Chapter 2: Chemical Bonding
In 1902, a young American chemist named Gilbert Newton Lewis was preparing a lecture for his students. He had a puzzle to explain. Why do some elements combine greedily with almost anything, while a handful of gases refuse to combine at all? On a sheet of notepaper he drew atoms as tiny cubes, with electrons sitting at the corners. An atom with all eight corners filled, he reasoned, would be satisfied. An atom with empty corners would go looking for company.
The sketch stayed in his drawer for years. Then, in 1916, Lewis published the idea that changed the way chemists think. Two atoms, he suggested, need not hand electrons over at all. They can hold a pair of electrons between them, and that shared pair belongs to both. With that single thought, the bond stopped being a mysterious hook drawn between two letters and became something made of electrons.
Lewis’s dots are still on the page today. When you draw a chlorine atom with seven dots around it, or an ammonia molecule with one pair left over, you are using a notation over a century old — and you are doing exactly what Lewis did, counting outer electrons to work out what an atom will do next.
Diagram
Fig. 2.0-1 — Two Ways Of Completing An Outer Shell Two panels side by side, each showing two atoms meeting. Left panel labelled ‘transfer’: a metal atom with one outer electron and a non-metal atom with seven; an arrow carries the single electron across; result shown as two separate charged particles pulled together by attraction. Right panel labelled ‘sharing’: two identical non-metal atoms overlapping slightly, with one pair of dots sitting in the overlap and belonging to both; result shown as a single joined molecule. Label the outer shells, the transferred electron, and the shared pair.
2.1 Why Atoms Combine
Look around a chemistry laboratory and you will notice something odd. Sodium has to be kept under oil, because it attacks the moisture in the air. Iron rusts quietly over months. Yet the helium sealed inside a balloon does nothing at all, however long you keep it.
Almost every element you will meet combines with something. A very small family of gases does not. If we can work out what makes that family different, we will have found the reason all the other elements combine — and that reason turns out to lie in the electrons of the outermost shell.
The electrons in the outermost shell of an atom are called its valence electrons, and they are the only electrons that take part in chemical change. The inner electrons are held tightly and stay where they are. So the whole of this chapter is really about what happens to the outermost few electrons when two atoms meet.
2.1.1 The Stability Of The Noble Gases
Helium, neon, argon, krypton, xenon and radon are the noble gases. They were once called the inert gases, because they were thought to take part in no reaction whatever. Their electron arrangements are set out in Table 2.1.
Table 2.1 Electronic configurations of the noble gases
| Noble gas | Atomic number | K | L | M | N | O | P | Valence electrons |
|---|---|---|---|---|---|---|---|---|
| He | 2 | 2 | 2 | |||||
| Ne | 10 | 2 | 8 | 8 | ||||
| Ar | 18 | 2 | 8 | 8 | 8 | |||
| Kr | 36 | 2 | 8 | 18 | 8 | 8 | ||
| Xe | 54 | 2 | 8 | 18 | 18 | 8 | 8 | |
| Rn | 86 | 2 | 8 | 18 | 32 | 18 | 8 | 8 |
Read down the last column and one pattern jumps out. Every one of these gases carries eight electrons in its outermost shell, except helium, which carries two.
Two electrons in the outermost shell is called a duplet, and eight electrons in the outermost shell is called an octet. Helium has a duplet because its only shell, the K shell, is full at two electrons. Every other noble gas has an octet. Figure 2.1-1 shows the shells of three of them.
Diagram
Fig. 2.1-1 — Complete Outermost Shells In The Noble Gases Three separate shell diagrams drawn side by side, each with a labelled nucleus and circular shells carrying dots: helium (2), neon (2, 8) and argon (2, 8, 8). Mark the outermost shell of each with a bracket labelled ‘complete outermost shell’; label helium’s as a duplet and neon’s and argon’s as octets.
Elements with a complete outermost shell do not react. Their atoms have no reason to lose an electron, no room to gain one comfortably, and nothing to gain by sharing. That is exactly what chemical stability means at this level.
Now turn the observation around. If a complete outermost shell makes an atom unreactive, then an incomplete outermost shell must be what makes every other atom reactive. An atom of sodium, with one electron in its outermost shell, and an atom of chlorine, with seven, are both one step away from the arrangement of a noble gas.
2.1.2 The Chemical Bond
When electrons are redistributed between two atoms, a force of attraction springs up between those atoms and holds them together. That force is the chemical bond.
The bond is not a stick or a hook, whatever the lines in a structural formula may suggest. It is electrical attraction — between charges, or between two nuclei and a pair of electrons lying between them.
2.1.3 Three Ways Of Attaining Stability
Atoms can redistribute their outermost electrons in three ways, and each way gives a different kind of bond. The whole of this chapter is built on these three.
- One or more electrons are transferred completely from one atom to the other. The bond formed is an electrovalent bond, also called an ionic bond.
- One, two or three pairs of electrons are shared between two atoms, each atom contributing half of every shared pair. The bond formed is a covalent bond, also called a molecular bond.
- A shared pair is contributed entirely by one of the two atoms. The bond formed is a coordinate bond, also called a dative bond.
Notice that all three end in the same place — every atom involved finishes with a complete outermost shell. What differs is the route taken. Which route two atoms take depends on the kind of atoms they are, and that is the question the next section opens with.
2.2 The Electrovalent (Or Ionic) Bond
Take an atom of a metal with one, two or three electrons in its outermost shell. It is only one, two or three electrons away from the arrangement of the noble gas below it in the table, and it can reach that arrangement by giving those electrons away.
Now take an atom of a non-metal with seven, six or five electrons in its outermost shell. It is one, two or three electrons short of a complete octet, and it can reach the arrangement of the next noble gas by accepting exactly that many.
Put the two together and the transaction completes itself. The metal atom hands over what it needs to lose, the non-metal atom accepts what it needs to gain, and afterwards both have the electron arrangement of the nearest noble gas. This complete handing over of electrons is what builds an electrovalent compound.
2.2.1 Cations, Anions And Electrovalency
Before the transfer, both atoms are electrically neutral, because each has as many protons in its nucleus as electrons around it. After the transfer, that balance is broken on both sides.
A metal atom that has lost electrons now has more protons than electrons, so it carries a positive charge. Such a charged particle is called a cation. A non-metal atom that has gained electrons now has more electrons than protons, so it carries a negative charge, and is called an anion.
An element whose atoms readily lose electrons to form positive ions is called an electropositive element. Sodium is one:
\[\mathrm{Na} - e^{-} \rightarrow \mathrm{Na^{+}}\quad\text{(cation)}\]
An element whose atoms readily accept electrons to form negative ions is called an electronegative element. Chlorine is one:
\[\mathrm{Cl} + e^{-} \rightarrow \mathrm{Cl^{-}}\quad\text{(anion)}\]
Look at what these two half changes have in common: the electron leaving the first line is the electron arriving in the second. Neither can happen on its own.
The cation and the anion now carry opposite charges, and opposite charges attract. This electrostatic attraction between the cation and the anion is the bond itself, and it is called an electrovalent bond or an ionic bond.
So the electrovalency of sodium is 1, because a sodium atom loses one electron. The electrovalency of magnesium is 2, and that of chlorine is 1.
2.2.2 Conditions For The Formation Of An Electrovalent Bond
An ionic bond forms only when the transfer of electrons is easy for both atoms. Three properties of the combining atoms decide that.
- Low ionisation potential. The ionisation potential of an atom is the energy needed to pull an electron away from it. If that energy is low for one of the atoms, it parts with its electron readily, and a cation forms easily.
- High electron affinity. The electron affinity of an atom is the energy released when it accepts an electron. A high value means the atom holds the added electron firmly, so an anion forms easily.
- Large electronegativity difference. Electronegativity is the tendency of an atom to pull shared electrons towards itself. If the two atoms differ widely in electronegativity, one pulls so much harder than the other that the electron is not merely pulled but taken. The greater the difference in electronegativity, the more ionic the resulting compound.
Read together, these three conditions describe a metal meeting a non-metal. The metals of groups 1, 2 and 13 hold their valence electrons loosely and tend to lose them. The non-metals of groups 15, 16 and 17 attract electrons strongly and tend to gain them. When atoms from these two sets meet, ionic bonds form.
Group 1 elements are the most electropositive of all, which is another way of saying that they are the most strongly metallic, and their metallic character increases down the group. Group 17 elements are the most electronegative, fluorine being the most electronegative element of all. Caesium sits at one extreme and fluorine at the other, so caesium fluoride, \(\mathrm{CsF}\), is the most ionic compound known.
Why Ionic Compounds Are Stable
An ionic solid is packed with ions of both kinds, and not all of the forces in it are attractive. Ions carrying like charges repel one another. But every ion is surrounded most closely by ions of the opposite charge, so the attractive forces between unlike charges far outweigh the repulsions between like ones. The net result is a strongly held, stable arrangement.
Some familiar electrovalent compounds are listed below.
| Formula | Name |
|---|---|
| \(\mathrm{NaCl}\) | Sodium chloride |
| \(\mathrm{MgCl_2}\) | Magnesium chloride |
| \(\mathrm{CaO}\) | Calcium oxide |
| \(\mathrm{KBr}\) | Potassium bromide |
| \(\mathrm{CaCl_2}\) | Calcium chloride |
2.2.3 Electron Dot (Lewis) Symbols
Drawing complete shells for every atom in a reaction quickly becomes tedious, and the inner shells never change anyway. Chemists therefore use a shorthand that shows only what matters.
In an electron dot symbol, also called a Lewis symbol, the symbol of the element stands for the nucleus together with all the inner shells, and dots placed around it stand for the valence electrons alone. Electrons that are already paired in the outermost shell are drawn as a pair of dots; an unpaired electron is drawn as a single dot. Some of these symbols are shown in Fig. 2.2-1.
Diagram
Fig. 2.2-1 — Electron Dot Symbols Of Some Atoms A row of five electron dot symbols, each with the element symbol at the centre and the correct number of valence electrons around it: hydrogen with one dot; magnesium with two dots; oxygen with six (two pairs and two single dots); chlorine with seven (three pairs and one single dot); neon with eight (four pairs). Label the row ‘valence electrons only — inner shells are not shown’.
When two different atoms appear in the same diagram, their electrons are drawn with different marks — dots for one atom and crosses for the other. This is only a bookkeeping device that lets you follow where each electron came from. Once a bond has formed, the electrons are identical and no experiment could tell them apart.
2.2.4 Structures Of Some Electrovalent Compounds
We can now follow three ionic compounds forming, electron by electron.
2.2.4.1 Sodium Chloride
The electron arrangement of a sodium atom is 2, 8, 1. It has one electron more than neon, whose arrangement is 2, 8. Losing that single electron is a far cheaper route to a complete shell than gaining seven, so a sodium atom gives it up.
\[\mathrm{Na}\ (2,8,1) - 1e^{-} \rightarrow \mathrm{Na^{+}}\ (2,8)\]
The atom is no longer neutral. Eleven protons remain in the nucleus but only ten electrons revolve around it, so the particle carries a net charge of \(+1\). It is now the sodium ion, \(\mathrm{Na^{+}}\), and its electron arrangement is that of neon.
This is worth pausing over. Sodium metal and the sodium ion are not two names for the same thing; they are chemically different substances, as Table 2.2 shows.
Table 2.2 Comparison of the properties of the sodium atom and the sodium ion
| Property | Sodium atom (Na) | Sodium cation (Na⁺) |
|---|---|---|
| Colour | Silvery white | Colourless |
| Toxicity | Poisonous | Non-poisonous |
| Chemical action | Very active | Inactive |
| Valence shell | Incomplete outermost shell | Complete outermost shell |
| Electrical state | Neutral | Positively charged |
| Existence | Combined state | Independent existence |
The chlorine atom is in the opposite position. Its electron arrangement is 2, 8, 7 — one electron short of argon, 2, 8, 8 — so it accepts one electron to complete its octet.
\[\mathrm{Cl}\ (2,8,7) + 1e^{-} \rightarrow \mathrm{Cl^{-}}\ (2,8,8)\]
A chlorine atom holds 17 protons and 17 electrons and is neutral. After accepting an electron it has one electron more than it has protons, so it carries a charge of \(-1\) and is called the chloride ion. Its properties, too, are quite unlike those of the atom it came from, as set out in Table 2.3.
Table 2.3 Comparison of the properties of the chlorine atom and the chloride ion
| Property | Chlorine atom (Cl) | Chloride anion (Cl⁻) |
|---|---|---|
| Colour | Yellowish green (as Cl₂ gas) | Colourless |
| Toxicity | Poisonous | Non-poisonous |
| Odour | Suffocating | Odourless |
| Chemical action | Very active | Inactive |
| Valence shell | Incomplete outermost shell | Complete outermost shell |
| Electrical state | Neutral | Negatively charged |
| Existence | Not independent | Independent |
Bring the two atoms together and one electron crosses from sodium to chlorine.
sodium + chlorine → sodium chloride
The sodium ion and the chloride ion, being oppositely charged, are then drawn towards each other by electrostatic attraction — often called the coulomb force — and an ionic compound results. The whole change is drawn in Fig. 2.2-2.
Diagram
Fig. 2.2-2 — Formation Of Sodium Chloride Upper half: shell diagrams of a sodium atom (2, 8, 1) and a chlorine atom (2, 8, 7) side by side, with a curved arrow carrying one electron from the sodium outer shell to the chlorine outer shell; label the arrow ‘transfer of one electron’. Lower half: the two resulting ions drawn in square brackets, sodium ion (2, 8) marked +, chloride ion (2, 8, 8) marked −, with facing arrows between them labelled ‘electrostatic (coulomb) force of attraction’. Beside this, show the electron dot form: Na with one dot and Cl with seven dots before transfer, and Na⁺ beside a bracketed Cl with eight electrons marked − after transfer, labelled NaCl.
2.2.4.2 Magnesium Chloride
A magnesium atom, atomic number 12, has the arrangement 2, 8, 2 and so has two valence electrons to lose. A chlorine atom has seven and needs only one.
\[\mathrm{Mg}\ (2,8,2) - 2e^{-} \rightarrow \mathrm{Mg^{2+}}\ (2,8)\]
Here is the difficulty. One chlorine atom can accept only one of the two electrons the magnesium atom must lose; accepting both would take it past an octet. So a second chlorine atom is needed to take the other.
\[2\,\mathrm{Cl}\ (2,8,7) + 2e^{-} \rightarrow 2\,\mathrm{Cl^{-}}\ (2,8,8)\]
The ratio of magnesium to chloride ions must therefore be 1 : 2, and the formula of magnesium chloride is \(\mathrm{MgCl_2}\). Notice that the formula was not chosen — it was forced by the electron count, as Fig. 2.2-3 shows.
Diagram
Fig. 2.2-3 — Formation Of Magnesium Chloride Left: shell diagram of a magnesium atom (2, 8, 2) with two chlorine atoms (2, 8, 7) drawn above and below it; two curved arrows carry one electron each from magnesium to the two chlorine atoms. Right: the magnesium ion (2, 8) in square brackets marked 2+, and two chloride ions (2, 8, 8) in square brackets each marked −, with arrows between them labelled ‘coulomb force’. Beneath, the electron dot form: Mg with two dots between two Cl symbols each with seven dots, giving Mg²⁺ and two bracketed chloride ions, labelled MgCl₂.
2.2.4.3 Calcium Oxide
A calcium atom, atomic number 20, has the arrangement 2, 8, 8, 2 and two valence electrons to lose. An oxygen atom has six valence electrons and needs two to reach an octet. This time the numbers match exactly, and one atom of each is enough.
\[\mathrm{Ca}\ (2,8,8,2) - 2e^{-} \rightarrow \mathrm{Ca^{2+}}\ (2,8,8)\]
\[\mathrm{O}\ (2,6) + 2e^{-} \rightarrow \mathrm{O^{2-}}\ (2,8)\]
Since a single oxygen atom accepts both the electrons a calcium atom gives away, the ions combine in a 1 : 1 ratio. The formula of calcium oxide is therefore \(\mathrm{CaO}\), and not \(\mathrm{Ca_2O_2}\) — a formula records the simplest ratio in which the ions combine, and Fig. 2.2-4 shows why that ratio is 1 : 1.
Diagram
Fig. 2.2-4 — Formation Of Calcium Oxide Left: shell diagrams of a calcium atom (2, 8, 8, 2) and an oxygen atom (2, 6), with two curved arrows carrying two electrons from calcium to oxygen. Right: calcium ion (2, 8, 8) in square brackets marked 2+ and oxide ion (2, 8) in square brackets marked 2−, with arrows between them labelled ‘coulomb force’. Beneath, the electron dot form: Ca with two dots beside O with six dots, giving Ca²⁺ and a bracketed oxide ion with eight electrons, labelled CaO.
2.2.5 Electrovalent Bond Formation As A Redox Process
Every electrovalent bond you have met in this section involved the same two events happening together: one atom lost electrons and another gained them. Those two events have names.
Loss of electrons is oxidation, and gain of electrons is reduction. The electropositive atom is oxidised; the electronegative atom is reduced. A change in which both occur is a redox reaction.
Take the formation of sodium chloride again, this time as a whole reaction rather than as a pair of atoms.
sodium + chlorine → sodium chloride
\[2\,\mathrm{Na} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{NaCl}\]
It can be split into two half reactions that show the electrons explicitly:
\[2\,\mathrm{Na} \rightarrow 2\,\mathrm{Na^{+}} + 2e^{-}\qquad\text{(oxidation)}\]
\[\mathrm{Cl_2} + 2e^{-} \rightarrow 2\,\mathrm{Cl^{-}}\qquad\text{(reduction)}\]
Add the two half reactions and the electrons cancel, because the two released by sodium are exactly the two taken by chlorine:
\[2\,\mathrm{Na} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{Na^{+}} + 2\,\mathrm{Cl^{-}}\]
Oxidation and reduction must always occur together, for a simple reason: the electrons lost by one substance have to go somewhere, and they go to the other.
That leads to the two terms used for the partners in such a reaction. Consider the reaction of aluminium with iron(III) oxide:
\[\mathrm{Fe_2O_3} + 2\,\mathrm{Al} \rightarrow 2\,\mathrm{Fe} + \mathrm{Al_2O_3}\]
Iron in the oxide is reduced to iron metal, which means it has accepted electrons:
\[\mathrm{Fe^{3+}} + 3e^{-} \rightarrow \mathrm{Fe}\qquad\text{(reduction)}\]
At the same time aluminium is oxidised, which means it has donated electrons:
\[\mathrm{Al} - 3e^{-} \rightarrow \mathrm{Al^{3+}}\qquad\text{(oxidation)}\]
Iron(III) oxide has taken electrons from aluminium, so it is the oxidising agent; aluminium has supplied them, so it is the reducing agent. The reaction as a whole is a redox reaction.
2.3 The Covalent (Or Molecular) Bond
Transfer works when one atom is eager to lose electrons and the other eager to gain them. But what happens when both atoms want to gain?
The atoms of non-metals usually carry five, six or seven electrons in the outermost shell — carbon, with four, and hydrogen, with one, are the exceptions you will meet most often. For such an atom, losing electrons is a poor bargain. Stripping four, five or six electrons from a small atom takes far more energy than any bond can pay back, so the atoms of these elements do not favour losing electrons.
That leaves one route open. Two atoms can each contribute electrons to a pair that lies between them, and count that pair as belonging to both. Each atom then reaches a complete outermost shell without either of them giving anything away permanently. This mutual sharing is covalent bonding, and in it each atom contributes an equal number of electrons to every shared pair.
Sharing takes place between two atoms of different non-metals, as in hydrogen chloride, and equally between two atoms of the same non-metal, as in a chlorine molecule. Both cases are worked through in full later in this section.
2.3.1 Single, Double And Triple Covalent Bonds
Two atoms need not stop at one shared pair. Depending on how many electrons each atom is short of, they may share one pair, two pairs or three.
A single covalent bond is formed by the sharing of one pair of electrons, each atom contributing one electron to that pair. It is drawn as a short line between the two symbols. Hydrogen, chlorine, hydrogen chloride, water, ammonia, methane and carbon tetrachloride are all built from single covalent bonds, as Fig. 2.3-1 shows.
Diagram
Fig. 2.3-1 — Molecules Built From Single Covalent Bonds Seven structural formulae drawn in a row and labelled beneath: H—H (hydrogen); Cl—Cl (chlorine); H—Cl (hydrogen chloride); H—O—H (water); ammonia drawn with nitrogen at the centre joined by single lines to three hydrogen atoms; methane drawn with carbon at the centre joined by single lines to four hydrogen atoms; carbon tetrachloride drawn with carbon at the centre joined by single lines to four chlorine atoms. Note under the last two that they contain four single covalent bonds each.
A double bond is formed by the sharing of two pairs of electrons between two atoms. It amounts to two single bonds acting together, and is drawn as two short lines. An oxygen molecule contains one double bond, written \(\mathrm{O{=}O}\), and a carbon dioxide molecule contains two.
A triple bond is formed by the sharing of three pairs of electrons, and is drawn as three short lines. A nitrogen molecule contains a triple bond, written \(\mathrm{N{\equiv}N}\).
Some molecules carry a mixture. A molecule of ethene, \(\mathrm{C_2H_4}\), has one double covalent bond and four single covalent bonds. A molecule of ethyne, \(\mathrm{C_2H_2}\), has one triple covalent bond and two single covalent bonds. These are collected in Fig. 2.3-2.
Diagram
Fig. 2.3-2 — Double And Triple Covalent Bonds Four structures drawn in a row and labelled: O=O (oxygen); O=C=O (carbon dioxide, showing two double bonds); N≡N (nitrogen); ethene drawn as two carbon atoms joined by a double line, each carrying two hydrogen atoms on single lines, labelled C₂H₄; ethyne drawn as H—C≡C—H, labelled C₂H₂. Mark the double bonds and the triple bonds with small arrows and the words ‘two shared pairs’ and ‘three shared pairs’.
2.3.2 Covalency
In an ionic compound we counted the electrons an atom loses or gains. In a covalent compound nothing is lost or gained, so we count something else: the number of electrons the atom puts into shared pairs.
On this count the covalency of hydrogen is 1, of oxygen 2, of nitrogen 3 and of carbon 4. Check each against the structures in Fig. 2.3-1 and you will find the same numbers of lines drawn at each atom.
2.3.3 Conditions For The Formation Of A Covalent Bond
Sharing, rather than transfer, is favoured when the following conditions hold.
- Both atoms should have four or more electrons in their outermost shells, that is, both should be non-metals. Hydrogen, beryllium, boron and aluminium are exceptions to the electron count.
- Both atoms should have high electronegativity, so that neither can strip electrons from the other.
- Both atoms should have high electron affinity.
- Both atoms should have high ionisation energy, so that neither parts with an electron easily.
- The electronegativity difference between the combining atoms should be zero or negligible.
The last condition is the one that ties this section to the previous one. A large electronegativity difference gives an ionic bond; a very small difference gives a covalent bond. The character of a bond is decided not by the elements’ names but by how far apart they stand on the electronegativity scale.
That is why the electrovalent character of chlorides and oxides falls off as you move across a period. Table 2.4 sets out the chlorides and oxides of the third period, and the change from ionic solids on the left to covalent gases and liquids on the right is plain to see.
Table 2.4 Bonding in the chlorides and oxides of the third period
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|---|
| Chloride | NaCl | MgCl₂ | AlCl₃ | SiCl₄ | PCl₃ / PCl₅ | S₂Cl₂ | Cl₂ |
| Bonding | Ionic solid | Ionic solid | Partially ionic, partially covalent solid | Covalent liquid | Covalent liquid / solid | Covalent liquid | Covalent gas |
| Oxide | Na₂O | MgO | Al₂O₃ | SiO₂ | P₂O₅ | SO₂ / SO₃ | Cl₂O₇ |
| Bonding | Ionic solid | Ionic solid | Ionic solid | Covalent solid | Covalent solid | Covalent gas / solid | Covalent gas |
Hydrogen, for its part, can combine with all the non-metals of groups 14 to 17 by covalent bonds.
2.3.4 Non-Polar Covalent Compounds
A shared pair sits between two nuclei, and both nuclei pull at it. If the two pulls are equal, the pair sits midway.
There are two ways a molecule can end up non-polar.
The first is the obvious one. When the two bonded atoms are identical, their electronegativity difference is zero, and the shared pair is attracted equally by both nuclei and lies at an equal distance from each. Hydrogen, chlorine and oxygen molecules are non-polar for this reason. Since no part of such a molecule carries even a partial charge, these molecules do not ionise in water.
The second is subtler. Two different atoms may differ so little in electronegativity, and be arranged so symmetrically, that the pulls cancel out over the molecule as a whole. Carbon tetrachloride and methane are of this kind. In methane, the shared pairs lie at equal distances from the carbon and the hydrogen atoms, because the electronegativities of carbon, 2.5, and hydrogen, 2.1, are nearly equal, and the four bonds point symmetrically outwards.
2.3.4.1 The Hydrogen Molecule
A hydrogen atom has one electron in its only shell, the K shell, which is full at two. It therefore needs one more electron to attain a duplet, the arrangement of helium.
It cannot get that electron from another hydrogen atom, since the other atom needs one just as badly. So the two atoms contribute one electron each, and the resulting pair is shared by both. Each atom now counts two electrons in its shell, and a single covalent bond, \(\mathrm{H{-}H}\), holds them together as a molecule of \(\mathrm{H_2}\).
Diagram
Fig. 2.3-3 — Formation Of A Hydrogen Molecule Two panels one above the other. Before combination: two hydrogen atoms drawn as H with a single dot and H with a single cross, labelled H-atom and H-atom. After combination: the two symbols brought together with the dot and cross side by side between them, labelled ‘one shared pair of electrons’, and beside it the structural form H—H (H₂) labelled ‘single covalent bond’. Add the note that a hydrogen molecule contains two atoms of hydrogen.
2.3.4.2 The Chlorine Molecule
A chlorine atom, 2, 8, 7, is one electron short of the argon arrangement, 2, 8, 8. Two chlorine atoms meeting each other are in exactly the same position, so neither can take from the other.
Each contributes one electron, forming one shared pair between them. Both atoms then have eight electrons in the outermost shell, and a single covalent bond, \(\mathrm{Cl{-}Cl}\), joins them into a molecule of \(\mathrm{Cl_2}\).
Diagram
Fig. 2.3-4 — Formation Of A Chlorine Molecule Before combination: two chlorine electron dot symbols side by side, one drawn with dots and one with crosses, each showing seven valence electrons (three pairs and one single), labelled Cl-atom and Cl-atom. After combination: the two symbols brought together with the single dot and single cross forming one pair between them, each atom now surrounded by eight electrons; label the central pair ‘one shared pair of electrons’ and write [Cl—Cl] (Cl₂) alongside, labelled ‘single covalent bond’.
2.3.4.3 The Oxygen Molecule
An oxygen atom, 2, 6, is two electrons short of the neon arrangement, 2, 8. One shared pair would leave both atoms one electron short, so one pair is not enough.
Each atom therefore contributes two electrons, giving two shared pairs between them. Both atoms attain an octet, and the double covalent bond formed is written \(\mathrm{O{=}O}\), giving a molecule of \(\mathrm{O_2}\).
Diagram
Fig. 2.3-5 — Formation Of An Oxygen Molecule Before combination: two oxygen electron dot symbols, one in dots and one in crosses, each showing six valence electrons, labelled O-atom and O-atom. After combination: the two symbols brought together with two pairs of electrons between them, each atom surrounded by eight electrons; label the central electrons ‘two shared pairs of electrons’ and write O=O (O₂) alongside, labelled ‘double covalent bond’. Add the note that an oxygen molecule contains two atoms of oxygen.
2.3.4.4 The Nitrogen Molecule
A nitrogen atom, 2, 5, is three electrons short of the neon arrangement, 2, 8. Two nitrogen atoms therefore contribute three electrons each, and three shared pairs form between them.
Both atoms attain an octet, and the triple covalent bond formed is written \(\mathrm{N{\equiv}N}\), giving a molecule of \(\mathrm{N_2}\).
Diagram
Fig. 2.3-6 — Formation Of A Nitrogen Molecule Before combination: two nitrogen electron dot symbols, one in dots and one in crosses, each showing five valence electrons, labelled N-atom and N-atom. After combination: the two symbols brought together with three pairs of electrons between them, each atom surrounded by eight electrons; label the central electrons ‘three shared pairs of electrons’ and write [N≡N] (N₂) alongside, labelled ‘triple covalent bond’.
2.3.4.5 The Carbon Tetrachloride Molecule
A carbon atom, 2, 4, needs four more electrons to reach the neon arrangement. A chlorine atom, 2, 8, 7, needs one to reach the argon arrangement.
One carbon atom therefore shares four electron pairs, one with each of four chlorine atoms. Carbon contributes one electron to each pair and each chlorine atom contributes the other. Carbon finishes with eight electrons around it, every chlorine atom finishes with eight, and the molecule \(\mathrm{CCl_4}\) contains five atoms in all, held by four single covalent bonds.
Each carbon–chlorine bond is itself polar, since chlorine is more electronegative than carbon. Yet the four bonds point symmetrically away from the central carbon atom, so their pulls cancel and the molecule as a whole is non-polar.
Diagram
Fig. 2.3-7 — Formation Of A Carbon Tetrachloride Molecule Before combination: a carbon electron dot symbol with four crosses at the centre, with four chlorine electron dot symbols in dots placed above, below, left and right, each showing seven valence electrons. After combination: the same arrangement with one cross and one dot paired between carbon and each chlorine atom, carbon surrounded by eight electrons and each chlorine by eight; label ‘one shared pair of electrons with each chlorine atom’. Alongside, the structural form with C at the centre joined by four single lines to four Cl atoms, labelled CCl₄ and ‘four single covalent bonds’.
2.3.4.6 The Methane Molecule
A carbon atom needs four electrons and a hydrogen atom needs one, so one carbon atom shares four electron pairs, one with each of four hydrogen atoms.
Carbon ends with an octet and each hydrogen atom with a duplet. A molecule of methane, \(\mathrm{CH_4}\), therefore contains five atoms in all — one of carbon and four of hydrogen — joined by four single covalent bonds.
Diagram
Fig. 2.3-8 — Formation Of A Methane Molecule Before combination: a carbon electron dot symbol with four dots at the centre, and four hydrogen atoms each with a single cross placed above, below, left and right. After combination: one dot and one cross paired between carbon and each hydrogen atom, carbon surrounded by eight electrons and each hydrogen by two. Alongside, the structural form with C at the centre joined by four single lines to four H atoms, labelled CH₄ and ‘four single covalent bonds’. Add the note that a methane molecule contains a total of five atoms.
2.3.5 Polar Covalent Compounds
Suppose now that the two bonded atoms differ appreciably in electronegativity. The shared pair is still shared — neither atom has taken it outright — but it no longer sits midway.
Hydrogen chloride is the standard example. Chlorine, with an electronegativity of 3, pulls the shared pair towards itself far harder than hydrogen, with an electronegativity of 2.1, can pull it back. Negative charge is therefore shifted towards the chlorine end of the bond.
The chlorine atom, having more than its fair share of the pair, develops a slight negative charge written \(\delta^{-}\). The hydrogen atom, left with less, develops a slight positive charge written \(\delta^{+}\). The bond is a polar covalent bond, and it is drawn with an arrow on the line pointing towards the atom that the shared pair has shifted to, as in Fig. 2.3-9.
Diagram
Fig. 2.3-9 — Polarity Of The Hydrogen Chloride Bond Three stages left to right. First, a hydrogen atom with a single cross beside a chlorine electron dot symbol with seven dots. Second, the shared pair formed between them, hydrogen with two electrons and chlorine with eight. Third, the structural form H—Cl with an arrow drawn on the bond line pointing towards chlorine, and the symbols δ+ written above hydrogen and δ− above chlorine. Caption note: the arrow shows the direction in which the shared pair has shifted.
The size of the electronegativity difference decides the kind of bond altogether:
- same electronegativity — the bond is non-polar;
- slightly different electronegativity — the bond is polar;
- much different electronegativity, a difference of more than 1.7 — the bond is ionic.
A polar covalent molecule has a slightly positive end and a slightly negative end, that is, two poles. It is therefore also called a dipole molecule.
Polarity has a consequence you can test with a bulb. In solution, the fractional charges of a polar covalent compound are pulled apart into full charges, and ions are produced.
The greater the electronegativity difference between the two atoms forming a bond, the more polar the molecule, and the more readily it ionises. Ammonia, water and hydrogen fluoride are other polar covalent compounds that ionise; their partial charges are shown in Fig. 2.3-10. Non-polar compounds, having no charge separation to begin with, do not ionise at all.
Diagram
Fig. 2.3-10 — Partial Charges In Some Polar Molecules Three small molecular sketches side by side. Ammonia: nitrogen at the apex marked δ−, with three hydrogen atoms below, each marked δ+. Water: oxygen at the apex marked δ−, with two hydrogen atoms below, each marked δ+. Hydrogen fluoride: H marked δ+ joined to F marked δ−. Label each sketch with the name of the molecule.
2.3.5.1 The Hydrogen Chloride Molecule
A hydrogen atom needs one electron to attain the helium duplet, and a chlorine atom, 2, 8, 7, needs one to attain the argon octet.
Each contributes one electron, and the single shared pair completes both shells at once. One molecule of hydrogen chloride therefore contains two atoms, one of hydrogen and one of chlorine, joined by one single covalent bond — a bond that is polar for the reason given above.
2.3.5.2 The Ammonia Molecule
A nitrogen atom, 2, 5, needs three electrons to attain the neon octet, and each hydrogen atom needs one.
One nitrogen atom therefore shares three electron pairs, one with each of three hydrogen atoms, and three single covalent bonds are formed. The molecule \(\mathrm{NH_3}\) contains four atoms in all — one of nitrogen and three of hydrogen. Nitrogen is more electronegative than hydrogen, so each bond is polar, with the nitrogen end slightly negative.
Diagram
Fig. 2.3-11 — Formation Of An Ammonia Molecule Before combination: a nitrogen electron dot symbol with five dots at the centre, with three hydrogen atoms each carrying a single cross placed to the left, right and below. After combination: one dot and one cross paired between nitrogen and each hydrogen atom, nitrogen surrounded by eight electrons and each hydrogen by two; label ‘one shared pair of electrons with each hydrogen atom’. Alongside, the structural form with N at the centre joined by three single lines to three H atoms, labelled NH₃ and ‘three single covalent bonds’.
2.3.5.3 The Water Molecule
An oxygen atom, 2, 6, needs two electrons to attain the neon octet, and each hydrogen atom needs one.
Each of the two hydrogen atoms therefore shares an electron pair with the oxygen atom. Hydrogen acquires a duplet, oxygen acquires an octet, and two single covalent bonds are formed. One molecule of water contains three atoms — one of oxygen and two of hydrogen.
Diagram
Fig. 2.3-12 — Formation Of A Water Molecule Before combination: an oxygen electron dot symbol with six dots at the centre, with a hydrogen atom carrying a single cross on either side, labelled H-atom, O-atom, H-atom. After combination: one dot and one cross paired on each side of the oxygen atom, oxygen surrounded by eight electrons and each hydrogen by two; label ‘one shared pair of electrons on each side’. Alongside, the structural form H—O—H labelled H₂O and ‘two single covalent bonds’.
2.3.6 Non-Polar Compared With Polar Covalent Compounds
The two kinds of covalent compound differ in one thing only — whether the shared pair sits midway — but that one difference decides how they behave in water. Table 2.5 sets the two side by side.
2.4 The Effect Of Electricity On Electrovalent And Covalent Compounds
So far the two kinds of bonding have been described in terms of electrons we cannot see. It is time to put the description to a test that anyone can watch, using a bulb.
An electric current is a flow of charge. A substance can carry a current only if it contains charged particles that are free to move. Ionic compounds are built from charged particles; covalent compounds are built from neutral molecules. If the bonding picture is right, the two should behave quite differently in a circuit.
Objective And Materials
The objective is to test the effect of electricity on electrovalent, polar covalent and non-polar covalent compounds.
You will need distilled water, beakers, electrodes of platinum or graphite, a 6 V battery, a 6 V light bulb in a holder, connecting wires, an ionic compound such as sodium chloride or magnesium chloride, a polar covalent compound such as hydrogen chloride, and a non-polar covalent compound such as carbon tetrachloride or sugar.
Procedure
Arrange an electrolytic cell as shown in Fig. 2.4-1. The circuit carries a 6 V battery, an ammeter, a bulb and two platinum electrodes, all connected in series.
Diagram
Fig. 2.4-1 — Circuit For Testing The Conductivity Of Compounds A simple circuit drawn as a closed loop. Along the top, a 6 volt battery with its positive and negative terminals marked, an ammeter on the left arm and a bulb on the right arm. At the bottom, a beaker labelled ‘electrolytic cell’ holding a liquid labelled ‘water or some liquid’, with two platinum wire electrodes dipping into it; label the left electrode anode and mark it (+), the right electrode cathode and mark it (−). Caption note: by convention the anode is drawn on the left and the cathode on the right.
Place each compound in turn — the electrovalent one, the polar covalent one and the non-polar covalent one — in a separate beaker, either as an aqueous solution or in the molten state. Dip the electrodes into the compound and watch the bulb. A glowing bulb means the compound is conducting.
Two ordinary precautions belong to this method. Rinse the electrodes in distilled water between one compound and the next, or traces of the first will confuse the second reading. And keep hands dry while handling the circuit, as you would with any electrical apparatus. The whole scheme is set out in FC 2.1.
Flowchart
FC 2.1 — Testing A Compound For Conductivity Set up circuit with battery, ammeter, bulb and electrodes in series → place the compound in the beaker, molten or in aqueous solution → dip the electrodes → observe the bulb → branch A: bulb glows → mobile ions are present → the compound is ionic, or a polar covalent compound that has ionised; branch B: bulb does not glow → no mobile ions are present → the compound is non-polar covalent
Observations
Electrovalent compounds such as sodium chloride and magnesium chloride, whether in the aqueous state or in the molten state, are good conductors of electricity: the bulb glows.
Polar covalent compounds such as hydrogen chloride in aqueous solution also show conductivity: the bulb glows.
Non-polar covalent compounds such as carbon tetrachloride, and a solution of sugar, do not conduct electricity: the bulb does not glow.
Explanation
Each observation follows from the bonding.
Electrovalent compounds such as sodium chloride are made of ions to begin with. In the solid the ions are locked in place, but in the molten state or in aqueous solution they dissociate — that is, they separate and become free to move. These moving ions, \(\mathrm{Na^{+}}\) and \(\mathrm{Cl^{-}}\), carry the current, and the bulb glows brightly.
Polar covalent compounds such as hydrogen chloride contain no ions at all before they are dissolved. In water they ionise, producing \(\mathrm{H^{+}}\) and \(\mathrm{Cl^{-}}\) ions, which are then free to move. The result is electrical conductivity, and again the bulb glows.
Non-polar covalent compounds such as carbon tetrachloride and sugar neither dissociate nor ionise. They remain as neutral molecules, which carry no charge and therefore cannot carry a current. The bulb stays dark.
2.5 Characteristic Properties Of Electrovalent And Covalent Compounds
Two compounds may look much alike in a bottle and behave quite differently on the bench. Sodium chloride is a hard solid that has to be heated fiercely before it melts; wax melts in a warm hand. That difference is not an accident of the substances. It follows directly from what holds their particles together.
Everything in this section comes from one comparison. In an electrovalent compound the particles are ions, and the forces between them are strong electrostatic attractions. In a covalent compound the particles are separate molecules, and the forces between those molecules are weak.
Nature Of The Particles
An electrovalent compound is made of ions, packed in a regular arrangement in which each ion is surrounded by ions of opposite charge. Strong electrostatic forces hold them in place and cannot be broken easily, so these compounds are hard, crystalline solids.
A covalent compound is made of molecules. The atoms inside each molecule are firmly bonded, but the forces between one molecule and the next are weak. Molecules therefore separate easily, and covalent compounds are gases, liquids or soft solids.
Melting And Boiling Points
Melting a solid means loosening the particles enough to let them slide past one another; boiling means separating them altogether. The energy needed depends on the strength of the forces holding them.
Electrovalent compounds are non-volatile: their melting and boiling points are high, because a great deal of energy is required to overcome the strong attraction between oppositely charged ions.
Covalent compounds are volatile: their melting and boiling points are low, because only the weak forces between molecules have to be overcome. The strong bonds inside each molecule are not broken at all when the substance melts or boils.
Conduction Of Electricity
Electrovalent compounds do not conduct electricity in the solid state. The ions are there, but the strong electrostatic forces in the solid keep them fixed, and a charge that cannot move cannot carry a current.
The same compounds are good conductors in the fused or aqueous state. Melting, or dissolving in water, weakens those forces, the ions become mobile, and they migrate through the liquid when a current is passed. Electrovalent compounds are also good conductors of heat.
Covalent compounds are non-conductors in the solid, molten and aqueous states alike, simply because they contain no free ions. Covalent compounds are poor conductors of heat as well.
Dissociation And Ionisation In Solution
An electrovalent compound is composed of ions before it ever meets water. Water is a polar compound, and it decreases the electrostatic force of attraction between those ions, so that in solution or in the molten state the ions separate and become mobile.
\[\mathrm{NaCl}\text{(aq)} \rightleftharpoons \mathrm{Na^{+}}\text{(aq)} + \mathrm{Cl^{-}}\text{(aq)}\]
Covalent molecules do not dissociate in this way, because there are no ions in them to be set free. Non-polar covalent compounds do not ionise even on passing an electric current.
Some covalent compounds are polar, however, and these ionise in their solutions and can then act as electrolytes:
\[\mathrm{HCl} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{Cl^{-}}\]
The particle written \(\mathrm{H_3O^{+}}\) is taken up in the next section; for the moment, notice only that the ions in this equation were made in solution, not merely released.
Solubility
Electrovalent compounds are soluble in water but insoluble in organic solvents. Water, being polar, weakens the attraction between the ions and frees them into solution, so the compound dissolves.
Covalent compounds are insoluble in water but dissolve in organic solvents. Organic solvents are non-polar, and non-polar molecules mix readily with other non-polar molecules. This is the origin of the old working rule that like dissolves like.
Rate Of Reaction
Electrovalent compounds show rapid reactions in aqueous solution. The free ions are already formed, so when two solutions are mixed the ions need only meet and unite.
Covalent compounds show slow reactions in aqueous solution. Old bonds must first be broken and new ones formed within the molecules, and that takes time.
2.6 The Coordinate (Or Dative) Bond
In every covalent bond so far, the two atoms contributed one electron each to the shared pair. That is not the only way a pair can come to lie between two atoms.
Suppose one atom already holds a pair of electrons that it is not using for bonding, and the other atom is short of a pair altogether. The first can offer its pair for both to share. The pair is then shared by two atoms, exactly as in a covalent bond, but only one of them paid for it.
A coordinate bond may equally be described as a bond formed between an ion and an atom of a polar covalent molecule that carries one or more unshared pairs of electrons. Because it has something of the covalent bond about it, in that the pair is shared, and something of the ionic bond, in that the pair came from one side only, it is sometimes called a co-ionic bond as well.
The atom that provides the electron pair is the donor. The atom or ion that shares the donated pair is the acceptor. A coordinate bond is drawn as an arrow, \(\rightarrow\), pointing from the donor to the acceptor.
The hydronium ion, \(\mathrm{H_3O^{+}}\), and the ammonium ion, \(\mathrm{NH_4^{+}}\), both contain a coordinate bond, and this section follows the formation of each.
2.6.1 Conditions For The Formation Of A Coordinate Bond
Two conditions must be met, one for each partner.
- One of the two atoms must have at least one lone pair of electrons. Ammonia and water both qualify.
- The other atom must be short of at least a pair of electrons. A hydrogen ion, \(\mathrm{H^{+}}\), qualifies, since it has no electrons at all.
The lone pair is the whole point. It is the pair that is offered to the other atom to make the coordinate bond, so an atom with no lone pair cannot form one.
2.6.2 Formation Of The Hydronium Ion And The Lone Pair Effect Of Oxygen
An oxygen atom has six electrons in its valence shell and a hydrogen atom has one. In a water molecule, two hydrogen atoms each share one electron with the oxygen atom, forming two single covalent bonds.
Count the oxygen atom’s electrons after that. Four of its six are used in the two shared pairs; the remaining two form a pair that is shared with nobody. So a water molecule carries two shared pairs and two lone pairs, as Fig. 2.6-1 shows.
Diagram
Fig. 2.6-1 — Shared Pairs And Lone Pairs In A Water Molecule A dot diagram of water: oxygen at the centre with a hydrogen atom on either side; the two pairs between oxygen and the hydrogen atoms labelled ‘shared pairs’, and the two pairs on the far side of the oxygen atom labelled ‘lone pairs’. Beside it, a second sketch of the same molecule drawn to shape, with the two O—H bonds meeting at the oxygen atom to give the bent, wedge-shaped molecule; mark the lone pairs above the oxygen atom and label the shape.
The lone pairs are not idle. They repel the bond pairs, and this repulsion pushes the two O—H bonds together, so that a water molecule is not straight but wedge shaped.
The O—H bonds are also polarised, because oxygen, with an electronegativity of 3.5, pulls the shared pairs much harder than hydrogen, with 2.1. The oxygen atom therefore carries a slight negative charge and each hydrogen atom a slight positive charge, which is why water is a polar molecule.
Now add an acid to water. Strong dipole interaction — attraction between the slightly positive and slightly negative parts — takes place between the polar water molecules and the polar acid molecules, and as a result a hydrogen ion is released from the acid.
That hydrogen ion has no electrons of its own, and a water molecule has two lone pairs going spare. The ion attaches itself at once to one of them, and the pair is then shared between oxygen and the newcomer. A coordinate bond has formed.
\[\mathrm{H^{+}} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}}\]
The particle produced is the hydronium ion, a hydrated proton. It carries a single positive charge, because the hydrogen ion brought a charge of \(+1\) with it and no electrons to cancel it.
Diagram
Fig. 2.6-2 — Formation Of The Hydronium Ion Left: dot diagram of a water molecule, with the oxygen atom’s two shared pairs and two lone pairs shown, labelled WATER. Centre: H⁺ labelled HYDROGEN ION, with an arrow from one lone pair of the oxygen atom towards it. Right: the resulting particle drawn in square brackets with a + sign outside, showing oxygen joined to three hydrogen atoms and retaining one lone pair, labelled HYDRONIUM ION. Below, the same ion drawn in line form: H—O—H with the third H joined by an arrow, the plain lines labelled ‘covalent bond’ and the arrow labelled ‘coordinate bond’.
So the hydronium ion contains one lone pair, two single covalent bonds and one coordinate bond.
2.6.3 Formation Of The Ammonium Ion And The Lone Pair Effect Of Nitrogen
A nitrogen atom has five valence electrons and a hydrogen atom has one. In ammonia, each of the three hydrogen atoms is bonded to the nitrogen atom by a shared pair, one electron of each pair coming from the hydrogen atom and one from the nitrogen atom.
That uses three of nitrogen’s five valence electrons. The other two remain as a lone pair on the nitrogen atom.
A hydrogen ion has no electrons in its outermost shell, so it cannot form an ordinary covalent bond with anything. What it can do is share the lone pair the nitrogen atom is holding. When it does, it acquires a share in two electrons and so reaches the helium arrangement, and the bond formed between the nitrogen atom of ammonia and the hydrogen ion is a coordinate bond.
\[\mathrm{NH_3} + \mathrm{H^{+}} \rightarrow \mathrm{NH_4^{+}}\]
The nitrogen atom is the donor and the hydrogen ion is the acceptor. The linkage forms the ammonium ion, which carries a single positive charge.
Diagram
Fig. 2.6-3 — Formation Of The Ammonium Ion Left: dot diagram of an ammonia molecule, nitrogen at the centre with three hydrogen atoms attached by shared pairs and one lone pair marked with a labelled arrow reading ‘lone pair’; label the molecule ‘ammonia molecule’. Centre: H⁺ labelled ‘hydrogen ion (proton)’, with a curved arrow from the lone pair towards it. Right: the resulting ion drawn in square brackets with a + sign outside, showing nitrogen joined to four hydrogen atoms, labelled ‘ammonium ion’. Beside it, the same ion in line form with three plain lines labelled ‘covalent bond’ and one arrow labelled ‘coordinate bond’.
One point deserves emphasis. Once the coordinate bond has formed, all four N—H bonds in the ammonium ion become identical. No experiment can pick out which of the four was made by donation. The distinction matters when we are working out how the ion was built, and not afterwards.
2.6.4 Formation Of The Hydroxyl Ion
The hydronium ion was made by adding a hydrogen ion to water. Removing one has an equally definite result.
When a hydrogen ion is taken away from a water molecule, it leaves without its electron, because it had none to take. The shared pair that joined it to the oxygen atom stays behind with oxygen, which is the more electronegative of the two.
\[\mathrm{H_2O} \rightarrow \mathrm{H^{+}} + \mathrm{OH^{-}}\]
The remaining particle has one more electron than it has protons, so it carries a charge of \(-1\). It is the hydroxyl ion, or hydroxide ion, \(\mathrm{OH^{-}}\).
Diagram
Fig. 2.6-4 — Structure Of The Hydroxyl Ion Left: dot diagram of a water molecule with both shared pairs and both lone pairs shown, and an arrow indicating the removal of one hydrogen atom without its electron. Right: the hydroxyl ion drawn in square brackets with a − sign outside, showing oxygen joined to one hydrogen atom by a shared pair and carrying three lone pairs. Label it ‘hydroxyl ion’.
2.6.5 Self-Ionisation Of Water
Put the last two subsections together and something interesting follows about water alone, with no acid added at all.
A water molecule can itself ionise, giving a hydrogen ion and a hydroxyl ion.
\[\mathrm{H_2O} \rightleftharpoons \mathrm{H^{+}} + \mathrm{OH^{-}}\]
But a bare hydrogen ion does not stay bare in water. It is transferred straight to the oxygen atom of a neighbouring water molecule, forming a hydronium ion by the coordinate bond described earlier.
\[\mathrm{H^{+}} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}}\]
The two steps together are usually written as one, between two water molecules:
\[\mathrm{H_2O} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{OH^{-}}\]
So water produces both hydronium ions and hydroxide ions from itself. The double arrow matters here: only a very small fraction of the molecules ionise at any moment, and the change runs in both directions.
2.6.6 Compounds Carrying More Than One Type Of Bond
A compound is not obliged to use one kind of bond throughout. Several familiar substances use two, and one uses all three.
When ammonium chloride is formed, the ammonium ion — which already contains three covalent bonds and one coordinate bond — is attracted to a chloride ion carrying the opposite charge. That attraction is an ionic bond.
Ammonium chloride is therefore a good example of a compound containing all three types of bond: covalent, coordinate and ionic.
Other combinations are common:
- compounds with both electrovalent and covalent bonds — sodium hydroxide, \(\mathrm{NaOH}\), and calcium carbonate, \(\mathrm{CaCO_3}\);
- a compound with both covalent and coordinate bonds — sulphuric acid, \(\mathrm{H_2SO_4}\), whose structure is drawn in Fig. 2.6-5.
Diagram
Fig. 2.6-5 — Structure Of A Sulphuric Acid Molecule A sulphur atom at the centre. Plain single lines join it to two —O—H groups, one on the left and one on the right, and these lines are labelled ‘covalent bonds’. Two arrows point from the sulphur atom to two further oxygen atoms, one drawn above and one below, and these arrows are labelled ‘coordinate bonds’. Label the whole structure H₂SO₄.
2.7 Summary
- The electrons in the outermost shell of an atom are its valence electrons, and only these take part in chemical change.
- The noble gases are unreactive because their outermost shells are complete — two electrons, a duplet, in helium, and eight electrons, an octet, in the others.
- Atoms of other elements combine by redistributing their valence electrons so as to acquire the electron arrangement of the nearest noble gas.
- The force of attraction that holds two atoms together in a molecule is called a chemical bond.
- Stability can be reached in three ways: by the transfer of electrons, giving an electrovalent or ionic bond; by the mutual sharing of one, two or three pairs, giving a covalent or molecular bond; and by the sharing of a pair provided by one atom alone, giving a coordinate or dative bond.
- An ion is a charged particle formed when an atom gains or loses electrons. Loss of electrons gives a positively charged cation; gain of electrons gives a negatively charged anion.
- An electrovalent compound is formed by the transfer of electrons from an atom of an electropositive element to an atom of an electronegative element, the oppositely charged ions being held by electrostatic attraction.
- Electrovalency is the number of electrons an atom loses or gains in forming an electrovalent bond.
- An electrovalent bond is favoured by a low ionisation potential in one atom, a high electron affinity in the other, and a large electronegativity difference between them. Bonds between metals and non-metals are therefore ionic.
- Fluorine is the most electronegative element and caesium is among the most electropositive, so caesium fluoride is the most ionic compound.
- In an electron dot symbol, the symbol of the element stands for the nucleus and inner shells, and dots around it stand for the valence electrons alone.
- In sodium chloride, one electron passes from sodium (2, 8, 1) to chlorine (2, 8, 7), giving Na⁺ (2, 8) and Cl⁻ (2, 8, 8).
- The properties of an ion are entirely different from those of the atom it came from, in colour, toxicity, chemical activity and electrical state.
- In magnesium chloride, one magnesium atom loses two electrons to two chlorine atoms, so the ions combine in the ratio 1 : 2 and the formula is \(\mathrm{MgCl_2}\).
- In calcium oxide, one calcium atom loses two electrons to one oxygen atom, so the ions combine in the ratio 1 : 1 and the formula is \(\mathrm{CaO}\).
- The formation of an electrovalent bond is a redox process: loss of electrons is oxidation and gain of electrons is reduction, and the two always occur together. \[2\,\mathrm{Na} \rightarrow 2\,\mathrm{Na^{+}} + 2e^{-}\qquad\text{(oxidation)}\] \[\mathrm{Cl_2} + 2e^{-} \rightarrow 2\,\mathrm{Cl^{-}}\qquad\text{(reduction)}\]
- An oxidising agent is an acceptor of electrons and a reducing agent is a donor of electrons.
- A covalent bond is formed by the mutual sharing of one or more pairs of electrons, each atom contributing an equal number of electrons. Non-metals bond in this way because losing several electrons costs too much energy.
- Sharing one pair gives a single bond, two pairs a double bond and three pairs a triple bond, drawn as one, two and three short lines.
- Covalency is the number of electrons of an atom that take part in forming shared pairs: hydrogen 1, oxygen 2, nitrogen 3, carbon 4.
- A covalent bond is favoured when both atoms are non-metals with high electronegativity, high electron affinity and high ionisation energy, and when the electronegativity difference between them is zero or negligible.
- Across a period, the electrovalent character of chlorides and oxides decreases, so that ionic solids on the left give way to covalent liquids and gases on the right.
- A covalent compound is non-polar when the shared pairs are equally distributed, either because the two atoms are identical or because the molecule is symmetrical — hydrogen, chlorine, oxygen, nitrogen, carbon tetrachloride and methane.
- A covalent compound is polar when the shared pair lies closer to the more electronegative atom, so that fractional charges δ+ and δ− develop — hydrogen chloride, hydrogen fluoride, ammonia and water. Such a molecule is a dipole molecule.
- A bond is non-polar when the electronegativity difference is zero, polar when it is small, and ionic when it exceeds 1.7.
- Ionisation is the process by which covalent compounds are converted into ions; polar covalent compounds ionise in water while non-polar ones do not.
- In the conductivity test, ionic compounds conduct when molten or in solution, polar covalent compounds conduct in aqueous solution because they ionise, and non-polar covalent compounds do not conduct at all. Only free ions can carry a current.
- Electrovalent compounds are hard crystalline solids of high melting and boiling point, soluble in water and insoluble in organic solvents, and react rapidly in solution, all because of the strong electrostatic forces between their ions.
- Covalent compounds are gases, liquids or soft solids of low melting and boiling point, insoluble in water but soluble in organic solvents, and react slowly, all because of the weak forces between their molecules.
- Electrovalent compounds conduct heat well; covalent compounds conduct it poorly.
- A coordinate bond is formed when the shared pair is provided entirely by one atom, the donor, and shared with an atom or ion short of a pair, the acceptor. It is drawn as an arrow from donor to acceptor.
- A lone pair is a pair of electrons in the outermost shell that is not shared with any other atom. A water molecule has two lone pairs and two shared pairs, and its lone pairs repel the bond pairs, making the molecule wedge shaped.
- A hydrogen ion attaches to a lone pair of a water molecule to give the hydronium ion, which contains two covalent bonds, one coordinate bond and one lone pair. \[\mathrm{H^{+}} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}}\]
- A hydrogen ion attaches to the lone pair of an ammonia molecule to give the ammonium ion, after which all four N—H bonds become identical. \[\mathrm{NH_3} + \mathrm{H^{+}} \rightarrow \mathrm{NH_4^{+}}\]
- Removing a hydrogen ion from water leaves the shared pair with oxygen and gives the hydroxyl ion, \(\mathrm{OH^{-}}\).
- Water ionises to a small extent by itself, producing hydronium and hydroxide ions. \[\mathrm{H_2O} + \mathrm{H_2O} \rightleftharpoons \mathrm{H_3O^{+}} + \mathrm{OH^{-}}\]
- Sodium hydroxide and calcium carbonate contain electrovalent and covalent bonds; sulphuric acid contains covalent and coordinate bonds; ammonium chloride contains all three types.
Flowchart
FC 2.2 — Deciding The Type Of Bond Between Two Atoms Two atoms meet, each with an incomplete outermost shell → is the electronegativity difference large, as between a metal and a non-metal? → yes: electrons are transferred, ions form, the bond is electrovalent → no: electrons are shared, the bond is covalent → within covalent, is the shared pair equally distributed? → yes: non-polar covalent; no: polar covalent → separately, if one atom supplies both electrons of the shared pair to an atom or ion short of a pair, the bond is coordinate
2.8 Key Terms
- Acceptor — the atom or ion that shares the electron pair donated in a coordinate bond.
- Ammonium ion — the ion NH₄⁺, formed when a hydrogen ion shares the lone pair of an ammonia molecule.
- Anion — a negatively charged ion, formed when an atom gains one or more electrons.
- Cation — a positively charged ion, formed when an atom loses one or more electrons.
- Chemical bond — the force of attraction that holds two atoms together in a molecule.
- Coordinate bond — a bond in which the shared pair of electrons is provided entirely by one of the two atoms but shared by both.
- Covalency — the number of electrons of an atom that take part in forming shared pairs.
- Covalent bond — a bond formed by the mutual sharing of one or more pairs of electrons between two atoms.
- Covalent compound — a compound whose atoms are held together by covalent bonds.
- Dative bond — another name for a coordinate bond.
- Dipole molecule — a molecule carrying a slight positive charge at one end and a slight negative charge at the other.
- Donor — the atom that provides the electron pair for a coordinate bond.
- Double bond — a bond formed by the sharing of two pairs of electrons between two atoms.
- Duplet — two electrons in the outermost shell, the complete arrangement of helium.
- Electron dot symbol — a symbol of an element surrounded by dots that stand for its valence electrons alone.
- Electronegative element — an element whose atoms readily accept electrons to form negative ions.
- Electropositive element — an element whose atoms readily lose electrons to form positive ions.
- Electrovalency — the number of electrons an atom loses or gains in forming an electrovalent bond.
- Electrovalent bond — the bond formed by the electrostatic attraction between a cation and an anion after electrons have been transferred.
- Electrovalent compound — a compound formed by the transfer of electrons from an electropositive atom to an electronegative atom.
- Hydronium ion — the ion H₃O⁺, formed when a hydrogen ion shares a lone pair of a water molecule.
- Ion — a charged particle formed when an atom gains or loses one or more electrons.
- Ionic bond — another name for an electrovalent bond.
- Ionisation — the process by which covalent compounds are converted into ions.
- Lone pair — a pair of electrons in the outermost shell of an atom that is not shared with any other atom.
- Molecular bond — another name for a covalent bond.
- Non-polar covalent compound — a covalent compound in which the shared pairs are equally distributed between the bonded atoms.
- Non-volatile — not easily vaporised; having high melting and boiling points.
- Octet — eight electrons in the outermost shell, the complete arrangement of the noble gases other than helium.
- Oxidation — loss of electrons by an atom or ion.
- Oxidising agent — a substance that accepts electrons.
- Polar covalent compound — a covalent compound in which the shared pair lies nearer one atom, giving fractional charges on the two atoms.
- Redox reaction — a reaction in which oxidation and reduction occur together.
- Reducing agent — a substance that donates electrons.
- Reduction — gain of electrons by an atom or ion.
- Single covalent bond — a bond formed by the sharing of one pair of electrons, each atom contributing one electron.
- Triple bond — a bond formed by the sharing of three pairs of electrons between two atoms.
- Valence electrons — the electrons in the outermost shell of an atom, which alone take part in chemical change.
- Volatile — easily vaporised; having low melting and boiling points.
2.9 Points To Ponder
- An atom is said to combine in order to reach the arrangement of a noble gas. But an atom has no wishes and makes no plans. What would a more careful way of saying the same thing sound like, and what is actually driving the change?
- The boundary between polar covalent and ionic is drawn at an electronegativity difference of 1.7. What would you expect a compound sitting very close to that line to be like, and does a sharp boundary describe nature or only describe our filing system?
- The formula NaCl gives a ratio of ions rather than a count of atoms in a particle, while the formula CH₄ gives a genuine count. Which other things you have learnt about these two kinds of compound follow from that single difference?
- Once the ammonium ion has formed, all four N—H bonds are identical, and no experiment can tell which one was the coordinate bond. Is the coordinate bond then a real feature of the ion, or only a description of how it was built? Does the distinction still earn its place?
- Water is used as the standard solvent in the laboratory, dissolves ionic compounds, ionises polar molecules and even ionises itself. Which properties of the water molecule, worked out in this chapter, are responsible for all of that?
- Suppose electronegativity differences among the elements were far smaller than they actually are, so that no pair of elements differed by more than a little. Which of the substances and properties described in this chapter would still exist, and which would not?
2.10 Objective Questions
Multiple Choice Type
Choose the correct option.
- The particles present in a crystal of magnesium chloride are: [1]
- molecules of MgCl₂
- magnesium atoms and chlorine atoms
- magnesium ions and chloride ions
- magnesium ions and chlorine molecules
- The number of shared pairs of electrons in a molecule of nitrogen is: [1]
- one
- two
- three
- four
- An element with the electron arrangement 2, 8, 2 will most probably form: [1]
- an anion carrying a charge of −2
- a cation carrying a charge of +2
- four covalent bonds
- no bond at all
- A compound conducts electricity when molten but not when solid. The compound is most likely to be: [1]
- non-polar covalent
- polar covalent
- electrovalent
- held by coordinate bonds only
- The covalency of nitrogen in an ammonia molecule is: [1]
- 1
- 2
- 3
- 5
- Which of the following molecules is polar? [1]
- methane
- carbon tetrachloride
- chlorine
- water
- In the formation of a coordinate bond, the donor atom must possess: [1]
- a positive charge
- an incomplete octet
- at least one lone pair of electrons
- at least one double bond
- A substance that accepts electrons in a reaction is called: [1]
- a reducing agent
- an oxidising agent
- an electrolyte
- a dipole
Assertion–Reason Type
For each pair, choose the answer from the codes given below. (1) Both A and R are true, and R is the correct explanation of A. (2) Both A and R are true, but R is not the correct explanation of A. (3) A is true but R is false. (4) A is false but R is true.
Assertion (A): Sodium chloride has a much higher melting point than carbon tetrachloride. Reason (R): Strong electrostatic forces act between the ions of an electrovalent compound. [1]
Assertion (A): A molecule of carbon tetrachloride is non-polar. Reason (R): The carbon–chlorine bond has no electronegativity difference across it. [1]
Assertion (A): All four N—H bonds in the ammonium ion are identical. Reason (R): The nitrogen atom in ammonia carries a lone pair of electrons. [1]
Assertion (A): A solution of sugar in water does not conduct electricity. Reason (R): Sugar molecules do not ionise in water. [1]
Assertion (A): Oxidation and reduction always occur together. Reason (R): The electrons lost by one substance must be gained by another. [1]
Match The Following
- Match each species in Column A with the correct description in Column B. [5]
| Column A | Column B |
|---|---|
| a. Oxygen molecule | i. contains three covalent bonds and one coordinate bond |
| b. Hydronium ion | ii. contains one double covalent bond |
| c. Ammonium ion | iii. contains a bond formed by electron transfer |
| d. Calcium oxide | iv. contains one lone pair, two covalent bonds and one coordinate bond |
| e. Methane | v. contains four single covalent bonds and no lone pair on the central atom |
Fill In The Blanks
An atom that loses electrons is said to be ……………. . [1]
The pair of electrons in the outermost shell that is not shared with any other atom is called a ……………. pair. [1]
Covalent compounds are ……………. in water but dissolve in ……………. solvents. [1]
A bond formed between two atoms whose electronegativity difference is greater than ……………. is ionic in character. [1]
True Or False
An electrovalent compound conducts electricity in the solid state. [1]
Helium attains stability with a duplet rather than an octet. [1]
Every covalent bond between two different atoms is polar. [1]
A hydrogen ion contains no electrons. [1]
Name The Following
The force of attraction that holds two atoms together in a molecule. [1]
The ion formed when a hydrogen ion attaches itself to a lone pair of a water molecule. [1]
A compound containing electrovalent, covalent and coordinate bonds. [1]
The most electronegative element. [1]
Give One Word Or Phrase
The number of electrons an atom loses or gains in forming an electrovalent bond. [1]
The conversion of covalent molecules into ions in solution. [1]
The atom that supplies both electrons of a shared pair. [1]
Odd One Out, And Give A Reason
Hydrogen, chlorine, hydrogen chloride, nitrogen. [1]
Sodium chloride, calcium oxide, methane, magnesium chloride. [1]
Neon, argon, krypton, chlorine. [1]
2.11 Check Your Concepts
Define a coordinate bond. [2]
State what is meant by the covalency of an element, and give the covalency of carbon. [2]
Name the two conditions that must be satisfied before a coordinate bond can form. [2]
Give a reason why the noble gases take part in almost no chemical reactions. [2]
Give a reason why a chloride ion is neither poisonous nor coloured, although a chlorine atom belongs to a poisonous, coloured gas. [2]
Give a reason why two chlorine atoms cannot form an ionic bond with each other. [3]
Explain, with an example, what is meant by a dipole molecule. [3]
Differentiate between a polar covalent compound and a non-polar covalent compound, stating the basis of your difference. [3]
Why does an aqueous solution of a polar covalent compound conduct electricity, while the pure compound does not? [3]
What would happen to the bulb in the circuit of Fig. 2.4-1 if solid sodium chloride were placed between the electrodes and then melted? Give a reason for the change. [3]
Differentiate between electrovalency and covalency, stating the basis of the difference, and give one example of each. [4]
Explain why electrovalent compounds have high melting points while covalent compounds have low ones, and why the two behave differently towards water and towards organic solvents. [4]
Explain, in terms of electron arrangements, why the formula of calcium oxide is CaO while that of magnesium chloride is MgCl₂. [4]
Explain what is meant by a lone pair of electrons, describe the two effects the lone pairs have on a water molecule, and state how a lone pair is used in forming the hydronium ion. [5]
Three solids are given to you: one electrovalent, one polar covalent and one non-polar covalent. Describe how you would distinguish between them using melting behaviour, solubility and electrical conductivity, and state the reason behind each result. [5]
2.12 Equations, Conversions & Identification
Write the balanced equation for the direct combination of magnesium with chlorine, and state which element is oxidised. [2]
Write the two half reactions that make up the formation of sodium chloride from sodium and chlorine, and label each as oxidation or reduction. [3]
Complete and balance the following, and state whether the change shown is oxidation or reduction. [3]
- \(\mathrm{Ca} - 2e^{-} \rightarrow\) …………….
- \(\mathrm{O} + 2e^{-} \rightarrow\) …………….
- \(\mathrm{Cl_2} + 2e^{-} \rightarrow\) …………….
Write the equation for the self-ionisation of water, and name both ions produced. [2]
How would you obtain the ammonium ion from ammonia gas? Write the equation and name the type of bond formed in the process. [2]
Write the equation for the ionisation of hydrogen chloride in water, and state why the pure gas does not conduct electricity although its solution does. [3]
Draw the electron dot structures of the following, showing all lone pairs: [4]
- a chlorine molecule
- a water molecule
- an oxygen molecule
- an ammonia molecule
Draw the electron dot diagram for the formation of calcium oxide from its atoms, showing the charges on both ions. [3]
An element X has the electron arrangement 2, 8, 1 and an element Y has 2, 6. [4]
- State the type of ion each will form, with its charge.
- Give the formula of the compound formed between X and Y.
- Name the type of bond present in it.
- State whether the compound will conduct electricity when molten, and why.
A colourless liquid P does not conduct electricity, dissolves grease readily and does not mix with water. A white solid Q melts only at a high temperature and its aqueous solution conducts electricity. Identify the type of bonding in P and in Q, and give one reason for each identification. [3]
Look at Fig. 2.4-1. Name the labelled parts of the apparatus and state the purpose of each: [3]
- the electrode on the left
- the electrode on the right
- the bulb
State what you would observe when the following are tested in the circuit of Fig. 2.4-1, and give the reason in each case. [3]
- molten magnesium chloride
- carbon tetrachloride
- an aqueous solution of hydrogen chloride
Name the bonds present in a molecule of sulphuric acid, and state how many of each kind there are. [2]