Chapter 1: The Language of Chemistry
In a Paris laboratory in the 1770s, Antoine Lavoisier did something that had rarely been done before. He sealed a substance inside a glass vessel with air, weighed the whole vessel on a fine balance, heated it, and weighed it again. The metal inside grew heavier. The vessel, sealed, weighed exactly what it had weighed before. Nothing had been created, and nothing had been lost — matter had only changed hands inside the glass.
That result was hard to describe, because the chemistry of the time had no proper language. The same white powder might be sold as flowers of zinc in one shop and as philosophical wool in another. Names told you where a substance came from, or what it looked like, but never what it was made of.
So in 1787 Lavoisier and three colleagues — Guyton de Morveau, Berthollet and Fourcroy — published a new method of naming. A substance would now be named after the elements inside it. Oil of vitriol became sulphuric acid. The name itself began to carry information.
You are about to learn the language that grew out of that decision: symbols for elements, formulae for compounds, and equations for changes. It is a language in which every letter, every small number and every sign means something exact. Fig. 1.0-1 shows where it began.
Diagram
Fig. 1.0-1 — Naming What A Substance Is Made Of Left panel: an eighteenth-century beam balance carrying a sealed round glass vessel with a small dish of metal inside it, a burner below; label the sealed vessel, the trapped air, the metal, and the balance pointer reading unchanged. Right panel: a simple two-column card headed ‘old name’ and ‘name based on composition’, with three paired entries — oil of vitriol / sulphuric acid, flowers of zinc / zinc oxide, spirit of salt / hydrochloric acid.
1.1 Matter, Elements And Compounds
Chemistry is the branch of science that studies matter. It asks what matter is made of, how it is built, and how it changes when we heat it, cool it, dissolve it or mix it with something else.
To answer those questions, chemists had to look below what the eye can see. Everything around you — the air in the room, the water in a glass, the metal of a spoon — is built from particles far too small to be seen. This section introduces the three names you will need for those particles and for the substances they build.
1.1.1 Molecules And Pure Substances
The basic unit of matter, in any state, is taken to be the molecule. A molecule is the smallest particle of a substance that can exist on its own and still be that substance. Break a molecule of water apart and you no longer have water.
A pure substance is one in which every molecule is alike. Because all its molecules are the same, a pure substance has a definite composition and a definite set of properties — a fixed melting point, a fixed boiling point, a fixed density.
A pure substance is also homogeneous, which means it is the same all the way through. Carbon dioxide gas is a pure substance: every single molecule in it is built from one carbon atom and two oxygen atoms, whether you sample the top of the jar or the bottom.
Elements and compounds are both pure substances. The difference between them lies in the kind of atoms inside their molecules, and that is where we turn next.
1.1.2 Atoms And Atomicity
A molecule may hold atoms of one kind or of different kinds. A pure substance whose molecules are built from only one kind of atom is called an element.
At present 118 elements are known to us. They are shared out very unevenly across the universe: hydrogen and helium together make up the great majority of its known mass, and every other element put together accounts for only about one per cent.
Every element is made of atoms. An atom is the smallest particle of an element that still shows the properties of that element, and it is the particle that takes part in a chemical reaction.
Notice the careful wording — may or may not exist independently. A helium atom is happy on its own. An oxygen atom is not; it is found joined to a second oxygen atom. So an atom is not always able to exist alone, while a molecule always is.
Counting The Atoms In A Molecule
When atoms of the same element join together, they form a molecule of that element. The number of atoms in one molecule of an element is called its atomicity. Table 1.1 lists the atomicity of some familiar elements, and Fig. 1.1-1 shows what those molecules look like.
Table 1.1 : Atomicity of some elements
| Atomicity | Name of the type | Examples |
|---|---|---|
| One | Monoatomic | Noble gases — helium, neon, argon, krypton, xenon and radon; metals; carbon |
| Two | Diatomic | Hydrogen \(\\mathrm{H\_2}\), oxygen \(\\mathrm{O\_2}\), nitrogen \(\\mathrm{N\_2}\), chlorine \(\\mathrm{Cl\_2}\) |
| Three | Triatomic | Ozone \(\\mathrm{O\_3}\) |
| Four | Tetratomic | Phosphorus \(\\mathrm{P\_4}\) |
| Eight | Octatomic | Sulphur \(\\mathrm{S\_8}\) |
Diagram
Fig. 1.1-1 — Molecules Of Elements With Different Atomicity Five simple ball models side by side, each labelled with its name, formula and atomicity: a single sphere for helium (He, monoatomic); two joined spheres for oxygen (\(\\mathrm{O\_2}\), diatomic); three spheres in a bent chain for ozone (\(\\mathrm{O\_3}\), triatomic); four spheres arranged as a tetrahedron for phosphorus (\(\\mathrm{P\_4}\), tetratomic); eight spheres in a puckered ring for sulphur (\(\\mathrm{S\_8}\), octatomic).
1.1.3 Compounds And Mixtures
When atoms of different elements combine chemically in a fixed ratio, they form the molecule of a compound. Sodium chloride, water and ammonia are all compounds, written \(\\mathrm{NaCl}\), \(\\mathrm{H\_2O}\) and \(\\mathrm{NH\_3}\).
The word chemically is doing real work in that sentence. The atoms in a compound are held by chemical bonds, so a compound cannot be broken down by any physical means such as filtering, sieving or letting it settle.
A Compound Is Not Its Elements
The properties of a compound are entirely different from those of the elements inside it. Sodium is a highly reactive, poisonous metal. Chlorine is a greenish-yellow poisonous gas. The compound they form together, sodium chloride, is neither poisonous nor reactive — it is the salt in your food.
This is worth pausing on. Once atoms combine chemically, they are no longer free to behave as they did before, so the compound starts with a fresh set of properties of its own.
When Substances Are Only Mixed
Two or more kinds of molecules can also be brought together without any chemical combination at all. Such a combination is a mixture. A mixture may contain only elements, or an element and a compound, or only compounds.
A mixture is called homogeneous when its composition is uniform throughout, and heterogeneous when the different molecules are not spread evenly. Salt solution is homogeneous; muddy water, with its silt settling at the bottom, is heterogeneous.
Mixtures are impure substances. They are formed by mixing two or more pure substances in any proportion, no chemical change takes place, and each substance keeps its own properties. Table 1.2 sets the three side by side.
1.2 Symbols Of Elements
The names of elements and compounds are shortened by using symbols and formulae. A chemist who writes \(\\mathrm{Fe}\) instead of ‘iron’ is not being lazy — the short form is exact, and it means the same thing in every country.
1.2.1 From Pictographs To A Universal System
Before 1600 A.D., alchemists represented the substances they worked with by pictures. A triangle stood for earth and a crescent for silver. These pictographs were private, and different workers used different ones.
John Dalton later used a set of small circular signs to represent elements — a plain circle for an oxygen atom, a circle with a dot at its centre for hydrogen, and so on. Fig. 1.2-1 shows the signs he used.
Diagram
Fig. 1.2-1 — Dalton’s Symbols Of Elements A two-column chart. Left column, the element names in order: hydrogen, carbon, oxygen, phosphorus, sulphur, platinum, iron, copper, silver, gold, lead, mercury. Right column, the matching sign drawn as a circle of the same size in each row: hydrogen a circle with a central dot; carbon a solid filled circle; oxygen an empty circle; phosphorus a circle containing a slanted stroke; sulphur a circle containing a cross; platinum a circle enclosing the letter P; iron a circle enclosing the letter I; copper a circle enclosing the letter C; silver a circle enclosing the letter S; gold a circle enclosing the letter G; lead a circle enclosing the letter L; mercury a circle enclosing a pattern of small dots.
Drawing a picture for every element could not last. As more elements were discovered, the signs became hard to remember and harder still to print. What was needed was a system built from letters.
Today that work is looked after by IUPAC, the International Union of Pure and Applied Chemistry. IUPAC guides research and naming in chemistry, and the symbols it approves are accepted everywhere, so each element has one unique symbol used by everyone in the world.
Why Symbols And Not Names
Element names depend on language, particularly in their endings. Magnesium becomes magnésium in French, magnesio in Spanish and magnij in Russian; Japanese reproduces the sound of the English word in katakana script.
The symbol does not change. Whatever a chemist calls the element aloud, \(\\mathrm{Mg}\) is written the same way in Jaipur, Paris and Tokyo. Symbols make the language of chemistry independent of cultural, linguistic and geographical boundaries.
1.2.2 How A Symbol Is Written
A symbol is the short form that stands for the atom of a particular element.
Each element is denoted by a symbol taken usually from the first letter of its name in English or in Latin, written as a capital. Sulphur is \(\\mathrm{S}\) and hydrogen is \(\\mathrm{H}\).
Trouble begins when several elements share a first letter. Carbon, cobalt, calcium, cadmium, chromium and chlorine all begin with C, so only one of them can have the single letter. Carbon takes \(\\mathrm{C}\); the others take two letters — the first written as a capital, followed by some prominent letter written small: \(\\mathrm{Co}\), \(\\mathrm{Ca}\), \(\\mathrm{Cd}\), \(\\mathrm{Cr}\) and \(\\mathrm{Cl}\).
The second letter is the reason capitals matter so much. \(\\mathrm{Co}\) is the element cobalt, while \(\\mathrm{CO}\) — two capitals — is the compound carbon monoxide, made of one carbon atom and one oxygen atom. A careless capital letter changes an element into a compound.
Symbols That Do Not Match The English Name
Some symbols do not agree with the English name at all, because they were taken from the older Latin or German name of the element. Tungsten is written \(\\mathrm{W}\), from its German name Wolfram, and copper is written \(\\mathrm{Cu}\), from its Latin name Cuprum. Table 1.3 lists the elements you will meet whose symbols come from these older names.
Table 1.3 : Symbols of certain elements based on their Latin names
| Name of element | Symbol | Latin name | |
|---|---|---|---|
| 1. | Gold | \(\\mathrm{Au}\) | Aurum |
| 2. | Silver | \(\\mathrm{Ag}\) | Argentum |
| 3. | Mercury | \(\\mathrm{Hg}\) | Hydrargyrum |
| 4. | Copper | \(\\mathrm{Cu}\) | Cuprum |
| 5. | Lead | \(\\mathrm{Pb}\) | Plumbum |
| 6. | Iron | \(\\mathrm{Fe}\) | Ferrum |
| 7. | Sodium | \(\\mathrm{Na}\) | Natrium |
| 8. | Potassium | \(\\mathrm{K}\) | Kalium |
| 9. | Tin | \(\\mathrm{Sn}\) | Stannum |
| 10. | Antimony | \(\\mathrm{Sb}\) | Stibium |
| 11. | Tungsten | \(\\mathrm{W}\) | Wolfram (German name) |
Table 1.4 gives the symbols of the common elements you will use throughout this book. It is worth learning them early, because every formula and every equation that follows is built out of them.
Table 1.4 : Symbols of some common elements
| Name in English | Symbol | Name in English | Symbol |
|---|---|---|---|
| Magnesium | \(\\mathrm{Mg}\) | Hydrogen | \(\\mathrm{H}\) |
| Aluminium | \(\\mathrm{Al}\) | Nitrogen | \(\\mathrm{N}\) |
| Calcium | \(\\mathrm{Ca}\) | Oxygen | \(\\mathrm{O}\) |
| Chromium | \(\\mathrm{Cr}\) | Fluorine | \(\\mathrm{F}\) |
| Manganese | \(\\mathrm{Mn}\) | Chlorine | \(\\mathrm{Cl}\) |
| Cobalt | \(\\mathrm{Co}\) | Bromine | \(\\mathrm{Br}\) |
| Nickel | \(\\mathrm{Ni}\) | Iodine | \(\\mathrm{I}\) |
| Zinc | \(\\mathrm{Zn}\) | Carbon | \(\\mathrm{C}\) |
| Silver | \(\\mathrm{Ag}\) | Sulphur | \(\\mathrm{S}\) |
| Barium | \(\\mathrm{Ba}\) | Phosphorus | \(\\mathrm{P}\) |
| Arsenic | \(\\mathrm{As}\) | Boron | \(\\mathrm{B}\) |
| Platinum | \(\\mathrm{Pt}\) | Silicon | \(\\mathrm{Si}\) |
| Uranium | \(\\mathrm{U}\) | Helium | \(\\mathrm{He}\) |
| Lithium | \(\\mathrm{Li}\) | Neon | \(\\mathrm{Ne}\) |
| Germanium | \(\\mathrm{Ge}\) | Argon | \(\\mathrm{Ar}\) |
1.2.3 Significance Of A Symbol
A symbol is not merely a short way of writing a name. It carries three pieces of information at once.
It represents: (i) the name of the element; (ii) one atom of that element; (iii) a definite mass of the element, equal to its atomic mass expressed in grams.
Take the symbol \(\\mathrm{N}\). It stands for the element nitrogen; it stands for one atom of nitrogen; and it stands for 14 grams of nitrogen, 14 being the atomic mass of the element.
The third meaning is the one students overlook, and it is the one that makes calculation possible. Because a symbol carries a mass, an equation written in symbols can later be read as a statement about weights in the laboratory.
1.3 Valency
Atoms do not combine in any proportion they please. One atom of sodium takes exactly one atom of chlorine; one atom of carbon takes exactly four atoms of hydrogen. Every element behaves as though it has a fixed number of hands with which to hold other atoms.
That number is called its valency. The valency of an element is a measure of its combining power with other elements — the property that decides how many other atoms one atom of it can combine with.
1.3.1 Valency In Terms Of Hydrogen And Oxygen
To compare combining powers, chemists needed one element to measure against. No element has a combining capacity smaller than that of hydrogen, so hydrogen was given a valency of one and made the standard.
Valency is therefore measured in terms of hydrogen atoms, or of oxygen atoms. It is the number of hydrogen atoms, or double the number of oxygen atoms, that can combine with one atom of the element.
Oxygen is doubled because one oxygen atom itself holds two hydrogen atoms, as water shows. So an oxygen atom is worth two hydrogen atoms in this counting.
Counting With Hydrogen
Look at what each element does when it combines with hydrogen. The three molecules are drawn in Fig. 1.3-1.
One atom of chlorine combines with one hydrogen atom to form a molecule of hydrogen chloride. So the valency of chlorine is one.
In an ammonia molecule, one atom of nitrogen combines with three atoms of hydrogen. So the valency of nitrogen is three.
In a methane molecule, one carbon atom combines with four hydrogen atoms. So the valency of carbon is four.
Diagram
Fig. 1.3-1 — Counting Valency From Hydrogen Compounds Three simple structural formulae drawn side by side, each with the compound name and the valency written below it. First: H–Cl, labelled hydrogen chloride, valency of chlorine = 1. Second: a central N with three single lines going down and outward to three H atoms, labelled ammonia, valency of nitrogen = 3. Third: a central C with four single lines to four H atoms placed above, below, left and right, labelled methane, valency of carbon = 4.
Counting With Oxygen
In sodium oxide, \(\\mathrm{Na\_2O}\), one oxygen atom combines with two sodium atoms. Each sodium atom therefore holds only half an oxygen atom’s worth of combining power: \[\\text{valency of Na} = 2 \\times \\frac{1}{2} = 1\]
In carbon dioxide, \(\\mathrm{CO\_2}\), one carbon atom combines with two oxygen atoms: \[\\text{valency of C} = 2 \\times 2 = 4\]
Both routes give the same answer for carbon, which is exactly what we should expect. Valency belongs to the element itself, not to the partner we happen to measure it against.
1.3.2 Modern Definition Of Valency
The hydrogen counting works, but it does not say why an element has the combining power it has. The reason lies inside the atom, among its electrons.
The number of electrons that an atom can lose, gain or share during a chemical reaction is called its valency.
An atom contains electrons, protons and neutrons. Each electron carries one unit of negative charge and each proton carries one unit of positive charge, while the neutron carries no charge at all.
Every atom normally contains equal numbers of electrons and protons. The two charges balance exactly, and so the atom as a whole is electrically neutral.
1.3.3 Ions, Cations And Anions
If one or more electrons are removed from an atom, or added to it, that balance is broken. The atom then has too few or too many electrons, and it carries an overall electric charge. It has become an ion.
Which electrons are lost or gained? Not any electron in the atom, but only those in the outermost shell, which is called the valence shell. Electrons in the outermost shell are called valence electrons, and they are the ones that take part in chemical change.
Atoms lose or gain these electrons for a reason. To become stable, an atom needs eight electrons in its outermost shell — an octet. Hydrogen is the exception; it becomes stable with two.
Atoms That Lose Electrons
An atom with one, two or three electrons in its valence shell finds it easier to give them away than to collect five, six or seven more. Losing them leaves the full shell below as the new outermost shell.
\[\\mathrm{Na}\\ (2,8,1) - e^{-} \\rightarrow \\mathrm{Na^{+}}\\ (2,8) \\quad \\text{monovalent}\] \[\\mathrm{Mg}\\ (2,8,2) - 2e^{-} \\rightarrow \\mathrm{Mg^{2+}}\\ (2,8) \\quad \\text{divalent}\] \[\\mathrm{Al}\\ (2,8,3) - 3e^{-} \\rightarrow \\mathrm{Al^{3+}}\\ (2,8) \\quad \\text{trivalent}\]
Positive ions are known as cations. Notice that the number of electrons lost is exactly the valency of the element.
Elements with one, two or three electrons in the valence shell are usually metals. Among the elements with four electrons in the outermost shell, carbon is a non-metal, silicon and germanium are metalloids, and the rest are metals. Elements with five, six or seven electrons in their outermost shell are normally non-metals.
Atoms That Gain Electrons
An atom with seven, six or five valence electrons has the shorter journey in the other direction. It gains one, two or three electrons to complete its octet.
\[\\mathrm{Cl}\\ (2,8,7) + e^{-} \\rightarrow \\mathrm{Cl^{-}}\\ (2,8,8) \\quad \\text{monovalent}\] \[\\mathrm{O}\\ (2,6) + 2e^{-} \\rightarrow \\mathrm{O^{2-}}\\ (2,8) \\quad \\text{divalent}\] \[\\mathrm{N}\\ (2,5) + 3e^{-} \\rightarrow \\mathrm{N^{3-}}\\ (2,8) \\quad \\text{trivalent}\]
Negative ions are known as anions. Fig. 1.3-2 shows one atom of each kind making the change.
Diagram
Fig. 1.3-2 — Formation Of A Cation And Of An Anion Two shell diagrams side by side, each drawn as a nucleus with concentric circles carrying dots for electrons. Left: a sodium atom labelled 2,8,1 with an arrow leaving the outermost shell carrying one electron, giving a smaller sodium ion labelled 2,8 and marked \(\\mathrm{Na^{+}}\); caption line ‘loss of one electron — cation’. Right: a chlorine atom labelled 2,8,7 with an arrow bringing one electron into the outermost shell, giving a chloride ion labelled 2,8,8 and marked \(\\mathrm{Cl^{-}}\); caption line ‘gain of one electron — anion’.
Table 1.5 collects the pattern. Read it from both ends and the reason becomes clear: an atom always takes the shorter route to a complete octet, so the valency first rises to four and then falls again.
Table 1.5 : Valency and the number of electrons in the outermost shell
| Number of electron(s) in outermost shell | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 |
|---|---|---|---|---|---|---|---|---|
| Valency | 1 | 2 | 3 | 4 | 3 | 2 | 1 | 0 |
An atom with eight electrons already has its octet. It has nothing to gain and nothing to give, so its valency is zero — which is why the noble gases you met in Table 1.1 are found as single, uncombined atoms.
1.3.4 Valency Of Elements
Depending on their valency, elements may be described as monovalent, divalent, trivalent or tetravalent.
Monovalent elements have a valency of one. They lose or gain one electron, and form one bond. Hydrogen, lithium, sodium, potassium, chlorine and fluorine are monovalent.
Divalent elements have a valency of two and can lose or gain two electrons. Magnesium, calcium, iron as \(\\mathrm{Fe^{2+}}\) and oxygen are divalent.
Trivalent elements have a valency of three. Boron, aluminium, nitrogen and phosphorus are common examples.
Tetravalent elements have a valency of four. They hold four electrons in their outermost shell and can form four covalent bonds. Carbon, silicon, germanium, tin and lead belong here.
1.3.5 Variable Valency
Certain elements do not stop at one valency. They exhibit more than one, and are said to show variable valency.
Iron is the everyday example. It shows valencies of 2 and 3, and so forms two different chlorides — \(\\mathrm{FeCl\_2}\), iron(II) chloride, and \(\\mathrm{FeCl\_3}\), iron(III) chloride. These are different compounds, with different colours and different reactions.
Why An Element Can Have Two Valencies
An atom of such an element can sometimes lose more electrons than are present in its valence shell. Electrons are then drawn from the penultimate shell as well — the shell next to the outermost one.
How many are removed depends on the conditions of the reaction. An element that can be stripped in two different ways ends up with two different valencies, and that is the origin of variable valency.
Naming The Two Valencies
Where an element shows two different positive valencies, the older names used the suffix ous for the lower valency and the suffix ic for the higher one. Copper(I) compounds were cuprous, copper(II) compounds cupric.
Modern chemists write a Roman numeral in place of these names. Thus \(\\mathrm{SnCl\_2}\), stannous chloride, is written tin(II) chloride, and \(\\mathrm{SnCl\_4}\), stannic chloride, is written tin(IV) chloride. This method is known as stock notation, in which a Roman numeral in brackets states the valency.
The Roman numeral is a clear improvement, because ous and ic only tell you which of two valencies is the lower. They cannot help at all with an element such as sulphur, which has three. Table 1.6 gives the elements you will meet most often.
Table 1.6 : Examples of variable valency
| Element | Valency | Name of compound formed | Formula |
|---|---|---|---|
| Iron | 2 | Ferrous [iron(II)] oxide | \(\\mathrm{FeO}\) |
| 3 | Ferric [iron(III)] oxide | \(\\mathrm{Fe\_2O\_3}\) | |
| Copper | 1 | Cuprous [copper(I)] oxide | \(\\mathrm{Cu\_2O}\) |
| 2 | Cupric [copper(II)] oxide | \(\\mathrm{CuO}\) | |
| Mercury | 1 | Mercurous [mercury(I)] oxide | \(\\mathrm{Hg\_2O}\) |
| 2 | Mercuric [mercury(II)] oxide | \(\\mathrm{HgO}\) | |
| Lead | 2 | Plumbous [lead(II)] oxide | \(\\mathrm{PbO}\) |
| 4 | Plumbic [lead(IV)] oxide | \(\\mathrm{PbO\_2}\) | |
| Tin | 2 | Stannous [tin(II)] chloride | \(\\mathrm{SnCl\_2}\) |
| 4 | Stannic [tin(IV)] chloride | \(\\mathrm{SnCl\_4}\) | |
| Nitrogen | 3 | Ammonia | \(\\mathrm{NH\_3}\) |
| 5 | Dinitrogen pentoxide | \(\\mathrm{N\_2O\_5}\) | |
| Phosphorus | 3 | Phosphorus trichloride | \(\\mathrm{PCl\_3}\) |
| 5 | Phosphorus pentachloride | \(\\mathrm{PCl\_5}\) | |
| Sulphur | 2 | Hydrogen sulphide | \(\\mathrm{H\_2S}\) |
| 4 | Sulphur dioxide | \(\\mathrm{SO\_2}\) | |
| 6 | Sulphur trioxide | \(\\mathrm{SO\_3}\) |
Variable valency is not a metallic privilege. Non-metals show it too: nitrogen and phosphorus exhibit valencies of 3 and 5, while sulphur exhibits valencies of 2, 4 and 6.
1.4 Formula
Atoms of elements combine to form molecules. Since every atom now has a symbol, a whole molecule can be written down in symbols too. This symbolic representation of a molecule is known as a formula, or molecular formula.
A molecular formula, also called a chemical formula, uses symbols to denote the molecule of an element or of a compound.
1.4.1 Formulae Of Elements And Compounds
A molecule of an element may contain one atom or several. The molecules of hydrogen, oxygen, nitrogen, chlorine, bromine and iodine each contain two atoms, and are written \(\\mathrm{H\_2}\), \(\\mathrm{O\_2}\), \(\\mathrm{N\_2}\), \(\\mathrm{Cl\_2}\), \(\\mathrm{Br\_2}\) and \(\\mathrm{I\_2}\).
In a compound, different atoms are united in a certain fixed ratio. The formula is written by placing the symbols of the elements present side by side, with their numbers written as subscripts.
Read the subscript as belonging only to the symbol just before it. So \(\\mathrm{NH\_4Cl}\) represents one molecule of ammonium chloride, containing one atom of nitrogen, four atoms of hydrogen and one atom of chlorine.
In the same way, \(\\mathrm{Na\_2CO\_3}\) denotes one molecule of sodium carbonate, which contains two atoms of sodium, one atom of carbon and three atoms of oxygen. Fig. 1.4-1 shows how this formula is read.
Diagram
Fig. 1.4-1 — Reading A Chemical Formula The formula Na₂CO₃ printed large across the centre. Three arrows point up from below to the three symbols in turn, labelled ‘sodium — 2 atoms’, ‘carbon — 1 atom’, ‘oxygen — 3 atoms’. A fourth arrow points down from above to the subscript 3, labelled ‘a subscript counts the atoms of the symbol it follows’. A note at the foot reads ‘one molecule of sodium carbonate’.
A Number In Front Is A Different Number
A subscript sits low and to the right, and belongs to one symbol. A number written in front of the whole formula is a different thing altogether — it counts complete molecules.
Thus \(2\\mathrm{H\_2O}\) represents two molecules of water, each molecule containing two atoms of hydrogen and one atom of oxygen. Altogether that is four hydrogen atoms and two oxygen atoms.
Brackets work in the same spirit. In \(\\mathrm{Ca(OH)\_2}\) the subscript 2 applies to everything inside the bracket, so the molecule holds two oxygen atoms and two hydrogen atoms, along with one calcium atom.
By looking at a formula, then, we understand the ratio in which the different atoms are united to form the molecule.
1.4.2 Significance Of A Molecular Formula
The molecular formula of a compound has quantitative significance. It does not merely name the compound; it carries numbers that can be worked with.
It represents: (i) both the molecule and the molecular mass of the compound; (ii) the respective numbers of different atoms present in one molecule of the compound; (iii) the ratios of the respective masses of the elements present in the compound.
Take the formula \(\\mathrm{CO\_2}\). It tells us that: (i) the molecular formula of carbon dioxide is \(\\mathrm{CO\_2}\); (ii) each molecule contains one carbon atom joined by chemical bonds with two oxygen atoms; (iii) the molecular mass of carbon dioxide is 44, since the atomic mass of carbon is 12 and that of oxygen is 16; and carbon and oxygen combine with each other in the ratio 3 : 8, because 12 parts of carbon combine with 32 parts of oxygen.
That last line is worth checking for yourself. Two oxygen atoms give \(16 \\times 2 = 32\), and the ratio \(12 : 32\) divides through by 4 to give \(3 : 8\).
This is the point at which a formula stops being a name and becomes a tool. From it you will later calculate molecular masses, percentage compositions and the masses that react in the laboratory.
1.5 Radicals
Look closely at the formula of a compound and you will usually find it falls into two parts. These two parts are known separately as radicals.
A molecule of potassium chloride has two parts, potassium and chloride, so potassium is one radical and chloride is the other. In the same way, magnesium sulphate has magnesium and sulphate as its two radicals.
The words single unit are the important ones. A sulphate radical contains five atoms, yet through an entire reaction it travels, combines and separates as one piece.
1.5.1 Basic And Acidic Radicals
Radicals carrying a positive charge — the cations — are called basic radicals. They are formed when a base loses its hydroxide ion, \(\\mathrm{OH^{-}}\).
For instance, \(\\mathrm{Na^{+}}\) is formed by the removal of \(\\mathrm{OH^{-}}\) from the base sodium hydroxide, \(\\mathrm{NaOH}\); and \(\\mathrm{Mg^{2+}}\) is formed by the removal of \(\\mathrm{OH^{-}}\) from the base magnesium hydroxide, \(\\mathrm{Mg(OH)\_2}\).
Radicals carrying a negative charge — the anions — are called acidic radicals. They are formed by the removal of the hydrogen ion, \(\\mathrm{H^{+}}\), from acids.
So \(\\mathrm{Cl^{-}}\) is formed by removing \(\\mathrm{H^{+}}\) from hydrochloric acid, \(\\mathrm{HCl}\), and \(\\mathrm{SO\_4^{2-}}\) is formed by removing \(\\mathrm{H^{+}}\) from sulphuric acid, \(\\mathrm{H\_2SO\_4}\).
The names make sense once you see where each radical comes from: the basic radical is the part left over from a base, and the acidic radical is the part left over from an acid.
1.5.2 Reactivity Of Radicals
A radical is not a comfortable particle. It carries a charge, and a charged particle is always looking for the opposite charge.
Basic radicals such as \(\\mathrm{Na^{+}}\) and \(\\mathrm{Mg^{2+}}\) are positively charged and electron-deficient. They tend to accept electrons in order to attain a stable configuration, and they are attracted towards negative charges.
Acidic radicals such as \(\\mathrm{Cl^{-}}\) and \(\\mathrm{SO\_4^{2-}}\) are negatively charged and electron-rich. They tend to donate electrons to achieve stability, and they are attracted towards positive charges.
Because they are charged and unstable on their own, both kinds of radical are highly reactive. Set them free and they will react to form compounds.
Radicals Build Salts
Put the two together and the charges cancel. Acidic and basic radicals combine to form salts:
\[\\mathrm{Na^{+}} + \\mathrm{Cl^{-}} \\rightarrow \\mathrm{NaCl}\] \[\\mathrm{Ca^{2+}} + \\mathrm{SO\_4^{2-}} \\rightarrow \\mathrm{CaSO\_4}\]
And when a salt is dissolved in water, it splits up again into its constituent radicals, which is why solutions of salts behave so differently from the dry solid.
Where do the two radicals come from in the first place? An acid reacts with a base to produce a salt and water, a change called neutralisation.
potassium hydroxide (base) + hydrochloric acid (acid) → potassium chloride (salt) + water
In the formation of potassium chloride, the potassium radical has been contributed by the base potassium hydroxide, while the chloride radical has been contributed by hydrochloric acid. FC 1.1 traces the whole journey.
Flowchart
FC 1.1 — Where The Two Radicals In A Salt Come From Potassium hydroxide (base) → loses OH⁻ → potassium radical K⁺ (basic, electropositive) → combines with → chloride radical Cl⁻ (acidic, electronegative) ← formed by loss of H⁺ ← hydrochloric acid (acid); the combination gives potassium chloride (salt) → dissolved in water → splits again into K⁺ and Cl⁻
Thus, in potassium chloride, potassium is an electropositive (or basic) radical and chloride is an electronegative (or acidic) radical. Similarly, in magnesium sulphate, magnesium is the electropositive radical while sulphate is the electronegative one, as Fig. 1.5-1 shows.
Diagram
Fig. 1.5-1 — A Salt Splitting Into Its Two Radicals A beaker of water with a few crystals of magnesium sulphate at the bottom, arrows leading from the crystals into the solution where two kinds of particle are drawn spread through the liquid: circles marked Mg²⁺ labelled ‘basic (electropositive) radical’ and larger circles marked SO₄²⁻ labelled ‘acidic (electronegative) radical’. Label the whole beaker ‘magnesium sulphate dissolved in water’.
1.5.3 Simple And Compound Radicals
A radical is called a simple radical when it is an atom only, such as sodium, \(\\mathrm{Na^{+}}\), or magnesium, \(\\mathrm{Mg^{2+}}\).
It is known as a compound radical, or complex radical, when it is made up of a group of two or more different atoms — that is, when it is polyatomic. Sulphate, \(\\mathrm{SO\_4^{2-}}\), is a compound radical, made up of one sulphur atom and four oxygen atoms.
1.5.4 Valency Of Radicals
Radicals have their own combining power, and it is according to this combining power that they form chemical formulae.
The charge on a radical is usually its valency. Depending on the charge, a radical may be monovalent with a valency of one, divalent with a valency of two, trivalent with a valency of three, or tetravalent with a valency of four.
Consider ammonium carbonate, \(\\mathrm{(NH\_4)\_2CO\_3}\). Ammonium, \(\\mathrm{NH\_4^{+}}\), is a basic radical with a combining power of 1, and carbonate, \(\\mathrm{CO\_3^{2-}}\), is an acidic radical with a combining power of 2. Two ammonium radicals are therefore needed to balance one carbonate radical, and that is why the formula carries the bracket and the subscript 2.
The lists you will need are given in Table 1.7 and Table 1.8. Do not try to memorise them in one sitting; use them as you work, and the common ones will stay with you.
Table 1.7 : List of some common electrovalent positive ions (basic radicals)
| Monovalent electropositive | Divalent electropositive | Trivalent electropositive | |||
|---|---|---|---|---|---|
| Ammonium | \(\\mathrm{NH\_4^{+}}\) | Argentic [silver(II)] | \(\\mathrm{Ag^{2+}}\) | Aluminium | \(\\mathrm{Al^{3+}}\) |
| Aurous [gold(I)] | \(\\mathrm{Au^{+}}\) | Barium | \(\\mathrm{Ba^{2+}}\) | Arsenic | \(\\mathrm{As^{3+}}\) |
| Argentous [silver(I)] | \(\\mathrm{Ag^{+}}\) | Calcium | \(\\mathrm{Ca^{2+}}\) | Auric [gold(III)] | \(\\mathrm{Au^{3+}}\) |
| Cuprous [copper(I)] | \(\\mathrm{Cu^{+}}\) | Cupric [copper(II)] | \(\\mathrm{Cu^{2+}}\) | Bismuth | \(\\mathrm{Bi^{3+}}\) |
| Hydrogen | \(\\mathrm{H^{+}}\) | Ferrous [iron(II)] | \(\\mathrm{Fe^{2+}}\) | Chromium | \(\\mathrm{Cr^{3+}}\) |
| Lithium | \(\\mathrm{Li^{+}}\) | Magnesium | \(\\mathrm{Mg^{2+}}\) | Ferric [iron(III)] | \(\\mathrm{Fe^{3+}}\) |
| Sodium | \(\\mathrm{Na^{+}}\) | Manganese | \(\\mathrm{Mn^{2+}}\) | Tetravalent electropositive | |
| Potassium | \(\\mathrm{K^{+}}\) | Mercuric [mercury(II)] | \(\\mathrm{Hg^{2+}}\) | Plumbic [lead(IV)] | \(\\mathrm{Pb^{4+}}\) |
| Mercurous [mercury(I)] | \(\\mathrm{Hg^{+}}\) | Nickel | \(\\mathrm{Ni^{2+}}\) | Platinic [platinum(IV)] | \(\\mathrm{Pt^{4+}}\) |
| Plumbous [lead(II)] | \(\\mathrm{Pb^{2+}}\) | Stannic [tin(IV)] | \(\\mathrm{Sn^{4+}}\) | ||
| Platinous [platinum(II)] | \(\\mathrm{Pt^{2+}}\) | ||||
| Stannous [tin(II)] | \(\\mathrm{Sn^{2+}}\) | ||||
| Zinc | \(\\mathrm{Zn^{2+}}\) |
Table 1.8 : List of some common electrovalent negative ions (acid radicals)
| Monovalent electronegative | Divalent electronegative | Trivalent electronegative | |||
|---|---|---|---|---|---|
| Acetate | \(\\mathrm{CH\_3COO^{-}}\) | Carbonate | \(\\mathrm{CO\_3^{2-}}\) | Arsenate | \(\\mathrm{AsO\_4^{3-}}\) |
| Bicarbonate or hydrogen carbonate | \(\\mathrm{HCO\_3^{-}}\) | Dichromate | \(\\mathrm{Cr\_2O\_7^{2-}}\) | Nitride | \(\\mathrm{N^{3-}}\) |
| Bisulphide or hydrogen sulphide | \(\\mathrm{HS^{-}}\) | Oxide | \(\\mathrm{O^{2-}}\) | Aluminate | \(\\mathrm{AlO\_3^{3-}}\) |
| Bisulphate or hydrogen sulphate | \(\\mathrm{HSO\_4^{-}}\) | Peroxide | \(\\mathrm{O\_2^{2-}}\) | Arsenite | \(\\mathrm{AsO\_3^{3-}}\) |
| Bisulphite or hydrogen sulphite | \(\\mathrm{HSO\_3^{-}}\) | Sulphate | \(\\mathrm{SO\_4^{2-}}\) | Phosphide | \(\\mathrm{P^{3-}}\) |
| Bromide | \(\\mathrm{Br^{-}}\) | Sulphite | \(\\mathrm{SO\_3^{2-}}\) | Phosphite | \(\\mathrm{PO\_3^{3-}}\) |
| Chloride | \(\\mathrm{Cl^{-}}\) | Sulphide | \(\\mathrm{S^{2-}}\) | Phosphate | \(\\mathrm{PO\_4^{3-}}\) |
| Permanganate | \(\\mathrm{MnO\_4^{-}}\) | Silicate | \(\\mathrm{SiO\_3^{2-}}\) | Borate | \(\\mathrm{BO\_3^{3-}}\) |
| Fluoride | \(\\mathrm{F^{-}}\) | Thiosulphate | \(\\mathrm{S\_2O\_3^{2-}}\) | Tetravalent electronegative | |
| Hydride | \(\\mathrm{H^{-}}\) | Zincate | \(\\mathrm{ZnO\_2^{2-}}\) | Carbide | \(\\mathrm{C^{4-}}\) |
| Hydroxide | \(\\mathrm{OH^{-}}\) | Plumbite | \(\\mathrm{PbO\_2^{2-}}\) | Ferrocyanide | \(\\mathrm{Fe(CN)\_6^{4-}}\) |
| Iodide | \(\\mathrm{I^{-}}\) | Stannate | \(\\mathrm{SnO\_3^{2-}}\) | ||
| Cyanide | \(\\mathrm{CN^{-}}\) | Manganate | \(\\mathrm{MnO\_4^{2-}}\) | ||
| Nitrate | \(\\mathrm{NO\_3^{-}}\) | Chromate | \(\\mathrm{CrO\_4^{2-}}\) | ||
| Nitrite | \(\\mathrm{NO\_2^{-}}\) | Oxalate | \(\\mathrm{(COO)\_2^{2-}}\) | ||
| Chlorite | \(\\mathrm{ClO\_2^{-}}\) | ||||
| Hypochlorite | \(\\mathrm{ClO^{-}}\) | ||||
| Chlorate | \(\\mathrm{ClO\_3^{-}}\) | ||||
| Perchlorate | \(\\mathrm{ClO\_4^{-}}\) | ||||
| Meta-aluminate | \(\\mathrm{AlO\_2^{-}}\) |
1.6 Compounds — Name And Formulae
You now have everything you need to build a formula: symbols for the elements, valencies for the radicals, and lists of the radicals themselves. What remains is a reliable way of putting them together.
1.6.1 Writing Formulae By The Criss-Cross Method
The steps below should be taken while attempting to write the formula of a compound. The method is called the criss-cross method, because the two valencies cross over each other.
- Write the symbols side by side, the basic radical first and then the acidic radical.
- Write the valency of each radical on top of its symbol.
- Ignore the (+) or (−) signs of the radicals.
- Divide the valency numbers by their highest common factor (H.C.F.), if any, to get the simplest ratio.
- Interchange the valencies of the radicals.
- Write the interchanged valency numbers to the lower right of the radicals. If the radical is a group of atoms and has a valency number more than 1, enclose it within parentheses.
Step (iv) is the one most often forgotten. A formula must give the simplest whole-number ratio, so if both valencies are even they must first be cut down. Step (vi) matters for the same practical reason: without brackets, \(\\mathrm{Al(OH)\_3}\) would read as though the 3 belonged only to the hydrogen.
Fig. 1.6-1 shows the crossing itself, and Table 1.9 works through six compounds step by step.
Diagram
Fig. 1.6-1 — The Criss-Cross Method Two worked panels. Upper panel: the symbols Al and SO₄ written side by side, with the valency 3 written above Al and 2 above SO₄; two long crossing arrows run from the 3 down to SO₄ and from the 2 down to Al; below, the result Al₂(SO₄)₃ with a callout on the bracket reading ‘brackets are needed because the sulphate radical is taken three times’. Lower panel: Pb with valency 4 and O with valency 2, an arrow marked ‘÷ H.C.F. 2’ converting them to 2 and 1, then the crossing arrows, then the result PbO₂.
Table 1.9 : Formulae written by the criss-cross method
| Name of compound | Symbols with valencies and charge | Exchange of valency | Formula |
|---|---|---|---|
| Magnesium chloride | \(\\mathrm{Mg^{2+}}\\ \\mathrm{Cl^{1-}}\) | \(\\mathrm{Mg}^{2}\\ \\mathrm{Cl}^{1} \\rightarrow \\mathrm{Mg\_1\\ Cl\_2}\) | \(\\mathrm{MgCl\_2}\) |
| Calcium oxide | \(\\mathrm{Ca^{2+}}\\ \\mathrm{O^{2-}}\); dividing by H.C.F. it becomes \(\\mathrm{Ca^{1+}}\\ \\mathrm{O^{1-}}\) | \(\\mathrm{Ca}^{1}\\ \\mathrm{O}^{1} \\rightarrow \\mathrm{Ca\_1\\ O\_1}\) | \(\\mathrm{CaO}\) |
| Aluminium hydroxide | \(\\mathrm{Al^{3+}}\\ \\mathrm{(OH)^{1-}}\) | \(\\mathrm{Al}^{3}\\ \\mathrm{(OH)}^{1} \\rightarrow \\mathrm{Al\_1\\ (OH)\_3}\) | \(\\mathrm{Al(OH)\_3}\) |
| Ferric oxide | \(\\mathrm{Fe^{3+}}\\ \\mathrm{O^{2-}}\) | \(\\mathrm{Fe}^{3}\\ \\mathrm{O}^{2} \\rightarrow \\mathrm{Fe\_2\\ O\_3}\) | \(\\mathrm{Fe\_2O\_3}\) |
| Sodium meta-aluminate | \(\\mathrm{Na^{+}}\\ \\mathrm{AlO\_2^{-}}\) | \(\\mathrm{Na}^{1}\\ \\mathrm{(AlO\_2)}^{1} \\rightarrow \\mathrm{Na\_1\\ (AlO\_2)\_1}\) | \(\\mathrm{NaAlO\_2}\) |
| Sodium aluminate | \(\\mathrm{Na^{+}}\\ \\mathrm{AlO\_3^{3-}}\) | \(\\mathrm{Na}^{1}\\ \\mathrm{(AlO\_3)}^{3} \\rightarrow \\mathrm{Na\_3\\ (AlO\_3)\_1}\) | \(\\mathrm{Na\_3AlO\_3}\) |
Notice that a subscript of 1 is never written, and a bracket around a single group is dropped. Chemists write \(\\mathrm{CaO}\), not \(\\mathrm{Ca\_1O\_1}\).
1.6.2 Important Compounds And Their Formulae
Table 1.10 collects the compounds you will meet again and again in this book and in the laboratory. Work through a few of them with the criss-cross method and check that the formula printed is the one you get.
Table 1.10 : Important compounds and their formulae
| Chemical name | Symbols with charge | Formula |
|---|---|---|
| Potassium chloride | \(\\mathrm{K^{1+}Cl^{1-}}\) | \(\\mathrm{KCl}\) |
| Potassium bromide | \(\\mathrm{K^{1+}Br^{1-}}\) | \(\\mathrm{KBr}\) |
| Potassium iodide | \(\\mathrm{K^{1+}I^{1-}}\) | \(\\mathrm{KI}\) |
| Potassium hydroxide | \(\\mathrm{K^{1+}OH^{1-}}\) | \(\\mathrm{KOH}\) |
| Potassium nitrite | \(\\mathrm{K^{1+}NO\_2^{1-}}\) | \(\\mathrm{KNO\_2}\) |
| Potassium nitrate | \(\\mathrm{K^{1+}NO\_3^{1-}}\) | \(\\mathrm{KNO\_3}\) |
| Potassium hydrogen carbonate | \(\\mathrm{K^{1+}HCO\_3^{1-}}\) | \(\\mathrm{KHCO\_3}\) |
| Potassium hydrogen sulphite | \(\\mathrm{K^{1+}HSO\_3^{1-}}\) | \(\\mathrm{KHSO\_3}\) |
| Potassium hydrogen sulphate | \(\\mathrm{K^{1+}HSO\_4^{1-}}\) | \(\\mathrm{KHSO\_4}\) |
| Potassium metaluminate | \(\\mathrm{K^{1+}AlO\_2^{1-}}\) | \(\\mathrm{KAlO\_2}\) |
| Potassium permanganate | \(\\mathrm{K^{1+}MnO\_4^{1-}}\) | \(\\mathrm{KMnO\_4}\) |
| Potassium sulphite | \(\\mathrm{K^{1+}SO\_3^{2-}}\) | \(\\mathrm{K\_2SO\_3}\) |
| Potassium sulphate | \(\\mathrm{K^{1+}SO\_4^{2-}}\) | \(\\mathrm{K\_2SO\_4}\) |
| Potassium carbonate | \(\\mathrm{K^{1+}CO\_3^{2-}}\) | \(\\mathrm{K\_2CO\_3}\) |
| Potassium dichromate | \(\\mathrm{K^{1+}Cr\_2O\_7^{2-}}\) | \(\\mathrm{K\_2Cr\_2O\_7}\) |
| Potassium zincate | \(\\mathrm{K^{1+}ZnO\_2^{2-}}\) | \(\\mathrm{K\_2ZnO\_2}\) |
| Potassium plumbite | \(\\mathrm{K^{1+}PbO\_2^{2-}}\) | \(\\mathrm{K\_2PbO\_2}\) |
| Sodium chloride | \(\\mathrm{Na^{1+}Cl^{1-}}\) | \(\\mathrm{NaCl}\) |
| Sodium hydroxide | \(\\mathrm{Na^{1+}OH^{1-}}\) | \(\\mathrm{NaOH}\) |
| Sodium nitrite | \(\\mathrm{Na^{1+}NO\_2^{1-}}\) | \(\\mathrm{NaNO\_2}\) |
| Sodium nitrate | \(\\mathrm{Na^{1+}NO\_3^{1-}}\) | \(\\mathrm{NaNO\_3}\) |
| Sodium hydrogen carbonate | \(\\mathrm{Na^{1+}HCO\_3^{1-}}\) | \(\\mathrm{NaHCO\_3}\) |
| Sodium hydrogen sulphite | \(\\mathrm{Na^{1+}HSO\_3^{1-}}\) | \(\\mathrm{NaHSO\_3}\) |
| Sodium hydrogen sulphate | \(\\mathrm{Na^{1+}HSO\_4^{1-}}\) | \(\\mathrm{NaHSO\_4}\) |
| Sodium metaluminate | \(\\mathrm{Na^{1+}AlO\_2^{1-}}\) | \(\\mathrm{NaAlO\_2}\) |
| Sodium sulphite | \(\\mathrm{Na^{1+}SO\_3^{2-}}\) | \(\\mathrm{Na\_2SO\_3}\) |
| Sodium sulphate | \(\\mathrm{Na^{1+}SO\_4^{2-}}\) | \(\\mathrm{Na\_2SO\_4}\) |
| Sodium carbonate | \(\\mathrm{Na^{1+}CO\_3^{2-}}\) | \(\\mathrm{Na\_2CO\_3}\) |
| Sodium zincate | \(\\mathrm{Na^{1+}ZnO\_2^{2-}}\) | \(\\mathrm{Na\_2ZnO\_2}\) |
| Sodium plumbite | \(\\mathrm{Na^{1+}PbO\_2^{2-}}\) | \(\\mathrm{Na\_2PbO\_2}\) |
| Silver chloride | \(\\mathrm{Ag^{1+}Cl^{1-}}\) | \(\\mathrm{AgCl}\) |
| Ammonium chloride | \(\\mathrm{NH\_4^{1+}Cl^{1-}}\) | \(\\mathrm{NH\_4Cl}\) |
| Ammonium sulphate | \(\\mathrm{NH\_4^{1+}SO\_4^{2-}}\) | \(\\mathrm{(NH\_4)\_2SO\_4}\) |
| Ammonium hydroxide | \(\\mathrm{NH\_4^{1+}OH^{1-}}\) | \(\\mathrm{NH\_4OH}\) |
| Calcium chloride | \(\\mathrm{Ca^{2+}Cl^{1-}}\) | \(\\mathrm{CaCl\_2}\) |
| Calcium hydroxide | \(\\mathrm{Ca^{2+}OH^{1-}}\) | \(\\mathrm{Ca(OH)\_2}\) |
| Calcium nitrate | \(\\mathrm{Ca^{2+}NO\_3^{1-}}\) | \(\\mathrm{Ca(NO\_3)\_2}\) |
| Calcium hydrogen carbonate | \(\\mathrm{Ca^{2+}HCO\_3^{1-}}\) | \(\\mathrm{Ca(HCO\_3)\_2}\) |
| Calcium hydrogen sulphite | \(\\mathrm{Ca^{2+}HSO\_3^{1-}}\) | \(\\mathrm{Ca(HSO\_3)\_2}\) |
| Calcium sulphite | \(\\mathrm{Ca^{2+}SO\_3^{2-}}\) | \(\\mathrm{CaSO\_3}\) |
| Calcium sulphate | \(\\mathrm{Ca^{2+}SO\_4^{2-}}\) | \(\\mathrm{CaSO\_4}\) |
| Calcium carbonate | \(\\mathrm{Ca^{2+}CO\_3^{2-}}\) | \(\\mathrm{CaCO\_3}\) |
| Calcium oxide | \(\\mathrm{Ca^{2+}O^{2-}}\) | \(\\mathrm{CaO}\) |
| Calcium silicate | \(\\mathrm{Ca^{2+}SiO\_3^{2-}}\) | \(\\mathrm{CaSiO\_3}\) |
| Calcium nitride | \(\\mathrm{Ca^{2+}N^{3-}}\) | \(\\mathrm{Ca\_3N\_2}\) |
| Magnesium chloride | \(\\mathrm{Mg^{2+}Cl^{1-}}\) | \(\\mathrm{MgCl\_2}\) |
| Magnesium hydroxide | \(\\mathrm{Mg^{2+}OH^{1-}}\) | \(\\mathrm{Mg(OH)\_2}\) |
| Magnesium nitrate | \(\\mathrm{Mg^{2+}NO\_3^{1-}}\) | \(\\mathrm{Mg(NO\_3)\_2}\) |
| Magnesium oxide | \(\\mathrm{Mg^{2+}O^{2-}}\) | \(\\mathrm{MgO}\) |
| Magnesium nitride | \(\\mathrm{Mg^{2+}N^{3-}}\) | \(\\mathrm{Mg\_3N\_2}\) |
| Zinc chloride | \(\\mathrm{Zn^{2+}Cl^{1-}}\) | \(\\mathrm{ZnCl\_2}\) |
| Zinc hydroxide | \(\\mathrm{Zn^{2+}OH^{1-}}\) | \(\\mathrm{Zn(OH)\_2}\) |
| Zinc nitrate | \(\\mathrm{Zn^{2+}NO\_3^{1-}}\) | \(\\mathrm{Zn(NO\_3)\_2}\) |
| Zinc sulphate | \(\\mathrm{Zn^{2+}SO\_4^{2-}}\) | \(\\mathrm{ZnSO\_4}\) |
| Zinc carbonate | \(\\mathrm{Zn^{2+}CO\_3^{2-}}\) | \(\\mathrm{ZnCO\_3}\) |
| Zinc oxide | \(\\mathrm{Zn^{2+}O^{2-}}\) | \(\\mathrm{ZnO}\) |
| Lead(II) chloride | \(\\mathrm{Pb^{2+}Cl^{1-}}\) | \(\\mathrm{PbCl\_2}\) |
| Lead(II) bromide | \(\\mathrm{Pb^{2+}Br^{1-}}\) | \(\\mathrm{PbBr\_2}\) |
| Lead(II) hydroxide | \(\\mathrm{Pb^{2+}OH^{1-}}\) | \(\\mathrm{Pb(OH)\_2}\) |
| Lead(II) nitrate | \(\\mathrm{Pb^{2+}NO\_3^{1-}}\) | \(\\mathrm{Pb(NO\_3)\_2}\) |
| Lead(II) sulphate | \(\\mathrm{Pb^{2+}SO\_4^{2-}}\) | \(\\mathrm{PbSO\_4}\) |
| Lead(II) oxide | \(\\mathrm{Pb^{2+}O^{2-}}\) | \(\\mathrm{PbO}\) |
| Manganese chloride | \(\\mathrm{Mn^{2+}Cl^{1-}}\) | \(\\mathrm{MnCl\_2}\) |
| Manganese sulphate | \(\\mathrm{Mn^{2+}SO\_4^{2-}}\) | \(\\mathrm{MnSO\_4}\) |
| Aluminium chloride | \(\\mathrm{Al^{3+}Cl^{1-}}\) | \(\\mathrm{AlCl\_3}\) |
| Aluminium sulphate | \(\\mathrm{Al^{3+}SO\_4^{2-}}\) | \(\\mathrm{Al\_2(SO\_4)\_3}\) |
| Aluminium hydroxide | \(\\mathrm{Al^{3+}OH^{1-}}\) | \(\\mathrm{Al(OH)\_3}\) |
| Aluminium sulphide | \(\\mathrm{Al^{3+}S^{2-}}\) | \(\\mathrm{Al\_2S\_3}\) |
| Aluminium oxide | \(\\mathrm{Al^{3+}O^{2-}}\) | \(\\mathrm{Al\_2O\_3}\) |
| Chromium chloride | \(\\mathrm{Cr^{3+}Cl^{1-}}\) | \(\\mathrm{CrCl\_3}\) |
| Chromium sulphate | \(\\mathrm{Cr^{3+}SO\_4^{2-}}\) | \(\\mathrm{Cr\_2(SO\_4)\_3}\) |
| Chromium oxide | \(\\mathrm{Cr^{3+}O^{2-}}\) | \(\\mathrm{Cr\_2O\_3}\) |
| Copper(I) chloride | \(\\mathrm{Cu^{1+}Cl^{1-}}\) | \(\\mathrm{CuCl}\) |
| Copper(I) oxide | \(\\mathrm{Cu^{1+}O^{2-}}\) | \(\\mathrm{Cu\_2O}\) |
| Copper(I) sulphide | \(\\mathrm{Cu^{1+}S^{2-}}\) | \(\\mathrm{Cu\_2S}\) |
| Copper(II) chloride | \(\\mathrm{Cu^{2+}Cl^{1-}}\) | \(\\mathrm{CuCl\_2}\) |
| Copper(II) hydroxide | \(\\mathrm{Cu^{2+}OH^{1-}}\) | \(\\mathrm{Cu(OH)\_2}\) |
| Copper(II) nitrate | \(\\mathrm{Cu^{2+}NO\_3^{1-}}\) | \(\\mathrm{Cu(NO\_3)\_2}\) |
| Copper(II) sulphate | \(\\mathrm{Cu^{2+}SO\_4^{2-}}\) | \(\\mathrm{CuSO\_4}\) |
| Copper(II) sulphide | \(\\mathrm{Cu^{2+}S^{2-}}\) | \(\\mathrm{CuS}\) |
| Copper(II) oxide | \(\\mathrm{Cu^{2+}O^{2-}}\) | \(\\mathrm{CuO}\) |
| Iron(II) chloride | \(\\mathrm{Fe^{2+}Cl^{1-}}\) | \(\\mathrm{FeCl\_2}\) |
| Iron(II) hydroxide | \(\\mathrm{Fe^{2+}OH^{1-}}\) | \(\\mathrm{Fe(OH)\_2}\) |
| Iron(II) nitrate | \(\\mathrm{Fe^{2+}NO\_3^{1-}}\) | \(\\mathrm{Fe(NO\_3)\_2}\) |
| Iron(II) sulphate | \(\\mathrm{Fe^{2+}SO\_4^{2-}}\) | \(\\mathrm{FeSO\_4}\) |
| Iron(II) sulphide | \(\\mathrm{Fe^{2+}S^{2-}}\) | \(\\mathrm{FeS}\) |
| Iron(II) oxide | \(\\mathrm{Fe^{2+}O^{2-}}\) | \(\\mathrm{FeO}\) |
| Iron(III) chloride | \(\\mathrm{Fe^{3+}Cl^{1-}}\) | \(\\mathrm{FeCl\_3}\) |
| Iron(III) sulphate | \(\\mathrm{Fe^{3+}SO\_4^{2-}}\) | \(\\mathrm{Fe\_2(SO\_4)\_3}\) |
| Iron(III) hydroxide | \(\\mathrm{Fe^{3+}OH^{1-}}\) | \(\\mathrm{Fe(OH)\_3}\) |
| Iron(III) sulphide | \(\\mathrm{Fe^{3+}S^{2-}}\) | \(\\mathrm{Fe\_2S\_3}\) |
| Iron(III) nitrate | \(\\mathrm{Fe^{3+}NO\_3^{1-}}\) | \(\\mathrm{Fe(NO\_3)\_3}\) |
| Iron(III) oxide | \(\\mathrm{Fe^{3+}O^{2-}}\) | \(\\mathrm{Fe\_2O\_3}\) |
1.7 Naming Compounds
Writing a formula and reading one are two halves of the same skill. This section runs the process backwards: given the formula, what is the compound called?
1.7.1 Steps In Naming A Compound
The following steps are to be taken while writing the name of a compound:
- Identify the basic and acidic radicals in the formula.
- Write the name of the compound in the same order of radicals as given in the formula.
Take \(\\mathrm{Ca(OH)\_2}\). The basic radical is \(\\mathrm{Ca^{2+}}\), calcium, and the acidic radical is \(\\mathrm{OH^{-}}\), hydroxide. The name of the compound is calcium hydroxide.
Take \(\\mathrm{Na\_2C\_2O\_4}\). The basic radical is \(\\mathrm{Na^{+}}\), sodium, and the acidic radical is \(\\mathrm{C\_2O\_4^{2-}}\), oxalate. The name of the compound is sodium oxalate.
Notice that the subscripts play no part in the name. They were fixed by the valencies, and the valencies are already implied by the two radical names.
1.7.2 Naming Different Types Of Compounds
Different types of compounds can be named as follows.
1. Compounds With A Metal And A Non-Metal
The metal is written first, followed by the non-metal, and the suffix ide is added to the non-metal.
calcium + nitrogen → calcium nitride, \(\\mathrm{Ca\_3N\_2}\)
So nitrogen becomes nitride, oxygen becomes oxide, sulphur becomes sulphide and chlorine becomes chloride.
2. Compounds Of Variable Valency
Elements like gold, iron, copper, mercury, lead and tin exhibit more than one valency. According to the latest IUPAC system of nomenclature, the valency is written in Roman numerals in brackets with the basic radical.
Gold(I) sulphate is \(\\mathrm{Au\_2SO\_4}\), formerly called aurous sulphate. Gold(III) chloride is \(\\mathrm{AuCl\_3}\), formerly called auric chloride. The older suffixes -ous and -ic were used to represent the lower and the higher valency respectively.
Acidic radicals show variable valency too, and their names change with it. Table 1.11 gives the pairs you will meet.
Table 1.11 : Acidic radicals with variable valency
| Acidic radical | Compound | Name of the compound |
|---|---|---|
| \(\\mathrm{O^{2-}}\) | \(\\mathrm{Na\_2O}\) | Sodium oxide |
| \(\\mathrm{O\_2^{2-}}\) | \(\\mathrm{Na\_2O\_2}\) | Sodium peroxide |
| \(\\mathrm{MnO\_4^{-}}\) | \(\\mathrm{KMnO\_4}\) | Potassium permanganate |
| \(\\mathrm{MnO\_4^{2-}}\) | \(\\mathrm{K\_2MnO\_4}\) | Potassium manganate |
| \(\\mathrm{\[Fe(CN)\_6]^{3-}}\) | \(\\mathrm{K\_3\[Fe(CN)\_6]}\) | Potassium ferricyanide |
| \(\\mathrm{\[Fe(CN)\_6]^{4-}}\) | \(\\mathrm{K\_4\[Fe(CN)\_6]}\) | Potassium ferrocyanide |
3. Compounds With Non-Metals Having Variable Valency
Here no Roman numeral is used. Instead the prefix tri, tetra, penta and so on is added, as the case may be, to the second non-metal, to say how many of its atoms are present.
\(\\mathrm{PCl\_3}\) is phosphorus trichloride and \(\\mathrm{PCl\_5}\) is phosphorus pentachloride. The prefix does the same work the Roman numeral does for a metal — it removes the ambiguity between two compounds of the same pair of elements.
4. Compounds With Different Numbers Of Oxygen Atoms
Oxygen is represented at the end of the formula, and the name of the compound depends on the number of oxygen atoms present in it.
The prefix hypo is used if the number of oxygen atoms is less than 2: \(\\mathrm{KClO}\) — potassium hypochlorite
The suffix ite is used if the number of oxygen atoms is 2: \(\\mathrm{KClO\_2}\) — potassium chlorite
The suffix ate is used if the number of oxygen atoms is 3: \(\\mathrm{KClO\_3}\) — potassium chlorate
The prefix per is used when the number of oxygen atoms is more than 3: \(\\mathrm{KClO\_4}\) — potassium perchlorate
1.7.3 Naming Of Acids
(a) Binary acids. Acids composed of two elements are named by adding the prefix hydro and the suffix ic to the name of the second element.
\(\\mathrm{HCl}\) is hydrochloric acid, and \(\\mathrm{HF}\) is hydrofluoric acid.
(b) Acids containing polyatomic radicals. These are named on the basis of the second element present in the molecule, and the prefix hydro is not used.
In \(\\mathrm{H\_2SO\_4}\) the second element is sulphur, hence the name sulphuric acid. In \(\\mathrm{HNO\_3}\) the second element is nitrogen, hence nitric acid. In \(\\mathrm{H\_3PO\_4}\) the second element is phosphorus, hence phosphoric acid.
If the number of oxygen atoms is less, the suffix ous is used instead of ic. Thus \(\\mathrm{H\_2SO\_3}\) is sulphurous acid and \(\\mathrm{HNO\_2}\) is nitrous acid.
The endings run in step with the endings of the radicals: an -ic acid gives an -ate salt, and an -ous acid gives an -ite salt. Table 1.12 shows the whole family for one element.
Table 1.12 : Acids, their acidic radicals and their sodium salts
| Acid | Acidic radical | Compound formed |
|---|---|---|
| \(\\mathrm{HClO}\), hypochlorous acid | \(\\mathrm{ClO^{-}}\), hypochlorite | \(\\mathrm{NaClO}\), sodium hypochlorite |
| \(\\mathrm{HClO\_2}\), chlorous acid | \(\\mathrm{ClO\_2^{-}}\), chlorite | \(\\mathrm{NaClO\_2}\), sodium chlorite |
| \(\\mathrm{HClO\_3}\), chloric acid | \(\\mathrm{ClO\_3^{-}}\), chlorate | \(\\mathrm{NaClO\_3}\), sodium chlorate |
| \(\\mathrm{HClO\_4}\), perchloric acid | \(\\mathrm{ClO\_4^{-}}\), perchlorate | \(\\mathrm{NaClO\_4}\), sodium perchlorate |
Names That Follow No Rule
There are certain compounds whose names do not follow any systematic rule at all. Such names are called trivial or common names, and they are widely accepted.
Nitrogen trihydride is called ammonia, \(\\mathrm{NH\_3}\). Hydrogen monoxide, or dihydrogen oxide, is called water, \(\\mathrm{H\_2O}\). Nobody asks for a glass of dihydrogen oxide, and no chemist objects.
1.7.4 Finding The Valency From The Formula
The valency of elements can be determined from a formula, using the valencies of the radicals and the fact that:
hydrogen \(\\mathrm{\[H]} = 1\); oxygen \(\\mathrm{\[O]} = 2\); chlorine \(\\mathrm{\[Cl]} = 1\)
The procedure is set out in FC 1.2. It is the criss-cross method run in reverse: instead of crossing valencies down into subscripts, we lift the subscripts back up into valencies.
Flowchart
FC 1.2 — Finding The Valency Of An Element From Its Formula Write the given formula (NO₂) → interchange the subscripts and write each as a superscript on the other symbol (N²O¹) → the valency of oxygen is known to be 2, so multiply both superscripts by 2 (N²ˣ²O¹ˣ²) → read off the result (N⁴O²) → the valency of nitrogen is 4
Thus, from the formula \(\\mathrm{NO\_2}\), we find that the valency of nitrogen is 4.
1.8 Free Radicals
The radicals of Section 1.5 were charged particles, pulled out of salts. There is a second kind of reactive particle that shares the name but not the charge, and this section is about it.
1.8.1 What A Free Radical Is
A free radical is an atom, molecule or ion that contains at least one unpaired valence electron.
Normally the electrons in a molecule exist in pairs. When a particle is left with one electron standing alone without a partner, it becomes highly unstable and exceptionally reactive. Such a particle is a free radical, and free radicals are short-lived species.
They are called free radicals because they can exist independently, and because their unpaired electron is free to react easily with other atoms or molecules. That is what makes them so much more reactive than ordinary stable particles.
A free radical is written with a dot to show the unpaired electron. Chlorine radical is \(\\mathrm{Cl}\\bullet\), hydroxyl radical is \(\\mathrm{OH}\\bullet\), and methyl radical is \(\\mathrm{CH\_3}\\bullet\).
How A Free Radical Is Formed
Free radicals are often formed when a covalent bond breaks evenly, so that each atom keeps one electron of the shared pair. This even splitting is called homolytic fission, or homolysis.
When a chlorine molecule splits in sunlight: \[\\mathrm{Cl\_2} \\rightarrow \\mathrm{Cl}\\bullet + \\mathrm{Cl}\\bullet\]
Each chlorine atom gets one unpaired electron and becomes a free radical. Fig. 1.8-1 shows the shared pair dividing.
Diagram
Fig. 1.8-1 — Homolytic Fission Of A Chlorine Molecule Left: two chlorine atoms drawn as circles joined by a shared pair of two dots between them, labelled Cl₂ and ‘shared pair of electrons’. A wavy arrow labelled ‘sunlight’ points to the right. Right: the two circles now separated, each carrying one dot on the facing side, each labelled Cl• and ‘one unpaired electron — a free radical’. Below the arrow, the caption line ‘the pair divides evenly, one electron to each atom’.
1.8.2 Existence Of Radicals In The Atmosphere
Atmospheric radicals are formed when stable molecules absorb energy, mainly from sunlight, and their chemical bonds break. This process produces highly reactive species with unpaired electrons.
The most important atmospheric radical is the hydroxyl radical, \(\\mathrm{OH}\\bullet\), often called the detergent of the atmosphere. It plays a vital role in removing air pollutants, including methane, \(\\mathrm{CH\_4}\), carbon monoxide and other harmful gases.
The name is well chosen. A detergent does not remove dirt by pushing it away; it changes it into something that water can carry off, and that is exactly what the hydroxyl radical does to a pollutant molecule.
Other major atmospheric radicals include the nitrate radical, \(\\mathrm{NO\_3}\\bullet\), and the chlorine radical, \(\\mathrm{Cl}\\bullet\).
The nitrate radical is especially active at night. It reacts with harmful gases and pollutants in the air, changing them into forms that can be washed out by rain.
The chlorine radical is released from chlorofluorocarbons, known as CFCs, and from related compounds. It takes part in reactions that destroy the ozone layer, as FC 1.3 traces.
Flowchart
FC 1.3 — How A Chlorine Radical Reaches The Ozone Layer CFC released at ground level → drifts slowly upward, unreactive on the way → reaches the upper atmosphere → absorbs ultraviolet radiation from the Sun → the C–Cl bond breaks by homolytic fission → chlorine radical Cl• set free → reacts with ozone molecules → ozone destroyed and the radical released again to attack further ozone
Thus atmospheric radicals play a crucial role both in cleaning the atmosphere and in influencing important environmental processes. The same unpaired electron that makes one radical useful makes another damaging.
1.8.3 Radicals And Their Effects On Living Organisms
Free radicals are also produced naturally inside the human body. They are formed during normal processes such as breathing, digestion and energy production in cells.
In the body, free radicals are mainly produced during cellular metabolism, when oxygen is used to release energy from food.
Beneficial Effects
In small and controlled amounts, free radicals are useful to the body:
- They help the immune system destroy bacteria and viruses. - They play a role in cell signalling and normal biological processes. - They assist in healing and defence mechanisms.
Harmful Effects
When free radicals are produced in excess, they can damage healthy cells. This condition is called oxidative stress. Excess free radicals can:
- damage cell membranes; - harm proteins and enzymes; - affect DNA and genetic material; - speed up ageing processes.
Such damage may increase the risk of various diseases over time. Notice that the problem is never the radical itself but the balance — too few and the immune system loses a tool, too many and healthy cells are attacked.
1.8.4 Reduction Or Elimination Of Free Radicals
In the atmosphere, restricting the use of chlorofluorocarbons reduces the formation of chlorine radicals, and so protects the ozone layer.
In the body, the defence is chemical. The body protects itself from harmful free radicals with substances called antioxidants, which react with a free radical and quench its unpaired electron before it can damage a cell.
Vitamins such as vitamin C and vitamin E act as antioxidants and help neutralise excess free radicals in the body.
1.9 Chemical Equation
A chemical change is a story with a beginning and an end: substances go in, different substances come out. Written in symbols and formulae, that story becomes a chemical equation.
1.9.1 Word Equation And Chemical Equation
Burning coal in air is a chemical reaction, in which a new substance, carbon dioxide, is formed. The reaction can be represented in two ways.
The first is a word equation:
carbon + oxygen \(\\xrightarrow{\\ \\text{heat}\\ }\) carbon dioxide
A word equation simply represents a change of reactants into products. It is where every equation should start, because it says in plain language what happened.
But a word equation does not give a true picture of the chemical reaction. It says nothing about how many particles of each substance take part, and nothing about the composition of those particles. Hence writing a chemical equation becomes necessary:
\[\\mathrm{C} + \\mathrm{O\_2} \\xrightarrow{\\ \\text{heat}\\ } \\mathrm{CO\_2}\]
1.9.2 Steps In Writing A Chemical Equation
- Write the symbols or the formulae of the reactants on the left hand side, with a (+) sign between them.
- Write the symbols or the formulae of the products on the right hand side, with a (+) sign between them.
- Put the sign of an arrow (\(\\rightarrow\)) in between the reactant side and the product side. Two arrow heads in opposite directions (\(\\rightleftharpoons\)) show that the products formed can recombine to form the reactants.
- Represent the reactants and the products in their molecular forms, because their atomic forms are usually neither stable nor capable of separate existence.
Step (iv) is the one that catches beginners. Oxygen gas must be written \(\\mathrm{O\_2}\) and not \(\\mathrm{O}\), because a lone oxygen atom is not what is present in the air.
sodium + water → sodium hydroxide + hydrogen \[2\\mathrm{Na} + 2\\mathrm{H\_2O} \\rightarrow 2\\mathrm{NaOH} + \\mathrm{H\_2}\]
A chemical equation tells us what substances are involved in a given reaction — the reactants — and what substances are formed as a result of the reaction — the products. Fig. 1.9-1 labels the parts.
Diagram
Fig. 1.9-1 — The Parts Of A Chemical Equation The equation CuSO₄ + 2NaOH → Cu(OH)₂ + Na₂SO₄ printed large across the centre. A horizontal brace under the left-hand side labelled ‘reactants’; a horizontal brace under the right-hand side labelled ‘products’. A callout arrow to the arrow symbol reading ‘read as: gives, or forms’. A callout arrow to the 2 in front of NaOH reading ‘coefficient — counts whole molecules’. A callout arrow to the plus sign reading ‘read as: reacts with, on the left; and, on the right’.
In this equation copper sulphate and sodium hydroxide, the reactants, react to produce copper hydroxide and sodium sulphate, the products.
1.9.3 Types Of Chemical Reactions By Reactants And Products
Chemical reactions may involve:
(i) One reactant and two or more products \[\\mathrm{CaCO\_3} \\rightarrow \\mathrm{CaO} + \\mathrm{CO\_2}\] \[2\\mathrm{Pb(NO\_3)\_2} \\rightarrow 2\\mathrm{PbO} + 4\\mathrm{NO\_2} + \\mathrm{O\_2}\]
(ii) Two reactants and one product \[\\mathrm{NH\_3} + \\mathrm{HCl} \\rightarrow \\mathrm{NH\_4Cl}\] \[\\mathrm{N\_2} + 3\\mathrm{H\_2} \\rightarrow 2\\mathrm{NH\_3}\] \[2\\mathrm{Cu} + \\mathrm{O\_2} \\rightarrow 2\\mathrm{CuO}\]
(iii) Two reactants and two products \[\\mathrm{AgNO\_3} + \\mathrm{NaCl} \\rightarrow \\mathrm{AgCl} + \\mathrm{NaNO\_3}\] \[\\mathrm{Na\_2SO\_4} + \\mathrm{BaCl\_2} \\rightarrow \\mathrm{BaSO\_4} + 2\\mathrm{NaCl}\]
(iv) Two reactants and three or more products \[\\mathrm{Cu} + 2\\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{CuSO\_4} + 2\\mathrm{H\_2O} + \\mathrm{SO\_2}\] \[2\\mathrm{KMnO\_4} + 16\\mathrm{HCl} \\rightarrow 2\\mathrm{KCl} + 2\\mathrm{MnCl\_2} + 8\\mathrm{H\_2O} + 5\\mathrm{Cl\_2}\]
1.9.4 Skeleton Equation And Balanced Equation
A skeleton equation is an equation that represents a chemical change but is unbalanced. In other words, the total numbers of atoms of each element on the two sides are not equal.
\[\\mathrm{KNO\_3} \\rightarrow \\mathrm{KNO\_2} + \\mathrm{O\_2}\]
In this equation the number of oxygen atoms in the reactant, three, is not equal to the number of oxygen atoms in the products formed, which is four.
In \(\\mathrm{CaCO\_3} \\rightarrow \\mathrm{CaO} + \\mathrm{CO\_2}\) the numbers of atoms of calcium, carbon and oxygen on both sides are the same, so the equation is balanced. The skeleton equation above is not, and its balanced form is:
\[2\\mathrm{KNO\_3} \\rightarrow 2\\mathrm{KNO\_2} + \\mathrm{O\_2}\]
Balancing an equation is the method of making the number of atoms of the different elements on both sides of the equation equal.
Why An Equation Must Be Balanced
An equation must be balanced in order to comply with the law of conservation of matter. An unbalanced equation would imply that atoms have been created or destroyed, and no chemical reaction can do that.
Note one thing carefully: while balancing, you may put a number in front of a formula, but you may never change a subscript inside it. Changing \(\\mathrm{H\_2O}\) to \(\\mathrm{H\_2O\_2}\) would balance the oxygen, but it would also change water into a different substance altogether.
1.9.5 Balancing By The Hit And Trial Method
There are two methods of balancing an equation: the hit and trial method, and the partial equation method.
The hit and trial method consists of counting the number of atoms of each element on both sides and trying to equalise them. Take the following steps:
- Count the number of times (frequency) an element occurs on either side.
- The element with the least frequency of occurrence is balanced first.
- When two or more elements have the same frequency, the metallic element is balanced first.
The reason behind step (ii) is practical. An element that appears in only one formula on each side can be fixed without disturbing anything else, while an element scattered across three formulae will have to be adjusted again later.
A Balanced Equation Is Not Always A Real One
Balancing is arithmetic; whether the reaction happens is chemistry. A balanced equation need not represent a real reaction.
\[\\mathrm{Cu} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{CuSO\_4} + \\mathrm{H\_2}\]
is perfectly balanced, but copper is less active than hydrogen, so it cannot displace hydrogen from the acid to produce hydrogen gas. Experiments show that copper reacts with concentrated sulphuric acid to give sulphur dioxide and not hydrogen, so the correct chemical equation is the one you balanced in Example 1.9.
The hit and trial method is very useful for balancing simple chemical equations, but it has some limitations: (i) it takes time to balance complicated equations; (ii) the mechanism of the reaction — the steps by which it has taken place — is not clear.
1.9.6 Balancing By The Partial Equation Method
Balancing can be done more easily by supposing the complex reaction to have taken place in steps. We write the equations for these individual steps and then add them. This is known as balancing by the partial equation method.
Consider the liberation of iodine from potassium iodide by reacting it with hydrogen peroxide, supposed to be completed in the following two steps.
Hydrogen peroxide decomposes to give water and nascent oxygen: \[\\mathrm{H\_2O\_2} \\rightarrow \\mathrm{H\_2O} + \\mathrm{O}\]
The nascent oxygen so produced oxidises potassium iodide in the presence of water to give iodine and potassium hydroxide: \[2\\mathrm{KI} + \\mathrm{H\_2O} + \\mathrm{O} \\rightarrow 2\\mathrm{KOH} + \\mathrm{I\_2}\]
Adding the two equations gives: \[2\\mathrm{KI} + \\mathrm{H\_2O\_2} \\rightarrow 2\\mathrm{KOH} + \\mathrm{I\_2}\]
Water and oxygen occur on the product side in the first step and on the reactant side in the second, so they cancel out. The resultant equation is a balanced chemical equation.
1.9.7 Information Conveyed By A Balanced Equation
Refer to the equation for the neutralisation of sulphuric acid by sodium hydroxide.
sodium hydroxide + sulphuric acid → sodium sulphate + water \[2\\mathrm{NaOH} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{Na\_2SO\_4} + 2\\mathrm{H\_2O}\]
The above equation tells us:
- about the actual result of the chemical change;
- about the reactants involved and the products formed as a result of the reaction;
- about the number of molecules of each substance taking part and formed. Here two molecules of sodium hydroxide and one molecule of sulphuric acid react to give one molecule of sodium sulphate and two molecules of water;
- about the chemical composition of the respective molecules. One molecule of sodium hydroxide contains one atom of sodium, one atom of oxygen and one atom of hydrogen;
- about molecular mass — that 80 parts by weight of sodium hydroxide react with 98 parts by weight of sulphuric acid to produce 142 parts by weight of sodium sulphate and 36 parts by weight of water:
\[2(23+16+1) + (2+32+64) = (46+32+64) + 2(2+16)\] \[80 + 98 = 142 + 36\] \[178 = 178\]
- it also proves the law of conservation of mass, since the total mass of the substances on either side of the equation is the same. According to the above equation, 178 gram of reactants is producing 178 gram of products.
Reading Masses And Volumes From An Equation
On the basis of experiments it is found that 22.4 litres of every gas, at 760 mm pressure and a temperature of \(0^\\circ\\text{C}\) — that is, at S.T.P. — weighs the same as its molar mass, which is its molecular mass expressed in grams. So 22.4 litres of hydrogen at S.T.P. will weigh 2 g, and 22.4 litres of ammonia at S.T.P. will weigh 17 g.
Now read the next equation with its masses written underneath:
\[\\mathrm{CaCO\_3} + 2\\mathrm{HCl} \\rightarrow \\mathrm{CaCl\_2} + \\mathrm{H\_2O} + \\mathrm{CO\_2}\] \[100 \\qquad 73 \\qquad 111 \\qquad 18 \\qquad 44\]
This balanced chemical equation conveys the following information:
- one molecule of calcium carbonate reacts with two molecules of hydrochloric acid to produce one molecule each of calcium chloride, water and carbon dioxide;
- 100 g of calcium carbonate reacts with 73 g of hydrochloric acid to produce 111 g of calcium chloride, 18 g of water and 44 g of carbon dioxide. Note that the masses of the reactants and the products are taken in grams;
- 100 g of calcium carbonate, on treatment with 73 g of hydrochloric acid, will produce 22.4 litres of carbon dioxide at S.T.P.
Statement (c) follows from the note above: 44 g is the molar mass of carbon dioxide, and one molar mass of any gas occupies 22.4 litres at S.T.P.
1.9.8 Limitations Of A Chemical Equation
A chemical equation does not tell us:
- the physical state of the reactants and the products, that is, whether the substances are solid, liquid or gas;
- the time taken for the completion of the reaction;
- whether heat is given out or absorbed during the reaction;
- the respective concentrations of the reactants and the products;
- the rate at which the reaction proceeds;
- whether the reaction is reversible or irreversible;
- whether the reaction is completed or it is not completed.
1.9.9 Making A Chemical Equation More Informative
Most of those gaps can be closed. A chemical equation can be made more informative by writing the following additional information into it.
(a) Physical states. These are shown by the letters \(\\text{(s)}\) for solid, \(\\text{(l)}\) for liquid, \(\\text{(g)}\) or an upward arrow for gas, and \(\\text{(aq)}\) for a solution in water.
Zinc, a solid metal, reacts with hydrochloric acid in aqueous state to produce zinc chloride in aqueous state and the gas hydrogen. \[\\mathrm{Zn}\\text{(s)} + 2\\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{ZnCl\_2}\\text{(aq)} + \\mathrm{H\_2}\\text{(g)}\]
(b) Concentration of acids. This can also be added to the reaction, and it matters, because the same metal and the same acid can give different products at different concentrations.
Zinc reacts with dilute sulphuric acid to produce zinc sulphate and hydrogen. \[\\mathrm{Zn}\\text{(s)} + \\mathrm{H\_2SO\_4}\\text{(dil)} \\rightarrow \\mathrm{ZnSO\_4}\\text{(aq)} + \\mathrm{H\_2}\\text{(g)}\]
Zinc reacts with concentrated sulphuric acid to produce zinc sulphate, sulphur dioxide and water. \[\\mathrm{Zn}\\text{(s)} + 2\\mathrm{H\_2SO\_4}\\text{(conc.)} \\rightarrow \\mathrm{ZnSO\_4}\\text{(aq)} + 2\\mathrm{H\_2O}\\text{(l)} + \\mathrm{SO\_2}\\text{(g)}\]
(c) Heating. Heating during the reaction can be shown by a sign of delta, \(\\Delta\), over the arrow.
Decomposition of potassium chlorate on heating produces potassium chloride and oxygen. \[2\\mathrm{KClO\_3} \\xrightarrow{\\ \\Delta\\ } 2\\mathrm{KCl} + 3\\mathrm{O\_2}\]
(d) A precipitate. Any reaction in which a precipitate is formed can show it by a downward arrow.
When a solution of silver nitrate reacts with dilute hydrochloric acid, a precipitate of silver chloride is formed. \[\\mathrm{AgNO\_3}\\text{(aq)} + \\mathrm{HCl}\\text{(aq)} \\rightarrow \\mathrm{HNO\_3}\\text{(aq)} + \\mathrm{AgCl}\\downarrow\]
(e) Reversibility. A reaction which is reversible can be shown by the sign of reversibility.
The reaction of steam with red hot iron to produce triferric tetraoxide and hydrogen is a reversible reaction. \[3\\mathrm{Fe} + 4\\mathrm{H\_2O} \\rightleftharpoons \\mathrm{Fe\_3O\_4} + 4\\mathrm{H\_2}\]
(f) Heat evolved or absorbed. Chemical reactions proceed with evolution or absorption of heat, and this information is provided by adding the term ‘heat’.
\[\\mathrm{C}\\text{(s)} + \\mathrm{O\_2}\\text{(g)} \\rightarrow \\mathrm{CO\_2}\\text{(g)} + \\text{Heat}\] \[\\mathrm{N\_2}\\text{(g)} + \\mathrm{O\_2}\\text{(g)} \\rightarrow 2\\mathrm{NO}\\text{(g)} - \\text{Heat}\]
(g) Temperature, pressure and catalyst. This information is provided above and below the arrow separating reactants and products.
Nitrogen reacts with hydrogen in the presence of the catalyst iron and the promoter molybdenum, at \(450^\\circ\\text{C}\) and 200 to 900 atmospheres pressure, to produce ammonia, with evolution of heat. \[\\mathrm{N\_2} + 3\\mathrm{H\_2} \\xrightarrow\[\\ 200-900\\ \\text{atm}\\ ]{\\ \\mathrm{Fe} + \\mathrm{Mo},\\ 450^\\circ\\text{C}\\ } 2\\mathrm{NH\_3} + \\text{heat}\]
1.10 Relative Atomic Mass
Atoms, being extremely small, cannot be seen or weighed directly. No balance in any laboratory can hold a single atom. Yet indirect methods of physics have enabled us to know the absolute mass of nearly all kinds of atoms.
The mass of a hydrogen atom is found to be \(1.66 \\times 10^{-24}\\ \\text{g}\), while that of a carbon atom is \(1.9926 \\times 10^{-23}\\ \\text{g}\). These masses are perfectly real, and perfectly useless for everyday work.
As these masses are too small, it is not convenient to use kilograms or grams as the unit. It has therefore been considered appropriate to use the mass of some standard atom as a unit, and then relate the masses of other atoms to it. The resulting masses of atoms are known as relative atomic mass, or atomic weight.
Choosing The Standard
In the beginning, the mass of the hydrogen atom was chosen as the unit, hydrogen being the lightest element, and the masses of other atoms were compared with it.
In 1961, carbon-12 was finally selected as the standard, because its adoption least affected the values of the atomic masses of the various elements on the old standard. One-twelfth of a carbon-12 atom is very nearly the mass of a hydrogen atom, so the numbers chemists were already using hardly changed.
Fig. 1.10-1 shows what that fraction is really saying.
Diagram
Fig. 1.10-1 — Comparing An Atom With The Standard A simple two-pan beam balance in balance. On the left pan, one sphere labelled ‘one atom of the element — oxygen’. On the right pan, sixteen identical small spheres stacked, each labelled ‘1/12 the mass of a carbon-12 atom = 1 u’. Below the balance the caption line ‘one oxygen atom is 16 times as heavy as the standard, so its relative atomic mass is 16’.
The Atomic Mass Unit
Atomic mass is expressed in atomic mass units, written a.m.u. The atomic mass unit is defined as \(\\frac{1}{12}\) the mass of a carbon atom, C-12. The mass of an atom of the carbon-12 isotope was given the atomic mass of 12 units, that is, 12 amu, or simply 12 u.
Thus the mass of a hydrogen atom is 1 amu, and those of oxygen and helium are 16 amu and 4 amu respectively.
Experimentally it is found that one atom of \(\\mathrm{C}^{12}\) has a mass of \(1.9926 \\times 10^{-23}\) grams. On dividing this mass by 12, the atomic mass unit is obtained:
\[1\\ \\text{amu or}\\ 1\\ \\text{u} = 1.6605 \\times 10^{-24}\\ \\text{g}\]
Two Points Worth Remembering
The atomic mass of an element is not its actual mass but its relative mass. The element carbon is used as the reference for expressing the atomic masses of elements.
The atomic mass of an element is the average mass of the atoms in it. That is why some elements are quoted with atomic masses that are not whole numbers — chlorine, for instance, is quoted as 35.5.
1.11 Relative Molecular Mass
An atom was compared with a standard to give its relative atomic mass. A molecule can be compared with the same standard, and the number that results is its relative molecular mass.
In practice we never have to make that comparison ourselves. A molecule is nothing but its atoms joined together, so its mass is simply the sum of their masses.
The relative molecular mass (R.M.M.) is obtained by adding together the relative atomic masses of all the various atoms present in a molecule.
Working It Out For Sulphuric Acid
One molecule of sulphuric acid, \(\\mathrm{H\_2SO\_4}\), has two atoms of hydrogen, one atom of sulphur and four atoms of oxygen. Take each element in turn, as Fig. 1.11-1 sets out.
Mass of 2 atoms of hydrogen is \(1 \\times 2 = 2\) amu Mass of 1 atom of sulphur is 32 amu Mass of 4 atoms of oxygen is \(16 \\times 4 = 64\) amu
So the mass of \(\\mathrm{H\_2SO\_4}\) is \(2 \\times 1 + 32 + 16 \\times 4 = 98\) amu.
Diagram
Fig. 1.11-1 — Building Up A Relative Molecular Mass The formula H₂SO₄ printed across the top. Three columns beneath it, one under each symbol. Column 1 headed ‘H’: ‘2 atoms × 1 = 2’. Column 2 headed ‘S’: ‘1 atom × 32 = 32’. Column 3 headed ‘O’: ‘4 atoms × 16 = 64’. A horizontal rule under the three columns, and beneath it the single line ‘total = 98 amu’, with a label at the side reading ‘relative molecular mass of sulphuric acid’.
Thus the molecular mass of a substance is the sum of the atomic masses of the constituent atoms present in one molecule of that substance.
Reading A Dot In A Formula
Some crystals carry water inside them, and the formula shows it after a dot. In \(\\mathrm{CuSO\_4!\\cdot!5H\_2O}\) the dot means that five molecules of water are attached to every molecule of copper sulphate in the crystal.
That water is part of the substance you weigh out, so it must be counted in the molecular mass. Add \(5 \\times 18\) for the five water molecules, exactly as you would add any other atoms.
Notice how the answer in part (b) keeps its decimal. Copper’s relative atomic mass is 63.5, not 64, and rounding it early would have carried an error right through to the end.
1.12 Percentage Composition
A formula tells you which elements are present and in what ratio of atoms. A farmer buying fertiliser, or a chemist checking a sample, usually wants something else: how much of the mass is the element that matters.
The percentage composition of a compound is the percentage by weight of each element present in it.
\[%\\ \\text{of an element} = \\frac{\\text{total mass of the element in one molecule}}{\\text{gram molecular mass of the compound}} \\times 100\]
The total percentage of all the elements in a compound is 100. That gives you a free check on every calculation you do in this section — add your answers and see whether they come to a hundred.
Look carefully at part (d) of that example. The percentage of water is not the same as the percentage of oxygen, even though the water contains oxygen — the water figure counts its hydrogen as well, and leaves out all the oxygen held in the two sulphate radicals.
1.13 Empirical Formula Of A Compound
A molecular formula states exactly how many atoms of each element are present in one molecule. There is a second, simpler formula that states only the ratio between them.
The empirical formula of hydrogen peroxide, \(\\mathrm{H\_2O\_2}\), is \(\\mathrm{HO}\). It indicates the simplest ratio, 1 : 1, between the hydrogen and oxygen atoms in its molecule, whereas its actual formula is \(\\mathrm{H\_2O\_2}\).
Similarly, the empirical formula of glucose, \(\\mathrm{C\_6H\_{12}O\_6}\), is \(\\mathrm{CH\_2O}\). It indicates that the ratio of carbon, hydrogen and oxygen atoms in a molecule of glucose is 1 : 2 : 1.
Getting from one to the other is the same H.C.F. step you met in the criss-cross method. Divide every subscript in the molecular formula by their highest common factor, and what is left is the empirical formula.
One Empirical Formula, Two Compounds
An empirical formula can be the same for two or more compounds. Glucose, \(\\mathrm{C\_6H\_{12}O\_6}\), and acetic acid, \(\\mathrm{CH\_3COOH}\) or \(\\mathrm{C\_2H\_4O\_2}\), both have the same empirical formula, \(\\mathrm{CH\_2O}\).
This is exactly why the empirical formula cannot replace the molecular formula. Glucose is a sugar and acetic acid is the acid in vinegar; a ratio alone cannot tell them apart. Table 1.14 sets the pairs side by side.
Table 1.14 : Empirical and molecular formulae of some compounds
| Compound | Molecular formula | Empirical formula |
|---|---|---|
| Water | \(\\mathrm{H\_2O}\) | \(\\mathrm{H\_2O}\) |
| Hydrogen peroxide | \(\\mathrm{H\_2O\_2}\) | \(\\mathrm{HO}\) |
| Glucose | \(\\mathrm{C\_6H\_{12}O\_6}\) | \(\\mathrm{CH\_2O}\) |
| Acetic acid | \(\\mathrm{C\_2H\_4O\_2}\) | \(\\mathrm{CH\_2O}\) |
Empirical Formula Mass
The empirical formula mass is the sum of the atomic masses of the various elements present in the empirical formula.
Thus, for hydrogen peroxide, \(\\mathrm{H\_2O\_2}\), the empirical formula is \(\\mathrm{HO}\) and its empirical formula mass is \(1 + 16 = 17\).
Compare that with the relative molecular mass of hydrogen peroxide, which is 34. The molecular mass is exactly twice the empirical formula mass, just as the molecular formula holds exactly twice the atoms of the empirical formula.
1.14 Summary
- A molecule is the smallest particle of a substance capable of independent existence. In a pure substance, every molecule is alike, so it has a definite composition and definite properties.
- An element is a simple and pure form of matter which cannot be decomposed into simpler substances. At present 118 elements are known.
- An atom is the smallest particle of an element which may or may not exist independently but can take part in a chemical reaction.
- The number of atoms in a molecule of an element is called its atomicity. Elements may be monoatomic, diatomic, triatomic, tetratomic or octatomic.
- A compound is a substance formed by the chemical combination of two or more elements in a fixed proportion. Its properties are entirely different from those of its constituent elements.
- A mixture is an impure substance formed by mixing pure substances in any proportion, with no chemical change; each constituent keeps its own properties.
- A symbol is the short form standing for the atom of a specific element, written with the first letter in capital. Some symbols come from Latin or German names, such as \(\\mathrm{Fe}\) from ferrum and \(\\mathrm{W}\) from Wolfram.
- A symbol represents the name of the element, one atom of it, and a definite mass of it equal to its atomic mass expressed in grams.
- Valency is the combining capacity of an element — the number of hydrogen atoms, or double the number of oxygen atoms, that combine with or displace one atom of that element.
- By the modern definition, valency is the number of electrons an atom can lose, gain or share during a chemical reaction.
- An ion is a positively or negatively charged atom or group of atoms formed by the loss or gain of electrons. Positive ions are cations and negative ions are anions.
- Atoms lose, gain or share electrons to obtain eight electrons in the outermost shell, hydrogen being stable with two.
- Elements are monovalent, divalent, trivalent or tetravalent according to their valency. Elements with one, two or three valence electrons are usually metals; those with five, six or seven are normally non-metals.
- Certain elements show variable valency, because electrons can also be lost from the penultimate shell. The valency is written as a Roman numeral in brackets, called stock notation, in place of the older -ous and -ic endings.
- A molecular formula is the symbolic representation of a molecule. A subscript counts atoms inside one molecule; a number written in front counts whole molecules.
- A molecular formula represents the molecule and its molecular mass, the number of atoms of each element in it, and the ratio of the masses of those elements.
- A radical is an atom or group of atoms that behaves as a single unit with a positive or negative charge. Basic radicals are cations left when a base loses \(\\mathrm{OH^{-}}\); acidic radicals are anions left when an acid loses \(\\mathrm{H^{+}}\).
- A radical made of a single atom is a simple radical; one made of a group of atoms is a compound or complex radical. The charge on a radical is usually its valency.
- Acidic and basic radicals combine to form salts, and a salt dissolved in water splits again into its radicals.
- Formulae are written by the criss-cross method: basic radical first, valencies written above, divided by their H.C.F. if any, then interchanged and written as subscripts, with brackets where a group is taken more than once.
- Compounds are named by identifying the basic and acidic radicals and writing them in the same order as the formula. A metal with a non-metal takes the suffix -ide.
- Where oxygen atoms differ, the names run hypo—ite, —ite, —ate and per—ate, as in potassium hypochlorite, chlorite, chlorate and perchlorate.
- Binary acids take the prefix hydro and the suffix -ic; oxyacids are named from the second element, with -ic for more oxygen and -ous for less. Some compounds keep trivial names, such as ammonia and water.
- The valency of an element can be found from a formula by interchanging the subscripts as superscripts and multiplying by the known valency of oxygen or hydrogen.
- A free radical is an atom, molecule or ion containing at least one unpaired valence electron, shown by a dot. It is formed by homolytic fission, as when \(\\mathrm{Cl\_2}\) splits in sunlight into two \(\\mathrm{Cl}\\bullet\).
- Atmospheric radicals form when sunlight breaks bonds in stable molecules. The hydroxyl radical removes pollutants, the nitrate radical acts mainly at night, and the chlorine radical from CFCs destroys ozone.
- Free radicals in the body arise during metabolism. In small amounts they aid immunity, signalling and healing; in excess they cause oxidative stress, damaging membranes, proteins and DNA. Antioxidants such as vitamins C and E neutralise them.
- A chemical equation is the symbolic representation of a chemical reaction. A word equation is written first, then the symbols and formulae, reactants on the left and products on the right.
- A skeleton equation is unbalanced. In a balanced equation the number of atoms of each element is the same on both sides, as required by the law of conservation of matter — matter is neither created nor destroyed in a chemical reaction.
- Equations are balanced by the hit and trial method, balancing first the element of least frequency and, where frequencies are equal, the metallic element. While balancing, coefficients may be changed but subscripts may not.
- A balanced equation need not represent a real reaction; what actually happens must be settled by experiment.
- A balanced equation tells us the reactants and products, the number of molecules of each, the composition of those molecules, and the masses that react and are formed, as in \(2\\mathrm{NaOH} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{Na\_2SO\_4} + 2\\mathrm{H\_2O}\), where \(80 + 98 = 142 + 36\).
- It does not tell us the physical states, the time taken, the heat change, the concentrations, the rate, the reversibility or whether the reaction is complete.
- These can be added: state symbols \(\\text{(s)}\), \(\\text{(l)}\), \(\\text{(g)}\), \(\\text{(aq)}\); dil or conc for acids; \(\\Delta\) over the arrow for heating; \(\\downarrow\) for a precipitate; \(\\rightleftharpoons\) for a reversible reaction; heat written in; and the catalyst, temperature and pressure on the arrow.
- The relative atomic mass of an element is the number of times one atom of it is heavier than \(\\frac{1}{12}\) the mass of an atom of carbon-12, adopted as the standard in 1961.
- One atomic mass unit is \(\\frac{1}{12}\) the mass of a carbon-12 atom, equal to \(1.6605 \\times 10^{-24}\\ \\text{g}\). Atomic mass is a relative, average value, not an actual mass.
- The relative molecular mass of a substance is the sum of the relative atomic masses of all the atoms in one molecule, so \(\\mathrm{H\_2SO\_4}\) has a relative molecular mass of 98.
- Percentage composition is the percentage by mass of each element in a compound: \[%\\ \\text{of an element} = \\frac{\\text{total mass of the element in one molecule}}{\\text{gram molecular mass of the compound}} \\times 100\] and all the percentages add up to 100.
- The empirical formula gives the simplest whole number ratio of atoms in a compound, such as \(\\mathrm{CH\_2O}\) for glucose. Two different compounds may share one empirical formula.
- The empirical formula mass is the sum of the atomic masses of the elements present in the empirical formula; for \(\\mathrm{HO}\) it is 17.
Flowchart
FC 1.4 — The Language Of Chemistry, End To End Element → its symbol (Berzelius system, one or two letters) → its valency, read from the valence electrons → formula of the compound, built by the criss-cross method → name of the compound, read from its two radicals → word equation for the change → skeleton equation in symbols and formulae → balanced equation, obeying the law of conservation of matter → masses and percentages calculated from relative atomic and molecular masses
1.15 Key Terms
- Acidic radical — a negatively charged radical, left behind when an acid loses its hydrogen ion.
- Anion — a negatively charged ion, formed when an atom or group of atoms gains electrons.
- Antioxidant — a substance, such as vitamin C or vitamin E, that neutralises excess free radicals in the body.
- Atom — the smallest particle of an element, which may or may not exist independently but can take part in a chemical reaction.
- Atomic mass unit (amu or u) — one-twelfth of the mass of an atom of carbon-12, equal to \(1.6605 \\times 10^{-24}\\ \\text{g}\).
- Atomicity — the number of atoms present in one molecule of an element.
- Balanced equation — an equation in which the number of atoms of each element is the same on both sides.
- Basic radical — a positively charged radical, left behind when a base loses its hydroxide ion.
- Cation — a positively charged ion, formed when an atom or group of atoms loses electrons.
- Chemical equation — the symbolic representation of a chemical reaction using symbols and formulae.
- Chemical formula — another name for the molecular formula of a substance.
- Compound — a substance formed by the chemical combination of two or more elements in a fixed proportion.
- Compound radical — a radical made of a group of two or more atoms, such as sulphate; also called a complex or polyatomic radical.
- Criss-cross method — the method of writing a formula by interchanging the valencies of the two radicals and writing them as subscripts.
- Divalent — having a valency of two.
- Element — a simple and pure form of matter which cannot be decomposed into simpler substances.
- Empirical formula — the formula giving the simplest whole number ratio of the atoms of the elements in a compound.
- Empirical formula mass — the sum of the atomic masses of the elements present in the empirical formula.
- Free radical — an atom, molecule or ion containing at least one unpaired valence electron, shown by a dot.
- Homolytic fission — the even breaking of a covalent bond, in which each atom keeps one electron of the shared pair.
- Ion — a positively or negatively charged atom, or group of atoms, formed by the loss or gain of electrons.
- Mixture — an impure substance formed by mixing two or more pure substances in any proportion, without chemical change.
- Molecular formula — the symbolic representation of a molecule, showing the symbols and numbers of the atoms present in it.
- Molecule — the smallest particle of a substance which is capable of independent existence.
- Monovalent — having a valency of one.
- Oxidative stress — the damage caused to healthy cells when free radicals are produced in excess in the body.
- Percentage composition — the percentage by mass of each element present in a compound.
- Product — a substance formed as a result of a chemical reaction, written on the right of the arrow.
- Pure substance — a substance in which every molecule is alike, having a definite composition and definite properties.
- Radical — an atom or group of atoms that behaves as a single unit and carries a positive or negative charge.
- Reactant — a substance that takes part in a chemical reaction, written on the left of the arrow.
- Relative atomic mass — the number of times one atom of an element is heavier than one-twelfth the mass of an atom of carbon-12.
- Relative molecular mass — the sum of the relative atomic masses of all the atoms present in one molecule of a substance.
- Simple radical — a radical consisting of a single atom, such as sodium or chloride.
- Skeleton equation — an equation that represents a chemical change but is not yet balanced.
- Stock notation — the writing of an element’s valency as a Roman numeral in brackets after its name.
- Symbol — the short form standing for the atom of a particular element.
- Tetravalent — having a valency of four.
- Trivalent — having a valency of three.
- Trivial name — a widely accepted common name that follows no systematic rule, such as ammonia or water.
- Valence electrons — the electrons present in the outermost shell of an atom.
- Valence shell — the outermost shell of an atom, from which electrons are lost or gained.
- Valency — the combining capacity of an element or radical, measured by the number of electrons lost, gained or shared.
- Variable valency — the property of certain elements of showing more than one valency, because electrons may also be lost from the penultimate shell.
1.16 Points To Ponder
Symbols were invented so that chemistry would read the same everywhere in the world. Yet a symbol such as \(\\mathrm{Fe}\) comes from a Latin word most students have never used. Is a symbol truly international if it hides the name of the element, or does hiding the name make it more international rather than less?
An element with four valence electrons could in principle lose four or gain four, and its valency would be 4 either way. What might decide which of the two it actually does, and what would you need to know about the atom to predict it?
Glucose and acetic acid share the empirical formula \(\\mathrm{CH\_2O}\), yet one is a sugar and the other is the acid in vinegar. What does this tell you about how much of a substance’s character is fixed by the ratio of its elements, and how much by something else?
Balancing an equation makes the atoms come out equal, but it can also be done for a reaction that never happens. If arithmetic cannot tell you whether a reaction is real, what can — and what would that mean for how a chemist should treat an equation found in a book?
The free radicals your body makes are useful in small amounts and damaging in large ones. Can a substance be called harmful or helpful on its own, or does the answer always depend on the quantity? Try to think of a second example from this chapter that behaves the same way.
Chemists changed their standard for atomic masses in 1961, choosing carbon-12 partly because it disturbed the existing numbers least. Should convenience carry weight in a scientific decision, or should a standard be chosen only on how well it is defined?
1.17 Objective Questions
- Choose the correct option: [8]
- The atomicity of a molecule of ozone is:
- one
- two
- three
- four
- An atom whose outermost shell contains six electrons will most probably:
- lose six electrons
- gain two electrons
- lose two electrons
- gain six electrons
- The correct formula of aluminium sulphide is:
- \(\\mathrm{AlS}\)
- \(\\mathrm{Al\_3S\_2}\)
- \(\\mathrm{Al\_2S\_3}\)
- \(\\mathrm{AlS\_3}\)
- The valency of manganese in \(\\mathrm{KMnO\_4}\) is:
- two
- four
- six
- seven
- The symbol \(\\mathrm{Sb}\) stands for the element whose Latin name is:
- stannum
- stibium
- natrium
- plumbum
- The number of oxygen atoms in one molecule of \(\\mathrm{Ca\_3(PO\_4)\_2}\) is:
- four
- six
- eight
- twelve
- A radical that carries no charge but has one unpaired electron is:
- \(\\mathrm{SO\_4^{2-}}\)
- \(\\mathrm{NH\_4^{+}}\)
- \(\\mathrm{OH}\\bullet\)
- \(\\mathrm{OH^{-}}\)
- The empirical formula of \(\\mathrm{N\_2O\_4}\) is:
- \(\\mathrm{NO}\)
- \(\\mathrm{NO\_2}\)
- \(\\mathrm{N\_2O}\)
- \(\\mathrm{N\_2O\_4}\)
- The atomicity of a molecule of ozone is:
- Choose the answer based on the codes given below: [4]
- Both A and R are true and R is the correct explanation of A.
- Both A and R are true but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
- Assertion (A): The formula of calcium oxide is written \(\\mathrm{CaO}\) and not \(\\mathrm{Ca\_2O\_2}\). Reason (R): The valencies must be divided by their highest common factor before they are interchanged.
- Assertion (A): A noble gas has a valency of zero. Reason (R): Its outermost shell already contains eight electrons, except in helium.
- Assertion (A): Copper does not liberate hydrogen from dilute sulphuric acid. Reason (R): The equation \(\\mathrm{Cu} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{CuSO\_4} + \\mathrm{H\_2}\) is not balanced.
- Assertion (A): Glucose and acetic acid have the same empirical formula. Reason (R): Two compounds with the same ratio of atoms must have the same properties.
- Match the substance in Column A with the correct entry in Column B. [5]
| Column A | Column B |
|---|---|
| a. \(\\mathrm{Hg}\) | i. tetratomic molecule |
| b. \(\\mathrm{P\_4}\) | ii. compound radical |
| c. \(\\mathrm{Cl}\\bullet\) | iii. hydrargyrum |
| d. \(\\mathrm{CO\_3^{2-}}\) | iv. empirical formula mass 17 |
| e. \(\\mathrm{HO}\) | v. free radical |
- Complete each statement: [5]
- The number of atoms present in one molecule of an element is called its …………… .
- An atom that has lost electrons carries a …………… charge and is called a …………… .
- In stock notation, the valency of the element is written as a …………… in brackets.
- A radical made of a single atom is called a …………… radical.
- The relative atomic mass of an element is compared with one-twelfth the mass of an atom of …………… .
- State whether each statement is true or false: [4]
- A subscript may be changed while balancing a chemical equation.
- Every balanced equation represents a reaction that actually takes place.
- The percentage of water in a hydrated salt is the same as its percentage of oxygen.
- Relative atomic mass is expressed in grams.
- Name: [5]
- the radical formed when sulphuric acid loses both its hydrogen ions
- an element that shows the valencies 2, 4 and 6
- the atmospheric radical known as the detergent of the atmosphere
- the compound whose trivial name is ammonia
- two vitamins that act as antioxidants in the human body
- Give one word or phrase for: [4]
- an equation that represents a chemical change but is not balanced
- the even breaking of a covalent bond so that each atom keeps one electron
- the water held inside a crystal in a fixed proportion
- the formula giving the simplest whole number ratio of atoms in a compound
- Pick the odd one out and give a reason for your choice: [3]
- helium, neon, oxygen, argon
- \(\\mathrm{NO\_3^{-}}\), \(\\mathrm{SO\_4^{2-}}\), \(\\mathrm{Cl^{-}}\), \(\\mathrm{CO\_3^{2-}}\)
- sodium, potassium, iron, lithium
1.18 Check Your Concepts
Define atomicity and give one example each of a diatomic and an octatomic molecule. [2]
State what the symbol \(\\mathrm{Mg}\) conveys, giving all three meanings. [2]
Give a reason why the symbol for cobalt must be written \(\\mathrm{Co}\) and never \(\\mathrm{CO}\). [2]
Explain, with the help of electron arrangements, why chlorine forms an anion while sodium forms a cation. [3]
Give a reason for each of the following: [3]
- hydrogen was chosen as the original standard for valency
- the number of oxygen atoms is doubled when valency is calculated from an oxide
- a chemical equation must be balanced
What would happen if, while balancing \(\\mathrm{H\_2} + \\mathrm{O\_2} \\rightarrow \\mathrm{H\_2O}\), a student changed \(\\mathrm{H\_2O}\) to \(\\mathrm{H\_2O\_2}\) instead of adding coefficients? Explain why this is not allowed. [3]
Differentiate between the following, stating the basis of difference in each case: [4]
- a compound and a mixture
- a simple radical and a compound radical
Explain why an element such as iron shows variable valency, while sodium does not. Illustrate your answer with one compound of each. [4]
Free radicals are produced in the human body during normal life processes. Explain how the same particles can be both useful and harmful, and state how the body limits the harm. [5]
1.19 Equations, Conversions & Identification
Balance the following equations: [4]
- \(\\mathrm{Al} + \\mathrm{Fe\_2O\_3} \\rightarrow \\mathrm{Al\_2O\_3} + \\mathrm{Fe}\)
- \(\\mathrm{Cu(NO\_3)\_2} \\rightarrow \\mathrm{CuO} + \\mathrm{NO\_2} + \\mathrm{O\_2}\)
- \(\\mathrm{Na\_2CO\_3} + \\mathrm{HCl} \\rightarrow \\mathrm{NaCl} + \\mathrm{H\_2O} + \\mathrm{CO\_2}\)
- \(\\mathrm{Fe} + \\mathrm{Cl\_2} \\rightarrow \\mathrm{FeCl\_3}\)
Write the balanced chemical equation for each of the following word equations: [4]
- zinc carbonate → zinc oxide + carbon dioxide
- aluminium + hydrochloric acid → aluminium chloride + hydrogen
- calcium hydroxide + carbon dioxide → calcium carbonate + water
- iron(III) hydroxide → iron(III) oxide + water
Complete and balance the following equations: [3]
- \(\\mathrm{Zn} + \\mathrm{H\_2SO\_4}\\text{(dil)} \\rightarrow \\ldots + \\ldots\)
- \(\\mathrm{NH\_3} + \\mathrm{HCl} \\rightarrow \\ldots\)
- \(\\mathrm{CaCO\_3} \\xrightarrow{\\ \\Delta\\ } \\ldots + \\ldots\)
Rewrite each equation so that it also carries the information given in brackets: [3]
- \(\\mathrm{Zn} + \\mathrm{HCl} \\rightarrow \\mathrm{ZnCl\_2} + \\mathrm{H\_2}\) [state symbols]
- \(\\mathrm{KClO\_3} \\rightarrow \\mathrm{KCl} + \\mathrm{O\_2}\) [balanced, and showing that heat is supplied]
- \(\\mathrm{AgNO\_3} + \\mathrm{HCl} \\rightarrow \\mathrm{HNO\_3} + \\mathrm{AgCl}\) [showing that a precipitate is formed]
Identify the substance in each case: [3]
- a radical of valency 1 that contains one nitrogen atom and four hydrogen atoms
- an element whose atoms have the arrangement 2, 8, 3 and which forms a trivalent cation
- a compound of relative molecular mass 18 whose empirical formula is the same as its molecular formula
Refer to Fig. 1.3-2, which shows the formation of two ions. State which particle is gaining an electron and which is losing one, name the ion formed in each case, and give the electron arrangement of each ion. [3]
Write the chemical formula of each of the following, showing the valencies you used: [4]
- ammonium dichromate
- copper(I) sulphide
- barium phosphate
- sodium thiosulphate
1.20 Practice With Numericals
[Relative atomic masses: \(\\mathrm{H} = 1\), \(\\mathrm{C} = 12\), \(\\mathrm{N} = 14\), \(\\mathrm{O} = 16\), \(\\mathrm{Na} = 23\), \(\\mathrm{Mg} = 24\), \(\\mathrm{Al} = 27\), \(\\mathrm{S} = 32\), \(\\mathrm{Cl} = 35.5\), \(\\mathrm{K} = 39\), \(\\mathrm{Ca} = 40\), \(\\mathrm{Fe} = 56\), \(\\mathrm{Cu} = 63.5\)]
Calculate the relative molecular mass of: [4]
- \(\\mathrm{Al\_2(SO\_4)\_3}\)
- \(\\mathrm{K\_2Cr\_2O\_7}\) [take \(\\mathrm{Cr} = 52\)]
- \(\\mathrm{Mg(NO\_3)\_2}\)
- \(\\mathrm{CH\_3COONa}\)
Calculate the percentage of calcium in calcium carbonate, and hence the mass of calcium present in 250 g of the compound. [3]
Ammonium sulphate is used as a nitrogenous fertiliser. Calculate the percentage of nitrogen in it, and state whether a sack of ammonium sulphate or a sack of urea of the same mass supplies more nitrogen. [3]
Calculate the percentage of water of crystallisation in \(\\mathrm{MgSO\_4!\\cdot!7H\_2O}\). [3]
Given that 1 amu \(= 1.6605 \\times 10^{-24}\\ \\text{g}\), calculate the mass in grams of one atom of sulphur, and state how many times heavier it is than one atom of carbon-12. [3]
Write the empirical formula of \(\\mathrm{C\_2H\_4}\) and of \(\\mathrm{H\_2O\_2}\), and calculate the empirical formula mass of each. [3]
From the balanced equation \(\\mathrm{CaCO\_3} \\rightarrow \\mathrm{CaO} + \\mathrm{CO\_2}\), calculate the mass of calcium oxide and the mass of carbon dioxide obtained by heating 50 g of calcium carbonate. Show that your two answers together account for the whole 50 g. [3]