is really just hydrogen, valency 1, combined with a negative ion. Hydrogen (H+, charge 1) and sulfate (SO4 2-, charge 2) criss-cross exactly the way calcium and oxide did: 2 hydrogens are needed to balance one sulfate's 2- charge, giving H2SO4, sulfuric acid, and there is no common factor to simplify away, since 2 and 1 share none. Hydrogen and nitrate (NO3-, charge 1) criss-cross to plain HNO3, nitric acid, since both charges are already 1. The '4' inside SO4 was never derived by any of this criss-crossing at all, it was simply sulfate's own fixed composition, one sulfur bonded to four oxygens, learned as a fact the same way hydroxide's 'OH' or carbonate's 'CO3' are learned; only the '2' in front of the H came from balancing charges.

Try dissolving camphor, common salt, copper sulfate, sugar and naphthalene, each in water, then in kerosene, then in petrol. The ionic compounds, common salt and copper sulfate, dissolve readily in water but not in kerosene or petrol; the covalent compounds, camphor and naphthalene, do the reverse, dissolving in kerosene and petrol but not water (sugar, also covalent, is an exception that dissolves in water too, without conducting electricity). Now test each solid, and then each water solution, for electrical conductivity using two electrodes wired to a bulb and a 9-volt battery. The solids themselves never light the bulb, ionic or covalent, since their ions or molecules aren't free to move. Dissolved in water, though, common salt and copper sulfate light the bulb brightly, since their ions are now free to carry charge through the solution, while sugar's water solution stays dark, since dissolved sugar produces no ions at all. Ionic compounds also generally have far higher melting and boiling points than covalent compounds, since breaking apart a whole crystal lattice of electrostatic attractions takes much more energy than separating individual covalent molecules from each other.

Once a covalent compound's formula is known, its molecular mass, the mass of one molecule, is simply the sum of the atomic masses of every atom in it: water, H2O, is (1 x 2) + (16 x 1) = 18 atomic mass units, and carbon dioxide, CO2, is (12 x 1) + (16 x 2) = 44 units. Ionic compounds don't form true molecules, so the equivalent idea is called a formula unit mass instead, the mass of the simplest whole-number ratio of ions the formula represents: sodium oxide, Na2O, is (23 x 2) + (16 x 1) = 62 units, and calcium nitrate, Ca(NO3)2, is 40 + [(14 + 16x3) x 2] = 164 units, remembering that everything inside the bracket is doubled by the subscript outside it.

Every idea in this chapter, conservation of mass, constant proportions, covalent and ionic bonding, the criss-cross method, molecular and formula unit mass, has been about compounds that already exist, quietly sitting there to be weighed and analysed. The next question is what happens the moment substances actually react with one another, not just combine once into a stable compound, but genuinely transform, the way rusting or burning did back in Class 7. Class 10 takes this further, showing precisely how atoms in reacting substances regroup into new combinations, and how to represent that regrouping exactly, atom for atom, using a balanced chemical equation.