Chapter 4: Atomic Structure And Chemical Bonding
In the 1850s, a German glassblower named Heinrich Geissler built some of the finest glass tubes anyone had ever seen. He sealed a metal plate into each end, pumped out most of the air with a mercury pump of his own design, and connected the plates to a high-voltage coil.
What happened next drew crowds. The thin gas left inside the tube lit up — pink, blue, greenish — in a steady, silent glow. Lecturers toured with these tubes. People bought them as ornaments. Nobody could say what the glow actually was.
Twenty years later, William Crookes began pumping the tubes far emptier than Geissler had. As the gas thinned, the coloured glow faded, and something stranger took its place: a patch of greenish light on the glass at the far end, directly opposite the negative plate. Put a metal cross in the way, and a sharp shadow of the cross appeared in that patch. Whatever was crossing the tube was travelling in straight lines, from the negative plate outward, and it could be blocked.
It was not light. It was not the gas. It came out of the metal itself.
Within a few decades, that faint green patch had been traced back to particles smaller than any atom — and the atom, which for two thousand years had meant ‘that which cannot be cut’, turned out to have parts.
Diagram
Fig. 4.0-1 — A Glowing Discharge Tube Sealed glass tube lying horizontally, connected at both ends to an induction coil; label the negative plate (cathode) at the left, the positive plate (anode) at the right, the side tube leading to a vacuum pump, a metal cross standing upright inside the tube, and the bright greenish patch on the far end of the glass with the sharp shadow of the cross falling on it.
4.1 Early Ideas About Matter
Take a lump of sugar and crush it. Crush the grains again, and again, as fine as you can. A question follows almost by itself: does this go on for ever, or do you finally reach a particle that cannot be broken any further?
Thinkers in India and in Greece asked exactly that question long before anyone could test it. Their answers were guesses, but they were remarkably good guesses, and chemistry still uses the word they left behind.
4.1.1 The Search For The Ultimate Particle
The earliest recorded answer comes from India. The sage Kanada, whose ideas are traditionally placed around the sixth century BCE, held that matter is made of minute, indestructible particles which he called paramanu — param meaning ultimate and anu meaning particle.
Kanada’s picture had two parts to it. A paramanu, he said, does not exist on its own; it joins with other paramanus to form a larger particle called an anu, which is what we would today call a molecule — a group of joined particles that can exist by itself. He also held that paramanus are of different kinds, each kind with its own properties.
In Greece, the philosopher Democritus (460 BCE – 370 BCE) reached a similar conclusion. He called the ultimate particle atomos, from the Greek for indivisible — that which cannot be cut. From atomos comes our word atom. The order in which these ideas appeared is set out in Fig. 4.1-1.
Diagram
Fig. 4.1-1 — A Timeline Of Ideas About The Ultimate Particle A horizontal timeline running left to right with four labelled markers: Kanada in India (about 6th century BCE) with the caption paramanu joins to form anu; Democritus in Greece (460–370 BCE) with the caption atomos, meaning indivisible; John Dalton in England (1808) with the caption first scientific atomic theory; and an open arrow at the right labelled the discovery of sub-atomic particles.
Both of these were philosophical ideas. Neither Kanada nor Democritus had a balance, a burner or an experiment to test what he had proposed. That step waited more than two thousand years.
4.1.2 Dalton’s Atomic Theory
The first scientific theory of the structure of matter was put forward by John Dalton in 1808. He was working from measurements of how substances combine, not from philosophy, and that is what makes his theory the starting point of modern chemistry.
Dalton pictured the atom as a solid, indivisible building block. Matter differs from matter, he argued, because the atoms that make it up are different.
The Postulates Of Dalton’s Theory
- Matter is made up of very small, indivisible particles called atoms.
- Atoms can neither be created nor destroyed.
- The atoms of one element are alike in all respects, and they differ from the atoms of every other element.
- Atoms of an element combine in small numbers to form molecules.
- Atoms of one element combine with atoms of another element in a simple ratio to form the molecules of a compound.
- The atom is the smallest unit of matter that can take part in a chemical reaction.
How The Theory Has Fared Since
Dalton’s theory held for almost a century, and then experiment overtook it. We now know that the atom is not indivisible at all, and that atoms of the same element need not be alike in every respect. Of his six postulates, the one that has survived untouched is the sixth: the atom is still the unit that takes part in a chemical reaction.
That is not a failure of the theory. A good theory is one that explains what is known and can be tested further — and it was the testing of Dalton’s atom that opened it up.
4.2 The Element And The Atom
Before we can take an atom apart, we need to be clear about two words that the rest of this chapter leans on constantly — element and atom. They are related, and they are easy to blur together.
4.2.1 What An Element Is
Some substances can be broken down into simpler ones. Heat chalk strongly and it gives off a gas, leaving a different white solid behind. Other substances refuse to break down at all. An element is a substance of this second kind — one made up of only one type of atom.
Carbon is an element. You may heat it, hammer it, dissolve it or pass electricity through it, and you will still have carbon; there is no ordinary way to split it into two simpler substances. Nor can carbon be built up by combining anything simpler.
The words ‘ordinary physical or chemical methods’ in that definition are doing real work, so it is worth saying what they exclude. In radioactivity — the spontaneous emission of highly penetrating rays by certain elements — and in high-energy nuclear reactions, one element genuinely can change into another. Those are nuclear changes, not the heating, dissolving and reacting of an ordinary laboratory.
4.2.2 What An Atom Is
Now take a small piece of zinc and grind it. The powder is still zinc — it still has the grey colour of zinc, and it still fizzes in dilute acid as zinc does. Grind the powder finer and the same is true. This is shown in Fig. 4.2-1.
Diagram
Fig. 4.2-1 — Dividing A Piece Of Zinc Four stages drawn left to right and joined by arrows: a lump of zinc, then coarse zinc granules, then fine zinc powder, then a single circle labelled one atom of zinc; label the first three stages still show the properties of zinc and the fourth stage the smallest particle that is still zinc.
But this cannot go on for ever. A stage comes at which the particles cannot be divided any further and still be zinc. Those particles are the atoms of zinc. An atom is the smallest particle of an element that shows all the properties of that element.
4.3 Discovery Of The Electron
Dalton’s atom was solid and indivisible. The first crack in that picture came from an unlikely direction — not from breaking a substance open, but from passing electricity through a gas.
An early clue came in 1833, when Michael Faraday showed that the flow of electricity is really the flow of charged particles. If electricity is carried by particles, and matter can be made to release electricity, then matter must contain charged particles of its own. Proving it took another sixty years and a piece of apparatus made of glass.
4.3.1 The Discharge Tube And Cathode Rays
Air and other gases are ordinarily poor conductors of electricity. Hold a charged object near a gas at ordinary pressure and nothing flows. The British scientist William Crookes found that this changes completely if two conditions are met together: the gas must be at a very low pressure, and the voltage across it must be very high.
Crookes worked with gas at a pressure of about \(0.01\ \text{mm}\) of mercury, and applied about \(10\,000\) volts from an induction coil. Under those conditions the gas conducts, and a stream of rays travels across the tube from the negative electrode to the positive one. Because they set out from the cathode, these are called cathode rays.
The apparatus used is called a discharge tube, or cathode ray tube. It is shown in Fig. 4.3-1.
Diagram
Fig. 4.3-1 — A Discharge Tube Horizontal hard-glass tube with a flat metal plate sealed into each end; label the left plate cathode (negative electrode) connected to the negative terminal, the right plate anode (positive electrode) connected to the positive terminal, the induction coil supplying 10 000 volts below the tube, the side tube leading to a vacuum pump, and the pressure inside marked 0.01 mm of mercury; show the cathode rays as straight arrows travelling from cathode to anode.
A discharge tube is a hard glass tube fitted with two metal plates called electrodes. The plate joined to the positive terminal of the supply is the anode; the plate joined to the negative terminal is the cathode. A side tube lets a vacuum pump draw the gas out until the pressure inside is very low.
4.3.2 Properties Of Cathode Rays
Cathode rays cannot be seen crossing the tube. Everything known about them was worked out from what they do, and each of the following observations tells us something.
- They travel from the cathode to the anode in straight lines.
- They cause a greenish yellow fluorescence on a soda-glass screen placed in the tube.
- They produce the shadow of an opaque object placed in their path, and they make a light paddle wheel rotate.
- They are affected by an electric field: they are deflected towards the positive plate and away from the negative plate. This shows that they carry a negative charge. See Fig. 4.3-2.
- When a beam of cathode rays falls on a hard metallic target such as tungsten, X-rays are produced.
- They penetrate through matter.
- They cause ionisation of the gas through which they pass.
- For the particles making up cathode rays, the ratio of charge to mass is always the same, \(e/m = 1.76\times10^{11}\ \text{C kg}^{-1}\), whatever gas is in the tube and whatever metal the cathode is made of.
Diagram
Fig. 4.3-2 — Deflection Of Cathode Rays In An Electric Field Discharge tube with two charged plates mounted one above the other in the path of the beam; label the upper plate negatively charged plate and the lower plate positively charged plate, the induction coil, and the side tube to the vacuum pump; show the beam entering straight and bending downwards towards the positively charged plate.
Property 3 is worth pausing on. A shadow means the rays are stopped by a solid object, and a turning paddle wheel means they strike it hard enough to push it. Rays that push are made of something with mass.
4.3.3 Thomson’s Conclusions
J. J. Thomson studied the cathode rays themselves — what they were made of and how they behaved. From property 8 in particular he drew a powerful conclusion: if the same particles come out of a tube filled with any gas, and out of a cathode made of any metal, then those particles must be present in every kind of matter.
Thomson concluded that:
- Cathode rays consist of negatively charged particles, now called electrons.
- These negatively charged particles are an integral part of all atoms.
- Electrons have a definite mass and a definite electric charge, and neither depends on the nature of the gas in the discharge tube.
The name itself came earlier. In 1874 G. J. Stoney had suggested the word electron — meaning ‘atom of negative electricity’ — for the particle of electricity that Faraday’s work implied. Thomson showed that the particle was real.
4.3.4 Properties Of The Electron
- Electrons from all sources are alike, with identical mass.
- They are a constituent part of all atoms.
- The mass of an electron is \(1/1837\) of the mass of a hydrogen atom, that is \(9.109\times10^{-31}\ \text{kg}\).
- An electron carries one unit of negative charge, of magnitude \(-1.602\times10^{-19}\ \text{C}\).
- The electron is extremely small; its radius is less than \(1\times10^{-15}\ \text{m}\).
The charge and the mass of an electron never change, whatever the cathode is made of and whatever gas the tube contains.
Now put two facts side by side. Every atom contains electrons, which are negatively charged; yet a piece of matter shows no charge at all. An atom must therefore hold, in addition, something positively charged — and in just the right amount, so that the total positive charge exactly balances the total negative charge of its electrons.
That reasoning is what sent scientists looking for a positive particle. It is also what happens inside the discharge tube itself: the high voltage breaks the gas atoms up into negatively charged particles and positively charged ones.
4.4 Discovery Of The Proton
If the discharge tube tears electrons out of gas atoms, then whatever is left behind must be positively charged. Finding it needed only one change to the apparatus — a cathode with holes in it.
4.4.1 Anode Rays In A Perforated Cathode Tube
E. Goldstein, a German scientist, placed a perforated cathode part-way along a discharge tube instead of sealing a plain plate into the end. With this arrangement he found a second set of rays, travelling in the direction opposite to the cathode rays — that is, from the anode side towards the cathode — and streaming out through the holes at the back of it. The set-up is shown in Fig. 4.4-1.
Diagram
Fig. 4.4-1 — Production Of Anode Rays Discharge tube with the anode at the left end and a perforated cathode, drawn as a plate with several holes, standing upright in the middle of the tube; label the anode, the cathode rays travelling leftwards from the perforated cathode towards the anode, the positive rays streaming rightwards through the holes of the cathode, and the induction coil below.
Goldstein called them canal rays, because they came through the holes, or canals, in the cathode. They are also called positive rays, since they are made of positively charged particles, and anode rays, since they travel away from the anode.
The name anode rays is slightly misleading, and it is worth clearing up at once. These rays do not come out of the anode metal at all. They are the positive remains of the gas atoms in the tube — what is left when the discharge knocks electrons out of them. Cathode rays, by contrast, do come out of the metal of the cathode itself.
4.4.2 Properties Of Anode Rays
- Anode rays travel in straight lines.
- They consist of minute material particles, and so they produce mechanical effects.
- They are made up of positively charged particles.
- They are deflected by electric and magnetic fields, but in the direction opposite to that of cathode rays. This confirms that they carry a positive charge; the particles are called protons.
- They produce fluorescence on a zinc sulphide screen.
- Their charge-to-mass ratio, \(e/m\), differs from gas to gas. It is much smaller than that of the electron, and it reaches its greatest value when hydrogen is the gas in the tube.
Property 6 is the odd one out, and it carries the most information. Recall that for cathode rays \(e/m\) was fixed, whatever the gas — the same particle came out of everything. Here the value changes with the gas, which means these positive particles are not all the same. They are the leftovers of whichever atoms happen to be in the tube. Hydrogen gives the largest \(e/m\) because the hydrogen atom is the lightest of all, so its positive residue is the lightest particle available.
The two kinds of rays are compared in Table 4.1.
4.4.3 Properties Of The Proton
- A proton carries one unit of positive charge, \(+1\), of value \(1.602\times10^{-19}\ \text{C}\) — equal in size to the charge on an electron, but opposite in sign.
- Its mass is the same as that of a hydrogen atom, that is \(1\) amu, or \(1.672\times10^{-24}\ \text{g}\). This is about 1837 times the mass of an electron.
- The proton lies in the central part of the atom, the nucleus — the small dense core that we shall come to in Section 4.6.
There is a neat way to see what a proton actually is. A hydrogen atom has one proton and one electron. Take the electron away, and only the proton is left:
hydrogen atom − one electron → proton
\[\mathrm{H} - e^{-} \rightarrow \mathrm{H^{+}}\]
So a proton is simply a hydrogen atom stripped of its electron. That is exactly what the discharge tube does to the hydrogen in it, and it is why the positive particle from hydrogen is the lightest one that can be produced in such a tube.
4.5 Thomson’s Plum Pudding Model
By the early 1900s two facts were settled: atoms contain electrons, and atoms are electrically neutral. The next task was to say how the negative and positive parts are arranged inside. J. J. Thomson offered the first serious answer in 1904.
4.5.1 The Model
Thomson pictured the atom as a single sphere of positive charge, spread evenly throughout, with the electrons stuck into it here and there. His model says:
- The atom is a sphere of uniform positive charge, with the electrons embedded in it.
- The total positive charge equals the total negative charge, so the atom as a whole is electrically neutral.
- The mass of the atom is uniformly distributed through the sphere.
Notice what point 3 claims: there is no heavy centre and no light outside. Every part of a Thomson atom is as massive as every other part. Fig. 4.5-1 shows the arrangement.
Diagram
Fig. 4.5-1 — Thomson’s Plum Pudding Model Of The Atom A large circle shaded evenly to show a sphere of uniform positive charge, with six small circles bearing minus signs scattered at random inside it; label the shaded sphere sphere of uniform positive charge and the small circles electrons.
4.5.2 Limitations Of The Model
Thomson’s picture accounted for neutrality, and for very little else.
- It made no room for a central nucleus in the atom.
- It could not explain how the positive and negative charges inside the atom are held apart without the atom collapsing. So it failed to explain why atoms are stable.
- It gave no insight into how the electrons are arranged, or how atoms bond with one another chemically.
- It could not explain the results of Rutherford’s gold foil experiment of 1911.
That last point was decisive. A model survives only as long as no experiment contradicts it, and within seven years an experiment did.
4.6 Discovery Of The Nucleus
Thomson’s atom was uniform throughout. To test that, you would need something small enough and fast enough to be sent into an atom, and then you would watch how it came out. In 1911 Lord Rutherford, a scientist from New Zealand, did exactly that.
4.6.1 Rutherford’s Gold Foil Experiment
The probe he used was the alpha particle, written \(\alpha\). Alpha particles are shot out naturally by radioactive substances. Each one is a doubly charged helium ion, \(\mathrm{He^{2+}}\), containing two protons and two neutrons — a helium atom from which both electrons have been removed. It is therefore heavy, fast and positively charged.
Rutherford directed a stream of alpha particles at a very thin sheet of gold, only about one millionth of a centimetre thick. Gold was chosen because he wanted the thinnest layer possible, and gold is the most malleable of metals — it can be beaten into the thinnest foil. Around the foil was a screen that gave a tiny flash of light wherever an alpha particle struck it, so the path of each particle could be traced. The arrangement is shown in Fig. 4.6-1.
Diagram
Fig. 4.6-1 — Deflection Of Alpha Particles By A Gold Leaf Source of alpha particles at the left inside a lead block, sending a narrow beam through a slit onto a vertical sheet of gold leaf; a circular fluorescent screen surrounds the foil; label the beam of alpha particles, the gold leaf, the screen, and three marked paths — 1 bounced back towards the source, 2 deflected slightly, 3 passed straight through.
He observed that:
- Most of the alpha particles passed straight through the foil.
- Some alpha particles were deflected slightly from their straight path.
- Very few — nearly one in ten thousand — were deflected through very large angles, or bounced straight back.
Each observation says something different, and the third is the startling one. A heavy, fast particle does not rebound off nothing. To turn one back, the foil must contain something both very massive and very concentrated. Fig. 4.6-2 shows what the paths look like close to such a centre.
Diagram
Fig. 4.6-2 — Paths Of Alpha Particles Near The Nucleus A large circle representing one atom, with a small circle at its centre marked with plus signs and labelled positively charged nucleus; a set of parallel horizontal arrows enters from the left; show most arrows passing straight through the outer part of the atom, one bending as it passes near the nucleus and labelled deflected α-particle, and one striking almost head-on and turning back the way it came.
The result also depends on what the foil is made of. A heavy metal such as platinum gives the same observations as gold. But with a light nucleus such as lithium, a fast alpha particle may simply push the light nucleus aside, and may not be turned back at all.
4.6.2 Rutherford’s Atomic Model
From these observations Rutherford put forward a new picture of the atom. Each of its points is tied to something he saw.
- The atom contains a large empty space, called the extra-nuclear space. This is why most alpha particles pass through the foil without deviating.
- There is a positively charged mass at the centre of the atom, called the nucleus, in which the entire mass of the atom is concentrated. The nucleus is therefore the densest part of the atom. Its positive nature is supported by the fact that positively charged alpha particles approaching it are deflected — like charges repel — which proved that the nucleus contains protons.
- The nucleus is very small compared with the size of the whole atom. This is proved by the fact that most alpha particles pass straight through.
- The electrons revolve around the nucleus at large distances from it and from each other, in circular orbits within the empty space of the atom.
- The atom as a whole is electrically neutral, because the total positive charge of the nucleus is balanced by the total negative charge of the electrons. That is, the number of protons and the number of electrons in an atom are equal.
- The electrons revolve around the nucleus in close circular paths called orbits. The force of attraction between the negatively charged electrons and the positively charged nucleus is counterbalanced by the centrifugal force acquired by the revolving electrons.
Point 6 is easier to picture with a comparison Rutherford himself drew. The solar system has the Sun at the centre, holding most of the mass, with the planets going round it at great distances. In the atom, the nucleus holds the main mass and the electrons revolve around it in orbits or shells. Fig. 4.6-3 sets the two side by side.
Diagram
Fig. 4.6-3 — Rutherford’s Model Compared With The Solar System Upper drawing: an atom with a small shaded central circle containing plus signs, labelled nucleus and proton, surrounded by two concentric circles labelled shells, with small circles bearing minus signs labelled electrons placed on the circles. Lower drawing: the Sun at the centre with several planets on elliptical orbits around it. A single caption line links the two: the nucleus holds the mass as the Sun does, and the electrons revolve as the planets do.
4.6.3 Drawbacks Of Rutherford’s Model
The very comparison that makes the model easy to picture is also its main weakness. Planets are not charged; electrons are.
According to the classical laws of mechanics and electrodynamics, an electrically charged particle that is in motion inevitably radiates energy. An electron moving continually around the nucleus should therefore lose energy all the time. As it lost energy it would be pulled gradually inwards, spiralling closer and closer, until it collided with the nucleus. The atom would collapse. This is shown in Fig. 4.6-4.
Diagram
Fig. 4.6-4 — An Atom Losing Energy On Rutherford’s Model A small circle at the centre marked with a plus sign and labelled positively charged nucleus; a single small circle labelled negatively charged electron starting on the outermost path and following a spiral drawn with arrowheads that winds inwards through several turns and ends at the nucleus.
If that were so, atoms would be highly unstable and matter would not exist in the form we know. But matter plainly does exist, and the atom is structurally stable.
So Rutherford’s model failed to explain two things:
- the stability of the atom;
- the distribution of the electrons around the nucleus, and why the electrons do not fall into the nucleus.
Answering that took a new idea about the orbits themselves, and it came two years later.
4.7 Bohr’s Atomic Model
Rutherford had the nucleus right and the electrons wrong. In 1913, Niels Bohr, a Danish physicist, explained the stability of the atom in a different way — by putting a strict condition on where an electron is allowed to be.
Bohr’s idea was that an electron cannot revolve just anywhere. Only certain orbits are permitted, and while an electron stays in one of them, the usual rule about a moving charge losing energy simply does not apply.
The Postulates Of Bohr’s Model
- The electrons revolving around the nucleus are confined to certain fixed orbits, called shells or energy levels, each of which is associated with a fixed amount of energy.
- While it is revolving around the nucleus in an orbit, an electron neither loses nor gains energy.
- An electron revolving in a particular orbit jumps to the next orbit on gaining a certain amount of energy, and falls back to the lower one on losing it.
Because each orbit carries a fixed amount of energy of its own, Bohr called it an energy level.
Naming The Shells
For convenience the energy levels are labelled K, L, M, N, or I, II, III, IV, and so on, counting outwards from the nucleus. The orbit closest to the nucleus is the K shell. It has the least energy, and the electrons present in it are called K electrons; the same naming follows for the successive shells and their electrons. Fig. 4.7-1 shows the arrangement.
Diagram
Fig. 4.7-1 — Energy Levels Around The Nucleus On Bohr’s Model A small central circle marked with a plus sign and labelled nucleus, surrounded by six concentric circles drawn closer together as they go outwards; label the circles from the innermost outwards as K, L, M, N, O, P, and mark the innermost three additionally as first energy level (K shell), second energy level (L shell) and third energy level (M shell); mark the shell numbers n = 1 to n = 6 along the left.
Notice from the figure that the shells are not equally spaced. The energy of the shells rises as you move outward, but the gaps between them get smaller.
4.8 Discovery Of The Neutron
Two particles had been found, and between them they should have accounted for everything about an atom. They did not. The problem showed up in a simple piece of arithmetic about helium.
4.8.1 The Missing Mass In The Nucleus
An atom of helium contains 2 protons. The mass of an electron is so small that it may be ignored, so the mass of a helium atom ought to be
mass = 2 × 1 amu = 2 amu
But the atomic mass of helium is found by experiment to be about 4.0 amu — twice as much as the protons can account for. Half the mass of the helium atom was unexplained.
Something else had to be present in the nucleus. Whatever it was, it had to satisfy two conditions at once: it must carry no electrical charge, since the atom’s charges already balance exactly, and it must have a mass equal to that of a proton, to make up the shortfall.
4.8.2 Chadwick’s Experiment And The Properties Of The Neutron
In 1932 James Chadwick found the missing particle. He bombarded light nuclei such as beryllium with alpha particles — that is, with helium nuclei — and found that an uncharged particle was thrown out. The change is shown in Fig. 4.8-1.
beryllium + alpha particle → carbon + a neutral particle
\[^{9}_{4}\mathrm{Be} + ^{4}_{2}\mathrm{He} \rightarrow ^{12}_{6}\mathrm{C} + ^{1}_{0}n\]
Diagram
Fig. 4.8-1 — Chadwick’s Discovery Of The Neutron A source of alpha particles at the left sending a narrow beam onto a small block labelled beryllium target; from the right of the block an unmarked straight arrow leaves, labelled neutrons — no charge, not deflected; below the block label the product carbon; show two charged plates, one positive and one negative, placed across the outgoing path with the arrow passing between them undeflected.
Because these particles were found to be neutral, they were named neutrons.
The properties of the neutron are few, and both of them matter:
- It is not deflected by any magnetic or electric field, which proves that it is electrically neutral.
- Its mass is \(1.676\times10^{-24}\ \text{g}\), that is 1 amu.
The Basic Structure Of Every Atom
With the electron, the proton and the neutron all identified, a single picture fits every atom:
- Electrons are negatively charged particles found outside the nucleus.
- Protons are positively charged particles found inside the nucleus.
- Neutrons are electrically neutral particles also found inside the nucleus, and they are slightly heavier than protons.
4.8.3 Dalton’s Theory Compared With The Modern Atomic Theory
We can now look back at the six postulates of Section 4.1.2 and see exactly where they stand. Table 4.2 sets Dalton’s claims against what is known today.
Dalton’s theory was not discarded; it was opened up. The atom is still the unit of chemical change, which is what a chemist uses it for. What changed is that we can now say what the atom is made of — and that is what the rest of this chapter is built on.
4.9 The Structure Of An Atom
We can now assemble the whole atom from the pieces the last five sections found. Three particles matter for chemistry — the electron, the proton and the neutron — and they are arranged in two clearly separate regions.
4.9.1 The Three Sub-Atomic Particles
An atom is built up of a number of sub-atomic particles. Modern theories recognise well over two hundred of them, with names such as the positron and the neutrino, but only three are needed to understand how atoms behave chemically. Their properties are set out in Table 4.3.
Table 4.3 — Properties Of The Sub-Atomic Particles
| Particle | Symbol | Charge (electronic unit) | Mass (amu) | Mass (grams) |
|---|---|---|---|---|
| Electron | \(_{-1}e^{0}\) or \(e^{-}\) | \(-1\) | \(1/1837\) | \(9.1\times10^{-28}\) |
| Proton | \(_{1}\mathrm{H^{1}}\) or \(p^{+}\) | \(+1\) | \(1\) | \(1.6\times10^{-24}\) |
| Neutron | \(_{0}n^{1}\) or \(n\) | \(0\) | \(1\) | \(1.6\times10^{-24}\) |
Read down the charge column and the mass column separately, because they tell different stories. The charges are \(-1\), \(+1\) and \(0\): equal and opposite for the first two, nothing for the third. The masses are 1, 1 and very nearly nothing: the two heavy particles are the ones in the nucleus.
4.9.2 The Nucleus And The Orbits
An atom has two structural parts: the nucleus, and the empty space around it in which lie the imaginary paths called orbits. Fig. 4.9-1 shows both.
Diagram
Fig. 4.9-1 — The Structure Of An Atom A small shaded circle at the centre labelled nucleus (contains protons and neutrons), surrounded by three concentric circles labelled orbits — electrons revolve in orbits or shells; place small filled dots on each circle to represent the electrons.
The protons and the neutrons are found together in the central part of the atom, the nucleus. Taken together, these two are called nucleons — the particles of the nucleus. Since the electrons weigh almost nothing, the nucleons account for practically the whole mass of the atom.
Main Features Of The Structure Of Atoms
- Atoms of all the elements, except hydrogen, are made up of three fundamental particles: electrons, protons and neutrons. The hydrogen atom has no neutron.
- The nucleus is located at the centre of the atom. It contains the protons and the neutrons, which account for the total mass of the atom.
- The nucleus is positively charged, because of the protons in it.
- The electrons lie outside the nucleus and have negligible mass.
- The number of electrons in an atom equals the number of protons in it, and so the atom is electrically neutral.
- The electrons revolve rapidly in fixed circular paths around the nucleus. These paths are called energy levels, or shells, or orbits.
- The atoms of different elements contain different numbers of electrons, protons and neutrons.
Feature 7 is the one to hold on to. If atoms differ only in how many of each particle they contain, then counting those particles is the same as identifying the element — which is exactly what the next section does.
4.10 Atomic Number And Mass Number
Atoms of different elements differ only in how many particles of each kind they contain. Two counts turn out to matter, and each has its own name and its own symbol.
4.10.1 Atomic Number
In 1913 Henry Moseley studied the X-rays given out by different elements. He found that each element has a property of its own tied to the number of protons in its atom, and that it is this number — not the weight of the atom — that decides which element it is.
That number is called the atomic number, and it is denoted by \(Z\).
- An atom of an element has its own characteristic number of protons in its nucleus, which distinguishes it from the atoms of every other element.
- The atomic number is also equal to the number of electrons in a neutral atom of the element, because in a neutral atom the number of protons equals the number of electrons.
- The number of protons gives the total positive charge present in the nucleus.
- The atomic number is unique for the atoms of each element.
Change the number of protons and you have a different element altogether. An atom with 6 protons is carbon; add one more proton and it is nitrogen. Nothing else about the atom is allowed a say in the matter.
4.10.2 Mass Number
The mass of an atom sits almost entirely in its nucleus, and the nucleus holds protons and neutrons. Counting both gives the mass number, denoted by \(A\).
The mass number is a whole-number approximation of the atomic mass measured in atomic mass units. The two are numerically close but they are not the same kind of quantity: the mass number is merely a count, while the atomic mass is a measured mass and carries a unit. The mass number of sodium is 23, and its atomic mass is 23 amu.
4.10.3 Symbolic Representation Of An Element
Both counts can be written on the symbol of the element itself, as \(^{A}_{Z}\mathrm{X}\) — the mass number above and the atomic number below. Table 4.4 shows how everything else is read off from those two figures.
Table 4.4 — Reading The Symbol \(^{A}_{Z}\mathrm{X}\)
| Symbolic representation | Atomic number [Z] | Mass number [A] |
|---|---|---|
| \(^{A}_{Z}\mathrm{X}\), where X is the symbol of the element, the superscript A is the mass number and the subscript Z is the atomic number | \(Z\) = number of protons [p] = number of electrons [e], so \(Z = p = e\) | \(A\) = number of neutrons [n] + number of protons [p], so \(A = n + p\); and since \(p = Z\), number of neutrons \(n = A - Z\) |
| Example: \(^{35}_{17}\mathrm{Cl}\) | \(Z = p = e = 17\) | \(A = n + p = 35\), so \(n = A - Z = 35 - 17 = 18\) |
The one relation to remember from all of this is the subtraction:
\[n = A - Z\]
where \(n\) is the number of neutrons, \(A\) the mass number and \(Z\) the atomic number. Protons and electrons are read off directly; only the neutrons have to be worked out.
4.11 Distribution Of Electrons In The Orbits — The Bohr–Bury Scheme
Knowing that an atom has, say, twelve electrons is not enough. What decides its chemistry is where those electrons sit — how many are in the first shell, how many in the second, and how many are left over in the outermost one. Bohr and Bury worked out the rules that govern this.
The distribution of electrons in the different shells of an atom is called its electronic configuration.
4.11.1 The Rules For Filling The Shells
The maximum possible number of electrons in a particular shell is given by the formula \(2n^{2}\), where \(n\) is the serial number of that shell counted outwards from the nucleus.
First, or K shell: \(2 \times 1^{2} = 2\) electrons Second, or L shell: \(2 \times 2^{2} = 8\) electrons Third, or M shell: \(2 \times 3^{2} = 18\) electrons Fourth, or N shell: \(2 \times 4^{2} = 32\) electrons, and so on.
Electrons are not accommodated in a given shell unless the shells inside it are filled. In other words, the shells are filled in a stepwise manner, from the inside outwards.
The outermost shell of an atom cannot accommodate more than 8 electrons, even if it has the capacity to hold more.
The outermost shell of a chemically stable atom holds a maximum of 8 electrons — except in hydrogen and helium, which have only one shell and can therefore hold a maximum of 2 electrons.
Rule 3 sounds like a contradiction of rule 1, so look at what actually happens. Calcium has 20 electrons, and its arrangement is
K = 2, L = 8, M = 8, N = 2
The M shell has a capacity of 18 electrons, yet only 8 are put into it before the N shell starts. The reason is that an atom becomes stable once it has 8 electrons in its outermost shell, and the outermost shell will not take a ninth.
4.11.2 Working Out An Electronic Configuration
Take magnesium. It has atomic number 12 and atomic mass 24, so its nucleus holds 12 protons and \(24 - 12 = 12\) neutrons, and 12 electrons surround it. Those 12 electrons are allotted to the shells in turn, as shown in Fig. 4.11-1.
K shell, or I shell = 2 electrons L shell, or II shell = 8 electrons M shell, or III shell = 2 electrons
So the electronic configuration of magnesium is 2, 8, 2.
Diagram
Fig. 4.11-1 — The Structure Of A Magnesium Atom Small shaded central circle labelled 12p, 12n, surrounded by three concentric circles; place 2 dots on the innermost circle, 8 dots evenly spaced on the middle circle and 2 dots on the outermost circle; label the shells K, L and M and write the configuration 2, 8, 2 below.
For the lighter elements, up to argon with atomic number 18, each inner shell is completely filled before any electron occupies an outer shell. Beyond argon the situation changes. Although the third shell can accommodate up to 18 electrons, the fourth shell begins to fill once the third has 8 in it — because at that moment the third shell is the outermost one.
Potassium, atomic number 19, shows this. Its 19 electrons are arranged as
K = 2, L = 8, M = 8, N = 1
4.11.3 Electron Arrangements Of The First Twenty Elements
Applying the same three steps to each element in turn gives Table 4.5. Read across a row to see one atom; read down the last columns to see a pattern building up.
Table 4.5 — Arrangement Of Electrons In The Atoms Of Elements With Atomic Numbers 1 To 20
| Element | Symbol | Atomic no. (Z) | Mass no. (A) | No. of neutrons (A − Z) | K shell | L shell | M shell | N shell |
|---|---|---|---|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | — | 1 | |||
| Helium | He | 2 | 4 | 2 | 2 | |||
| Lithium | Li | 3 | 7 | 4 | 2 | 1 | ||
| Beryllium | Be | 4 | 9 | 5 | 2 | 2 | ||
| Boron | B | 5 | 11 | 6 | 2 | 3 | ||
| Carbon | C | 6 | 12 | 6 | 2 | 4 | ||
| Nitrogen | N | 7 | 14 | 7 | 2 | 5 | ||
| Oxygen | O | 8 | 16 | 8 | 2 | 6 | ||
| Fluorine | F | 9 | 19 | 10 | 2 | 7 | ||
| Neon | Ne | 10 | 20 | 10 | 2 | 8 | ||
| Sodium | Na | 11 | 23 | 12 | 2 | 8 | 1 | |
| Magnesium | Mg | 12 | 24 | 12 | 2 | 8 | 2 | |
| Aluminium | Al | 13 | 27 | 14 | 2 | 8 | 3 | |
| Silicon | Si | 14 | 28 | 14 | 2 | 8 | 4 | |
| Phosphorus | P | 15 | 31 | 16 | 2 | 8 | 5 | |
| Sulphur | S | 16 | 32 | 16 | 2 | 8 | 6 | |
| Chlorine | Cl | 17 | 35, 37 | 18, 20 | 2 | 8 | 7 | |
| Argon | Ar | 18 | 40 | 22 | 2 | 8 | 8 | |
| Potassium | K | 19 | 39 | 20 | 2 | 8 | 8 | 1 |
| Calcium | Ca | 20 | 40 | 20 | 2 | 8 | 8 | 2 |
Two things in the table are worth noticing now, because the rest of the chapter turns on them. First, the last shell fills up from 1 to 8 and then a new shell begins — helium, neon and argon are the atoms at which a shell is complete. Second, chlorine is listed with two mass numbers, 35 and 37, and so with two different neutron counts. We shall return to that in Section 4.14.
The same information is drawn out atom by atom in Fig. 4.11-2.
Diagram
Fig. 4.11-2 — Orbital Diagrams Of The Atoms Of The First Twenty Elements A grid of twenty labelled panels, five rows of four, in order of atomic number from hydrogen to calcium. In each panel draw a shaded central circle carrying the number of protons and neutrons, for example 12p 12n for magnesium, surrounded by the correct number of concentric shell circles with the correct number of dots on each. Head each panel with the element name and its symbol written as mass number over atomic number, and print the electronic configuration below the drawing.
4.12 Valence Electrons, Ions And Valency
Look again at Table 4.5 and cover every column except the last one filled in each row. What is left — the number of electrons in the outermost shell — turns out to decide almost everything about how an element behaves.
4.12.1 The Valence Shell And Valence Electrons
The outermost shell of an atom is called its valence shell, and the electrons present in the valence shell are known as valence electrons.
Hydrogen has only one electron, in its only shell, so it has one valence electron. Carbon, with the configuration 2, 4, has four electrons in its valence shell and so has four valence electrons.
The chemical properties of an element are decided by its valence electrons, because these are the electrons that take part in chemical reactions. The inner electrons are held close to the nucleus and are not disturbed when atoms react.
This has a direct consequence. Elements with the same number of valence electrons show similar chemical properties, and elements with different numbers of valence electrons behave differently. Sodium and potassium each have one valence electron, and both react vigorously with water; neither behaves anything like sulphur, which has six.
4.12.2 Formation Of Cations And Anions
Elements with 1, 2 or 3 valence electrons — hydrogen and helium excepted — are metals. A metal atom can lose its valence electrons, and when it does, it is left with more protons than electrons.
Sodium shows this plainly. Its configuration is 2, 8, 1; losing the single outer electron leaves 2, 8.
sodium atom − one electron → sodium ion
\[\mathrm{Na} - e^{-} \rightarrow \mathrm{Na^{+}}\]
Diagram
Fig. 4.12-1 — Formation Of A Sodium Ion Two labelled orbital drawings joined by an arrow marked − e⁻. On the left, sodium atom (2, 8, 1): shaded centre marked 11p, 12n with three shells carrying 2, 8 and 1 electrons. On the right, sodium ion (2, 8): the same shaded centre marked 11p, 12n with two shells carrying 2 and 8 electrons, the whole drawing enclosed in square brackets with a superscript plus sign outside.
The particle produced is charged, and a charged particle of this kind is called an ion.
A cation is smaller than the atom it came from. The atom has lost a whole shell of electrons, so what is left is drawn in more tightly by the same nuclear charge.
Because a metal atom hands its electrons over, metals are good reducing agents — a reducing agent being a substance that donates electrons to something else. We shall come back to this in Section 4.15.4.
Elements with 4, 5, 6 or 7 electrons in their valence shells are non-metals. These atoms gain electrons instead, and so form negatively charged ions. Chlorine, 2, 8, 7, takes one electron to become 2, 8, 8.
chlorine atom + one electron → chloride ion
\[\mathrm{Cl} + e^{-} \rightarrow \mathrm{Cl^{-}}\]
Diagram
Fig. 4.12-2 — Formation Of A Chloride Ion Two labelled orbital drawings joined by an arrow marked + e⁻. On the left, chlorine atom (2, 8, 7): shaded centre marked 17p, 18n with three shells carrying 2, 8 and 7 electrons. On the right, chloride ion (2, 8, 8): the same shaded centre with three shells carrying 2, 8 and 8 electrons, the added electron drawn as a cross, the whole drawing enclosed in square brackets with a superscript minus sign outside.
An anion is larger than the atom it came from. The same nuclear charge now has to hold more electrons, so its grip on each one is weaker and the electron cloud spreads out. Since a non-metal atom takes electrons, non-metals are good oxidising agents — an oxidising agent being a substance that accepts electrons.
Whatever happens, the nucleus is untouched. When an atom changes into an ion:
- the number of protons remains the same;
- the number of neutrons remains the same;
- only the number of electrons changes.
4.12.3 Valency From The Valence Electrons
The number of electrons an atom deals with in this way is its valency.
Working out a valency from a configuration takes one decision: is it quicker for this atom to give electrons away, or to take them in?
- An element with 1, 2 or 3 valence electrons loses them, so its valency is 1, 2 or 3 respectively.
- An element with 4, 5, 6 or 7 valence electrons gains electrons instead, and needs \(8 -\) (valence electrons) of them. Its valency is therefore 4, 3, 2 or 1 respectively.
Applying the same reasoning down the list of the first twenty elements gives Table 4.7.
Table 4.7 — Valence Electrons, Valency And Classification Of The First Twenty Elements
| Element | Symbol | Valence electrons | Valency | Metal / non-metal / noble gas |
|---|---|---|---|---|
| Hydrogen | H | 1 | 1 | Non-metal |
| Helium | He | 2 | 0 | Noble gas |
| Lithium | Li | 1 | 1 | Metal |
| Beryllium | Be | 2 | 2 | Metal |
| Boron | B | 3 | 3 | Non-metal (solid) |
| Carbon | C | 4 | 4 | Non-metal (solid) |
| Nitrogen | N | 5 | 3 | Non-metal (gas) |
| Oxygen | O | 6 | 2 | Non-metal (gas) |
| Fluorine | F | 7 | 1 | Non-metal (gas) |
| Neon | Ne | 8 | 0 | Noble gas |
| Sodium | Na | 1 | 1 | Metal (solid) |
| Magnesium | Mg | 2 | 2 | Metal (solid) |
| Aluminium | Al | 3 | 3 | Metal (solid) |
| Silicon | Si | 4 | 4 | Non-metal (solid) |
| Phosphorus | P | 5 | 3 | Non-metal (solid) |
| Sulphur | S | 6 | 2 | Non-metal (solid) |
| Chlorine | Cl | 7 | 1 | Non-metal (gas) |
| Argon | Ar | 8 | 0 | Noble gas |
| Potassium | K | 1 | 1 | Metal (solid) |
| Calcium | Ca | 2 | 2 | Metal (solid) |
Three entries in the valency column stand out: helium, neon and argon, all with a valency of zero. They combine with nothing. Why that should be so is the subject of the next section.
4.13 The Reason For The Chemical Activity Of An Atom
Sodium reacts the moment it touches water. Argon, sealed inside a light bulb, does nothing at all for years. Both are made of atoms with nuclei and shells, so the difference must lie in the one thing that varies between them — the arrangement of the valence electrons.
4.13.1 The Octet And The Duplet
Begin with the atoms that do nothing. Six inert gases occur among the naturally found elements: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). They are also called the noble or rare gases. They combine neither with other atoms nor among themselves, so their combining capacity — their valency — is zero. Table 4.8 shows how their electrons are arranged.
Table 4.8 — Electronic Configurations Of The Inert Gases
| Inert gas | Atomic no. | K | L | M | N | O | P |
|---|---|---|---|---|---|---|---|
| Helium | 2 | 2 | |||||
| Neon | 10 | 2 | 8 | ||||
| Argon | 18 | 2 | 8 | 8 | |||
| Krypton | 36 | 2 | 8 | 18 | 8 | ||
| Xenon | 54 | 2 | 8 | 18 | 18 | 8 | |
| Radon | 86 | 2 | 8 | 18 | 32 | 18 | 8 |
Read the last filled figure in each row. Every one of these atoms, apart from helium, has 8 electrons in its outermost shell. This arrangement of eight is called an octet.
Helium is the exception, and its exception is not a special favour. Helium has only one shell, the K shell, and by the \(2n^{2}\) rule that shell can hold a maximum of 2 electrons. With 2 electrons the shell is entirely full. This arrangement of two is called a duplet, and it makes helium just as stable as the others.
An atom whose valence shell is already complete has nothing to gain and nothing to lose. There is no room for another electron and no advantage in giving one up. That is precisely why the inert gases are chemically inactive.
So the chemical activity of an atom depends on the number of electrons in its valence shell. Chemically active atoms have an incomplete octet — fewer than 8 electrons in the valence shell, or fewer than 2 in the case of hydrogen. It is the valence electrons of atoms that bring about the formation of molecules.
4.13.2 Why Atoms Combine
In 1916 Kossel and Lewis, working independently, put forward the octet theory of valency, also called the electronic theory of valency. Their starting point was the observation we have just made: the rare gases have a stable electronic configuration, and nothing else does.
Their conclusion follows directly from it. The atoms of all elements other than the inert gases combine because they have incomplete valence shells, and tend to attain a stable configuration.
An atom does this by reaching the electronic configuration of the inert gas nearest to it — that is, 8 electrons in its valence shell. This is known as the octet rule. For helium, and for atoms that need to match helium, the valence shell holds 2 electrons instead, and this is the duplet rule.
An atom has three ways of getting there — by gaining electrons, by losing electrons, or by sharing electrons with another atom, of the same element or of a different one. This is the origin of valency:
The number of electrons gained, lost or shared by an atom in attaining the octet in its outermost shell gives the combining capacity of the element, that is, its valency.
4.13.3 The Chemical Bond
Whichever route is taken, electrons are redistributed between the atoms. As a result of that redistribution, a force of attraction develops between the atoms and binds them together into a molecule. This force is called a chemical bond.
The chemical combination of atoms therefore involves a redistribution of electrons that leaves each atom with a stable electronic configuration. From here the chapter divides in two. When electrons are transferred outright from one atom to another, the result is an electrovalent bond, treated in Section 4.15. When electrons are shared between atoms instead, the result is a covalent bond, treated in Section 4.16.
4.14 Isotopes
Dalton held that the atoms of an element are alike in all respects. Table 4.5 already contained a quiet contradiction of that: chlorine was listed with two mass numbers, 35 and 37. It is time to explain it.
4.14.1 What Isotopes Are
Two or more atoms of the same element may contain the same number of protons but differ in the number of neutrons. Since the number of protons is unchanged, the number of electrons is unchanged too, and the electrons are arranged in exactly the same way. Such atoms therefore have identical chemical properties while having different masses.
Atoms of an element that show similar chemical properties but possess different masses are said to show isotopy, and the varieties of such atoms are called isotopes.
Since \(A = n + p\) and the number of protons is fixed for an element, the only thing that can differ between isotopes is the number of neutrons. Some elements have several isotopes, as Table 4.9 shows.
Table 4.9 — Some Elements And Their Isotopes
| Element | Number of isotopes | Symbols |
|---|---|---|
| Hydrogen | Three | \(^{1}_{1}\mathrm{H}\), \(^{2}_{1}\mathrm{H}\), \(^{3}_{1}\mathrm{H}\) |
| Carbon | Three | \(^{12}_{6}\mathrm{C}\), \(^{13}_{6}\mathrm{C}\), \(^{14}_{6}\mathrm{C}\) |
| Chlorine | Two | \(^{35}_{17}\mathrm{Cl}\), \(^{37}_{17}\mathrm{Cl}\) |
| Oxygen | Two | \(^{16}_{8}\mathrm{O}\), \(^{18}_{8}\mathrm{O}\) |
| Potassium | Two | \(^{39}_{19}\mathrm{K}\), \(^{41}_{19}\mathrm{K}\) |
4.14.2 Structures Of The Isotopes
4.14.2.1 Isotopes Of Hydrogen
Hydrogen has three isotopes, named protium, deuterium (or heavy hydrogen) and tritium. Each has the same atomic number, 1, but the mass numbers are 1, 2 and 3 respectively. Since isotopes differ only in their number of neutrons, protium has no neutron, deuterium has one and tritium has two. They are written as \(^{1}_{1}\mathrm{H}\), \(^{2}_{1}\mathrm{H}\) and \(^{3}_{1}\mathrm{H}\), the subscript giving the atomic number and the superscript the mass number.
Diagram
Fig. 4.14-1 — Isotopes Of Hydrogen Three separate circular drawings side by side, each with one shell carrying a single electron marked e⁻. Label the nuclei 1p for protium, 1p 1n for deuterium and 1p 2n for tritium, and print the names and the symbols ¹₁H, ²₁H and ³₁H beneath the drawings.
Table 4.10 — Particles In The Three Isotopes Of Hydrogen
| Number of particles | Protium | Deuterium | Tritium |
|---|---|---|---|
| Number of electrons | 1 | 1 | 1 |
| Number of protons | 1 | 1 | 1 |
| Atomic number | 1 | 1 | 1 |
| Number of neutrons | 0 | 1 | 2 |
| Mass number | 1 | 2 | 3 |
Every row but two is identical across the three columns, and those two — neutrons and mass number — are the whole story of isotopy.
4.14.2.2 Isotopes Of Carbon
Carbon, atomic number 6, has three isotopes, of mass numbers 12, 13 and 14. They are written \(^{12}_{6}\mathrm{C}\), \(^{13}_{6}\mathrm{C}\) and \(^{14}_{6}\mathrm{C}\), or as C-12, C-13 and C-14. All three have 6 protons and 6 electrons; the numbers of neutrons are 6, 7 and 8 respectively.
Diagram
Fig. 4.14-2 — Structures Of The Isotopes Of Carbon Three separate drawings side by side, each with a shaded nucleus and two shells labelled K and L carrying 2 and 4 electrons respectively. Mark the nuclei 6p 6n, 6p 7n and 6p 8n, and print the symbols ¹²₆C, ¹³₆C and ¹⁴₆C beneath them.
4.14.2.3 Isotopes Of Chlorine
The two isotopes of chlorine, of mass numbers 35 and 37, have the same atomic number, 17 — that is, the same number of protons and the same number of electrons. They differ only in the number of neutrons in their nuclei.
For \(^{35}_{17}\mathrm{Cl}\): number of neutrons \(= 35 - 17 = 18\). For \(^{37}_{17}\mathrm{Cl}\): number of neutrons \(= 37 - 17 = 20\).
Diagram
Fig. 4.14-3 — Structures Of The Isotopes Of Chlorine Two drawings side by side, each with a shaded nucleus and three shells labelled K, L and M carrying 2, 8 and 7 electrons respectively. Mark the nuclei 17p 18n and 17p 20n, and print the symbols ³⁵₁₇Cl and ³⁷₁₇Cl beneath them.
4.14.3 Properties Of Isotopes
Isotopes have similar chemical properties. They have the same atomic number, so the same number of protons and electrons, and hence the same electronic configuration. Since chemical properties are decided by the electronic configuration of an atom, the isotopes of an element are chemically alike. A chlorine atom of mass number 37 undergoes exactly the same reactions as a chlorine atom of mass number 35.
Isotopes differ in a few physical properties, such as density and boiling point. Physical constants depend on the mass of the atom, and isotopes have different mass numbers because they have different numbers of neutrons. The boiling point of protium is 20.38 K, while that of deuterium is 23.5 K. The heavier the isotope of an element, the higher its boiling point.
Isotopes such as tritium and carbon-14 show a further physical difference: they are radioactive. The extra neutrons in their nuclei make those nuclei unstable, so the nuclei break up spontaneously, emitting certain kinds of radiation. Carbon-14 and tritium are therefore called radioisotopes.
4.14.4 Fractional Atomic Masses
Mass numbers are always whole numbers, because you cannot have part of a proton or part of a neutron. Yet the atomic mass of chlorine is given as 35.5. An atom of mass 35.5 does not exist — so where does the figure come from?
It comes from isotopes. The atomic mass of an element is the weighted average of the masses of all its naturally occurring isotopes, in the proportions in which they actually occur. Any natural sample of chlorine contains 75 per cent of \(^{35}_{17}\mathrm{Cl}\) and 25 per cent of \(^{37}_{17}\mathrm{Cl}\) — that is, the two are in the ratio 3 : 1.
4.14.5 Uses Of Isotopes
- The isotope of cobalt \(^{60}_{27}\mathrm{Co}\) is used in radiotherapy for treating cancer and other diseases.
- \(^{14}_{6}\mathrm{C}\) is used for determining the age of historical and geological material.
- An isotope of uranium, \(^{235}_{92}\mathrm{U}\), is used as a fuel in nuclear reactors.
- An isotope of iodine, \(^{131}_{53}\mathrm{I}\), is used in the treatment of goitre and other disorders of the thyroid gland.
- Radioactive isotopes are used in industry to detect leakage in underground oil pipelines, gas pipelines and water pipes.
Use 5 is worth a moment’s thought, because it shows how the two kinds of property work together. A radioisotope added to the oil behaves chemically exactly like the rest of the oil, so it travels wherever the oil travels; but it announces its position through the ground by its radiation. A detector run along the surface finds the leak without a single trench being dug.
4.14.6 Isobars
Isotopes have the same atomic number and different mass numbers. The reverse arrangement also occurs.
Since the properties of an element depend on its atomic number, and the atomic numbers here are different, isobars are different elements altogether. They have both different physical properties and different chemical properties.
Examples: \(^{40}_{18}\mathrm{Ar}\) and \(^{40}_{20}\mathrm{Ca}\); and \(^{54}_{24}\mathrm{Cr}\) and \(^{54}_{26}\mathrm{Fe}\).
4.15 The Electrovalent (Ionic) Bond
The first of the two routes to a complete octet is the blunt one: one atom hands its spare electrons over to another, and keeps none of them.
4.15.1 Formation Of An Electrovalent Bond
Atoms of metallic elements have 1, 2 or 3 valence electrons. Atoms of non-metallic elements have 5, 6 or 7. Put the two together and the arrangement suits both sides — the metal has exactly the electrons the non-metal is short of. The metal atom loses its valence electrons to the non-metal atom, and after the transfer both atoms have the electronic configuration of the nearest inert gas.
A metallic element whose atom readily loses electrons to form a positively charged ion is called an electropositive element. A non-metallic element whose atom readily accepts electrons to form a negatively charged ion is called an electronegative element. Section 4.12.2 showed both changes happening: sodium losing an electron to become a cation, chlorine gaining one to become an anion.
Now the two ions are in the same place and carry opposite charges. They attract each other, and that attraction is the bond.
Bonds formed between metals and non-metals are ionic or electrovalent bonds. Some familiar examples are sodium chloride (\(\mathrm{NaCl}\)), magnesium chloride (\(\mathrm{MgCl_2}\)), calcium oxide (\(\mathrm{CaO}\)), potassium bromide (\(\mathrm{KBr}\)) and calcium chloride (\(\mathrm{CaCl_2}\)). Caesium fluoride, \(\mathrm{CsF}\), is the most ionic compound of all, because caesium has the greatest tendency of any element to lose an electron and fluorine the greatest tendency to gain one.
Features Common To Ionic Bonding
- Metal atoms always lose their outer electrons to form positive ions.
- Non-metal atoms, with the single exception of hydrogen, always gain electrons to become negative ions.
- In both cases the ion formed has a more stable electron arrangement — that of the noble gas nearest to the element concerned.
- Ionic bonds result from the attraction between oppositely charged ions.
Why Are Ionic Compounds Stable?
An ionic compound is packed with ions, and ions of like charge repel one another, so there is a repulsive force present as well. But the electrostatic force of attraction between opposite charges is much stronger than the repulsion between like charges, and it is the balance between them, weighted heavily towards attraction, that makes ionic compounds stable.
4.15.2 Electron Dot (Lewis) Symbols
Drawing every shell of every atom is slow, and most of what is drawn takes no part in the bonding. Only the outermost electrons matter. Lewis therefore introduced a shorthand in which only the valence electrons are shown.
The electron dot symbol for an atom consists of the symbol of the element surrounded by dots representing the electrons of its outermost shell. Paired electrons are drawn as a pair of dots; an unpaired electron is drawn as a single dot. The electron dot symbol of hydrogen is one dot beside the H; that of oxygen is six dots around the O, arranged as two pairs and two singles.
Symbols other than dots — crosses or small circles — may be used, so that the electrons contributed by different atoms can be told apart in the same diagram. If dots stand for the outermost electrons of nitrogen (2, 5) and crosses for the electrons of hydrogen, then the ammonia molecule, \(\mathrm{NH_3}\), can be drawn with its nitrogen electrons and hydrogen electrons clearly distinguished. The bonding in ammonia itself is treated in Section 4.16.5.7; here the diagram is only an illustration of the notation.
In an ionic compound the transfer of an outermost electron is shown by a curved arrow running from the atom that loses it to the atom that gains it.
4.15.3 Structures Of Some Electrovalent Compounds
4.15.3.1 Sodium Chloride
The electronic configuration of a sodium atom is 2, 8, 1. It has one electron in excess of the stable configuration of the nearest noble gas, neon (2, 8). So an atom of sodium tends to give up the electron from its outermost shell:
\[\mathrm{Na} \rightarrow \mathrm{Na^{+}} + 1e^{-}\]
with the configuration changing from 2, 8, 1 in the atom to 2, 8 in the cation. After giving up one electron the sodium atom is no longer electrically neutral: eleven protons remain in its nucleus but only ten electrons revolve around it, so it carries a net positive charge of \(+1\). This is the sodium ion, \(\mathrm{Na^{+}}\), and its configuration is that of neon.
The electronic configuration of chlorine is 2, 8, 7 — one electron less than the nearest noble gas, argon (2, 8, 8). The chlorine atom therefore tends to acquire an electron:
\[\mathrm{Cl} + 1e^{-} \rightarrow \mathrm{Cl^{-}}\]
with the configuration changing from 2, 8, 7 to 2, 8, 8. A chlorine atom is electrically neutral, with 17 protons and 17 electrons. After acquiring an electron from the sodium atom it has one electron more than the number of protons in its nucleus, and so a charge of \(-1\). This is the chloride ion, \(\mathrm{Cl^{-}}\), and it now has an octet resembling argon.
So when an atom of sodium combines with an atom of chlorine, one electron is transferred from the sodium atom to the chlorine atom, and sodium chloride is formed. Fig. 4.15-1 shows the transfer.
Diagram
Fig. 4.15-1 — Orbit Structure Of Electrovalent Bonding In Sodium Chloride Left: sodium atom (2, 8, 1) and chlorine atom (2, 8, 7) drawn as shell diagrams, sodium’s electrons as crosses and chlorine’s as circles, with a curved arrow running from sodium’s single outer electron to the outer shell of chlorine. Right: sodium ion, drawn with shells 2, 8, enclosed in square brackets with a plus sign outside, and chloride ion, drawn with shells 2, 8, 8, enclosed in square brackets with a minus sign outside; label the two ions sodium ion (2, 8) and chloride ion (2, 8, 8) and mark the arrows between them coulomb forces.
The cation \(\mathrm{Na^{+}}\) and the anion \(\mathrm{Cl^{-}}\) are then drawn towards each other by their opposite charges — the coulomb force between them — and an ionic compound is formed.
4.15.3.2 Magnesium Chloride
Magnesium, atomic number 12, has the configuration 2, 8, 2; chlorine, atomic number 17, has 2, 8, 7. Magnesium has two valence electrons and chlorine has seven. The magnesium atom reaches a stable arrangement of 8 electrons by losing both outermost electrons, becoming the magnesium ion \(\mathrm{Mg^{2+}}\):
\[\mathrm{Mg} - 2e^{-} \rightarrow \mathrm{Mg^{2+}}\]
with the configuration changing from 2, 8, 2 to 2, 8.
But a chlorine atom, holding 7 electrons in its outermost shell, can accept only 1 of the 2 electrons that a magnesium atom donates. So for every magnesium atom forming a magnesium ion, there must be two chlorine atoms to form two chloride ions:
\[2\mathrm{Cl} + 2e^{-} \rightarrow 2\mathrm{Cl^{-}}\]
each going from 2, 8, 7 to 2, 8, 8. The ratio of magnesium ions to chloride ions in magnesium chloride must therefore be 1 : 2, and the formula of the compound is \(\mathrm{MgCl_2}\). This is shown in Fig. 4.15-2.
Diagram
Fig. 4.15-2 — Orbit Structure Of The Formation Of Magnesium Chloride Left: magnesium atom (2, 8, 2) with a chlorine atom (2, 8, 7) drawn above it and another below it, with curved arrows carrying one outer electron of magnesium to each chlorine. Right: magnesium ion drawn with shells 2, 8 in square brackets carrying a 2+ sign, and two chloride ions each drawn with shells 2, 8, 8 in square brackets each carrying a minus sign; label the attraction between them coulomb force, and label the groups magnesium atom (2, 8, 2), two chlorine atoms (2, 8, 7), magnesium ion (2, 8) and two chloride ions (2, 8, 8).
4.15.3.3 Calcium Oxide
Calcium, atomic number 20, has the configuration 2, 8, 8, 2; oxygen, atomic number 8, has 2, 6. A calcium atom has 2 valence electrons and an oxygen atom has 6, so oxygen requires 2 electrons to attain the octet — the configuration of neon (2, 8).
In the presence of oxygen, each calcium atom loses its 2 valence electrons. The calcium atom forms an ion of charge \(+2\) and the oxygen atom an ion of charge \(-2\):
\[\mathrm{Ca} - 2e^{-} \rightarrow \mathrm{Ca^{2+}}\]
going from 2, 8, 8, 2 to 2, 8, 8; and
\[\mathrm{O} + 2e^{-} \rightarrow \mathrm{O^{2-}}\]
going from 2, 6 to 2, 8.
Here one oxygen atom accepts both of the electrons donated by one calcium atom, so only one atom of each is needed. The formula of calcium oxide is therefore \(\mathrm{CaO}\), and not \(\mathrm{Ca_2O_2}\). Fig. 4.15-3 shows the transfer, and Fig. 4.15-4 gives the same three compounds in the electron dot shorthand.
Diagram
Fig. 4.15-3 — Orbit Structure Of The Formation Of Calcium Oxide Left: calcium atom drawn with shells 2, 8, 8, 2, its outer electrons as crosses, beside an oxygen atom drawn with shells 2, 6, its electrons as circles, with two curved arrows carrying both outer electrons of calcium into the outer shell of oxygen. Right: calcium ion with shells 2, 8, 8 in square brackets carrying a 2+ sign, and oxide ion with shells 2, 8 in square brackets carrying a 2− sign, with the attraction between them labelled coulomb force.
Diagram
Fig. 4.15-4 — Electron Dot Structures Of Three Electrovalent Compounds Three short rows, one below the other. Row 1, sodium chloride: Na with a single dot and Cl with seven electrons, a curved arrow from the sodium dot to chlorine, giving Na⁺ and Cl⁻ with the chloride shown in square brackets with eight electrons. Row 2, magnesium chloride: Mg with two dots between two Cl atoms each with seven electrons, giving Mg²⁺ and two chloride ions each in square brackets with eight electrons. Row 3, calcium oxide: Ca with two dots beside O with six electrons, giving Ca²⁺ and an oxide ion in square brackets with eight electrons and a 2− sign. Use dots for the metal’s electrons and circles for the non-metal’s throughout.
4.15.4 Electrovalent Bond Formation As A Redox Change
Look at what happens to each partner separately. In the formation of an electrovalent bond, the electropositive atom loses electrons and undergoes oxidation, while the electronegative atom gains electrons and undergoes reduction. A change in which both happen together is called a redox reaction.
Take the formation of sodium chloride from its elements:
sodium + chlorine → sodium chloride
\[2\mathrm{Na} + \mathrm{Cl_2} \rightarrow 2\mathrm{Na^{+}} + 2\mathrm{Cl^{-}}\]
The same reaction can be written as two half reactions:
\[2\mathrm{Na} \rightarrow 2\mathrm{Na^{+}} + 2e^{-} \qquad \text{(oxidation)}\]
\[\mathrm{Cl_2} + 2e^{-} \rightarrow 2\mathrm{Cl^{-}} \qquad \text{(reduction)}\]
Oxidation and reduction always occur simultaneously, and the reason is simply that electrons have to go somewhere. The electrons lost by one substance must be gained by the other, in the same number, at the same moment.
This works for reactions in solution too. Drop a strip of zinc into blue copper sulphate solution and a red-brown deposit forms on the zinc while the blue colour fades:
zinc + copper sulphate solution → zinc sulphate solution + copper
\[\mathrm{Zn}\text{(s)} + \mathrm{CuSO_4}\text{(aq)} \rightarrow \mathrm{ZnSO_4}\text{(aq)} + \mathrm{Cu}\text{(s)}\]
Stripping away the sulphate, which takes no part, leaves the electron transfer itself:
\[\mathrm{Zn} + \mathrm{Cu^{2+}} \rightarrow \mathrm{Zn^{2+}} + \mathrm{Cu}\]
\[\mathrm{Cu^{2+}} + 2e^{-} \rightarrow \mathrm{Cu} \qquad \text{(reduction)}\]
\[\mathrm{Zn} - 2e^{-} \rightarrow \mathrm{Zn^{2+}} \qquad \text{(oxidation)}\]
Zinc, acting as the reducing agent, reduces the copper(II) ion to copper. At the same time the copper(II) ion, acting as the oxidising agent, oxidises zinc to the zinc(II) ion. The net change is a redox reaction.
From this a general statement follows: an oxidising agent is an acceptor of electrons, and a reducing agent is a donor of electrons.
4.15.5 Properties Of Electrovalent Compounds
Electrovalent compounds have strong electrostatic forces of attraction between their ions, and these cannot be separated easily. Every property below follows from that one fact.
- They are hard, brittle, crystalline solids consisting of ions.
- They are non-volatile, with high melting points and high boiling points.
- They do not conduct electricity in the solid state, but they are good conductors of electricity in the fused state or in aqueous solution.
The last property is the most revealing. An electric current through a liquid is carried by moving ions. In a solid ionic compound the ions are locked in place in the crystal and cannot move, so nothing is carried. Melt the solid, or dissolve it in water, and the ions are set free to move — and the compound conducts at once.
4.16 The Covalent (Molecular) Bond
Transfer works when one atom is willing to give and the other to take. But what happens when both atoms want to take, and neither will give? Two chlorine atoms each need one more electron; neither will part with any. They settle it by sharing.
4.16.1 Formation Of A Covalent Bond
The atoms of non-metals usually have 5, 6 or 7 electrons in their outermost shell — carbon, with 4, and hydrogen, with just 1, being the exceptions. The atoms of such elements do not favour the loss of their electrons, for reasons of energy, and so transfer between them is not possible. These atoms can complete their octets only by mutually sharing one or more pairs of electrons, and each atom contributes an equal number of electrons to the sharing.
So whenever a non-metal combines with another non-metal to attain a stable configuration, the sharing of electrons takes place between their atoms and a covalent bond is formed.
Hydrogen is a non-metal and chlorine is a non-metal. When they combine to form hydrogen chloride, electrons are shared between the hydrogen and chlorine atoms and a covalent bond is formed. A covalent bond can equally well be formed between two atoms of the same non-metal: two chlorine atoms combine by sharing one electron each, forming a chlorine molecule, \(\mathrm{Cl_2}\).
Two summarising statements are worth holding on to:
- A covalent bond is formed when both combining atoms have four or more electrons in their outermost shells, that is, between non-metals — hydrogen, beryllium and boron being exceptions to the electron count.
- Hydrogen can combine with all non-metals by means of covalent bonds.
4.16.2 Single, Double And Triple Covalent Bonds
Covalent bonds are of three types, according to how many pairs of electrons are shared.
A single covalent bond is formed by the sharing of one pair of electrons between the atoms, each atom contributing one electron. It is denoted by a short line, —, drawn between the two atoms, so a hydrogen molecule may be written as H—H.
Molecules of chlorine, hydrogen chloride, water, ammonia, methane and carbon tetrachloride are all examples of single covalent bonds. Their structures are drawn in Fig. 4.16-1.
Diagram
Fig. 4.16-1 — Structural Formulae Of Some Molecules With Single Covalent Bonds Six small structures arranged in two rows of three, each drawn with plain lines for bonds and labelled beneath. Row 1: Cl—Cl labelled chlorine; H—Cl labelled hydrogen chloride; H—O—H labelled water. Row 2: a nitrogen atom bonded to three hydrogens, two in the line H—N—H and one below, labelled ammonia; a carbon atom bonded to four hydrogens, one above, one below, one on each side, labelled methane; a carbon atom bonded to four chlorines in the same arrangement, labelled carbon tetrachloride.
A double bond is formed by the sharing of two pairs of electrons between two atoms. It is a combination of two single bonds, so it is represented by two short lines, =. The oxygen molecule, \(\mathrm{O_2}\), contains a double bond and can be written as O=O.
A triple bond is formed by the sharing of three pairs of electrons between two atoms. It is a combination of three single bonds, so it is represented by three short lines, ≡. The nitrogen molecule, \(\mathrm{N_2}\), contains a triple bond and is written as N≡N.
Some molecules have a combination of single bonds with a double or a triple bond. The ethene molecule, \(\mathrm{C_2H_4}\), has one double covalent bond and four single covalent bonds. The ethyne molecule, \(\mathrm{C_2H_2}\), has one triple covalent bond and two single covalent bonds. Both are shown in Fig. 4.16-2.
Diagram
Fig. 4.16-2 — Structures Of Ethene And Ethyne Upper half, ethene: on the left an electron dot diagram with two carbon atoms sharing two pairs between them, each carbon also joined to two hydrogens by a shared pair, dots for carbon electrons and crosses for hydrogen electrons; on the right the structural formula with H and H on each carbon and a double line between the carbons, labelled ethene (C₂H₄). Lower half, ethyne: on the left the electron dot diagram with three shared pairs between the two carbons and one shared pair to each hydrogen; on the right the structural formula H—C≡C—H labelled ethyne (C₂H₂).
4.16.3 Covalency
So the covalency of hydrogen is 1, that of oxygen is 2, that of nitrogen is 3 and that of carbon is 4. Compare these with the valencies in Table 4.7 and they agree — the atom needs the same number of electrons whether it is going to gain them or share them.
4.16.4 Non-Polar And Polar Covalent Compounds
Sharing a pair of electrons does not always mean sharing it evenly.
Covalent compounds are non-polar when the shared pair of electrons is equally distributed between the two atoms. No charge separation takes place; the molecule is symmetrical and electrically neutral.
This is what happens when the two atoms joined by the bond are identical. The shared pair is then at an equal distance from both, and is equally attracted by the two nuclei. Hydrogen (\(\mathrm{H_2}\)), chlorine (\(\mathrm{Cl_2}\)) and oxygen (\(\mathrm{O_2}\)) are perfectly non-polar compounds for this reason.
Covalent compounds are said to be polar when the shared pair of electrons is not at an equal distance between the two atoms. One atom pulls the pair closer to itself, and this results in the development of fractional positive and negative charges on the two atoms. Such molecules ionise in water. Hydrogen chloride is the standard example.
4.16.5 Structures Of Some Covalent Molecules
Table 4.12 collects what each atom in this list has, and what it needs, before we take the molecules one at a time.
Table 4.12 — Electronic Configurations Of The Atoms In Nine Covalent Molecules
| Molecule | Atoms involved | Electronic configuration | Nearest noble gas | Electrons needed by each atom |
|---|---|---|---|---|
| Hydrogen, \(\mathrm{H_2}\) | Hydrogen | H [1] | Helium [2] | 1 |
| Chlorine, \(\mathrm{Cl_2}\) | Chlorine | Cl [2, 8, 7] | Argon [2, 8, 8] | 1 |
| Oxygen, \(\mathrm{O_2}\) | Oxygen | O [2, 6] | Neon [2, 8] | 2 |
| Nitrogen, \(\mathrm{N_2}\) | Nitrogen | N [2, 5] | Neon [2, 8] | 3 |
| Hydrogen chloride, \(\mathrm{HCl}\) | Hydrogen, chlorine | H [1]; Cl [2, 8, 7] | Helium [2]; argon [2, 8, 8] | 1; 1 |
| Water, \(\mathrm{H_2O}\) | Hydrogen, oxygen | H [1]; O [2, 6] | Helium [2]; neon [2, 8] | 1; 2 |
| Ammonia, \(\mathrm{NH_3}\) | Nitrogen, hydrogen | N [2, 5]; H [1] | Neon [2, 8]; helium [2] | 3; 1 |
| Carbon tetrachloride, \(\mathrm{CCl_4}\) | Carbon, chlorine | C [2, 4]; Cl [2, 8, 7] | Neon [2, 8]; argon [2, 8, 8] | 4; 1 |
| Methane, \(\mathrm{CH_4}\) | Carbon, hydrogen | C [2, 4]; H [1] | Neon [2, 8]; helium [2] | 4; 1 |
4.16.5.1 Hydrogen Molecule
Hydrogen is a non-polar compound. A hydrogen atom has one electron in its only shell, and it needs one more electron to attain the duplet. To meet this need, a hydrogen atom shares its electron with another hydrogen atom. One electron each is contributed by the two atoms, and the resulting pair is mutually shared by both, forming a hydrogen molecule.
hydrogen atom + hydrogen atom → hydrogen molecule
\[\mathrm{H} + \mathrm{H} \rightarrow \mathrm{H_2}\]
One shared pair of electrons means a single covalent bond, written H—H.
Diagram
Fig. 4.16-3 — Formation Of A Hydrogen Molecule Upper row, before combination: two separate circles each labelled H atom, each with one electron on its shell, the two electrons enclosed in a dotted box between them; below this the electron dot structure H• + •H. Lower row, after combination: the two circles overlapping, with the shared pair of two dots drawn in the overlap, labelled hydrogen molecule; below it H:H, and beside it one shared pair of electrons, single covalent bond.
4.16.5.2 Chlorine Molecule
Chlorine is a non-polar compound. To attain the configuration of argon, a chlorine atom needs one electron. When two chlorine atoms come closer, each contributes one electron and they form one shared pair of electrons between them. Both atoms thus attain an octet, and a single covalent bond, Cl—Cl, is formed.
\[\mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\]
Diagram
Fig. 4.16-4 — Formation Of A Chlorine Molecule Upper row, before combination: two circles labelled Cl-atom, each with seven electrons drawn on its outer shell, the two electrons that will be shared enclosed in a dotted box between them; below this the electron dot structure showing each Cl with seven electrons. Lower row, after combination: the two circles overlapping with one shared pair in the overlap, labelled chlorine molecule; below it the dot structure with the shared pair between the two Cl symbols, and the labels one shared pair of electrons, single covalent bond.
4.16.5.3 Oxygen Molecule
Oxygen is a non-polar compound. To attain the configuration of the nearest inert gas, neon, an oxygen atom needs two electrons. When two oxygen atoms come closer, each contributes two electrons, so they have two shared pairs of electrons between them. Both atoms attain an octet, and a double covalent bond, O=O, is formed.
\[\mathrm{O} + \mathrm{O} \rightarrow \mathrm{O_2}\]
Diagram
Fig. 4.16-5 — Formation Of An Oxygen Molecule Upper row, before combination: two circles labelled O-atom, each with six electrons on the outer shell, the four electrons to be shared enclosed in a dotted box between them; below this the electron dot structure of two O atoms with six electrons each. Lower row, after combination: the two circles overlapping with two shared pairs drawn in the overlap, labelled oxygen molecule; below it the dot structure with two shared pairs, and the labels two shared pairs of electrons, double covalent bond.
4.16.5.4 Nitrogen Molecule
Nitrogen is a non-polar compound. To attain the configuration of neon, a nitrogen atom needs three electrons. When two nitrogen atoms come closer, each contributes three electrons, so they have three shared pairs of electrons between them. Both atoms attain an octet, and a triple covalent bond, N≡N, is formed.
\[\mathrm{N} + \mathrm{N} \rightarrow \mathrm{N_2}\]
Diagram
Fig. 4.16-6 — Formation Of A Nitrogen Molecule Upper row, before combination: two circles labelled N-atom, each with five electrons on the outer shell, the six electrons to be shared enclosed in a dotted box between them; below this the electron dot structure of the two atoms. Lower row, after combination: the two circles overlapping with three shared pairs drawn in the overlap, labelled nitrogen molecule; below it the dot structure with three shared pairs, and the labels three shared pairs of electrons, triple covalent bond.
4.16.5.5 Hydrogen Chloride Molecule
Hydrogen chloride is a polar compound. To attain the configuration of the nearest noble gas, a hydrogen atom needs one electron and a chlorine atom also needs one electron. In the hydrogen chloride molecule, the hydrogen atom shares an electron pair with the chlorine atom, so that hydrogen acquires a duplet configuration and chlorine an octet. A single covalent bond, H—Cl, results.
\[\mathrm{H} + \mathrm{Cl} \rightarrow \mathrm{HCl}\]
One molecule of hydrogen chloride contains a total of two atoms: one atom of hydrogen and one atom of chlorine.
Diagram
Fig. 4.16-7 — Formation Of A Hydrogen Chloride Molecule Upper row, before combination: a small circle labelled H atom with one electron drawn as a cross, and a larger circle labelled chlorine atom with seven electrons drawn as dots, joined by dotted lines; below this the electron dot structure H× beside Cl with seven dots. Lower row, after combination: the two circles overlapping with the shared pair — one cross and one dot — in the overlap, labelled hydrogen chloride molecule; below it the dot structure and the form H—Cl, with the labels one shared pair of electrons, single covalent bond.
4.16.5.6 Water Molecule
Water is a polar compound. To attain the configuration of the nearest noble gas, a hydrogen atom needs one electron and an oxygen atom needs two. In a water molecule, each of the two hydrogen atoms shares an electron pair with the oxygen atom, so that each hydrogen acquires a duplet configuration and oxygen an octet. Two single covalent bonds result, written H—O—H.
\[2\mathrm{H} + \mathrm{O} \rightarrow \mathrm{H_2O}\]
One molecule of water contains a total of three atoms: one atom of oxygen and two atoms of hydrogen.
Diagram
Fig. 4.16-8 — Formation Of A Water Molecule Upper row, before combination: a central circle labelled O-atom with six electrons drawn as dots, with a small circle labelled H-atom on each side, each hydrogen electron drawn as a cross and joined to the oxygen by dotted lines; below this the electron dot structure H× + O with six dots + ×H. Lower row, after combination: the three circles overlapping in a row, with one shared pair in each overlap, labelled water molecule; below it the dot structure and the form H—O—H, with the label two single covalent bonds.
4.16.5.7 Ammonia Molecule
Ammonia is a polar compound. To attain the electronic configuration of the nearest noble gas, a nitrogen atom needs three electrons and a hydrogen atom needs one. When a molecule of ammonia is formed, one atom of nitrogen shares three electron pairs, one with each of the three atoms of hydrogen. Three single covalent bonds result.
\[\mathrm{N} + 3\mathrm{H} \rightarrow \mathrm{NH_3}\]
The ammonia molecule contains a total of four atoms: one atom of nitrogen and three atoms of hydrogen.
Diagram
Fig. 4.16-9 — Formation Of An Ammonia Molecule Upper row, before combination: a central circle labelled N with five electrons drawn as dots, with three circles labelled H-atom placed above, left and right of it, each hydrogen electron drawn as a cross and joined by dotted lines; below this the electron dot structure with H× above, and H× + N with five dots + ×H in the line. Lower row, after combination: the four circles overlapping, with one shared pair in each of the three overlaps, labelled ammonia molecule; below it the structural form with H—N—H in the line and a third H below the nitrogen, labelled three single covalent bonds.
4.16.5.8 Carbon Tetrachloride Molecule
Carbon tetrachloride is a non-polar compound. To attain the configuration of the nearest noble gas, a carbon atom needs four electrons and a chlorine atom needs one. When a molecule of carbon tetrachloride is formed, one atom of carbon shares four electron pairs, one with each of the four atoms of chlorine. Four single covalent bonds result.
\[\mathrm{C} + 4\mathrm{Cl} \rightarrow \mathrm{CCl_4}\]
One molecule of carbon tetrachloride contains five atoms in all: one atom of carbon and four atoms of chlorine.
Diagram
Fig. 4.16-10 — Formation Of A Carbon Tetrachloride Molecule Upper row, before combination: a central circle labelled C with four electrons drawn as crosses, with four circles labelled Cl-atom placed above, below, left and right, each with seven electrons drawn as dots and joined to the carbon by dotted lines; below this the corresponding electron dot structure. Lower row, after combination: the five circles overlapping, with one shared pair in each of the four overlaps, labelled carbon tetrachloride molecule; beside it the structural form with C at the centre and Cl above, below, left and right, labelled four single covalent bonds.
4.16.5.9 Methane Molecule
Methane is a non-polar compound. To attain the configuration of the nearest noble gas, a carbon atom needs four electrons and a hydrogen atom needs one. When a molecule of methane is formed, one atom of carbon shares four electron pairs, one with each of the four atoms of hydrogen. Four single covalent bonds result.
\[\mathrm{C} + 4\mathrm{H} \rightarrow \mathrm{CH_4}\]
A methane molecule contains a total of five atoms: one atom of carbon and four atoms of hydrogen.
Diagram
Fig. 4.16-11 — Formation Of A Methane Molecule Upper row, before combination: a central circle labelled C with four electrons drawn as dots, with four circles labelled H-atom placed above, below, left and right, each hydrogen electron drawn as a cross and joined by dotted lines; below this the corresponding electron dot diagram. Lower row, after combination: the five circles overlapping, with one shared pair in each of the four overlaps, labelled methane molecule; beside it the structural form with C at the centre and H above, below, left and right, labelled four single covalent bonds.
4.16.6 Properties Of Covalent Compounds
Covalent compounds are made up of molecules, and the force of attraction between one molecule and the next is weak. Every property below follows from that.
- They are either gases, or liquids, or soft solids.
- They are volatile, with low melting points and low boiling points.
- They are non-conductors of electricity in the solid, molten or aqueous state.
- On passing an electric current, non-polar covalent compounds do not ionise. Some covalent compounds are polar in nature, and these do form ions in their solutions and act as electrolytes.
- They are insoluble in water, but dissolve in organic solvents.
Notice the difference between the weak force between molecules and the strong bond inside one. Boiling a covalent liquid separates the molecules from one another; it does not break the covalent bonds, and the molecules that leave as vapour are the same molecules that were there in the liquid.
4.17 Radical Formation On Breaking A Covalent Bond
A covalent bond is a shared pair of electrons — one electron from each atom. Chlorine and methane, drawn in Fig. 4.17-1, both show this: every line in a structural formula stands for one such pair.
Diagram
Fig. 4.17-1 — Shared Pairs In Chlorine And In Methane Two rows. Row 1, chlorine (Cl₂): on the left the electron dot structure of two Cl atoms with the shared pair between them, on the right the same written as Cl—Cl. Row 2, methane (CH₄): on the left the electron dot structure with carbon’s four electrons drawn as dots and the four hydrogen electrons as crosses, forming four shared pairs; on the right the structural formula with C at the centre and an H above, below, left and right.
Now ask what happens when such a bond is broken. The pair can be split in two ways, and one of them concerns us here: the bond may break symmetrically, so that the two shared electrons are divided equally between the two atoms. This is called homolytic bond fission, or homolytic cleavage.
During homolytic fission, each atom retains one electron from the original covalent bond. The result is two neutral species, each carrying an unpaired electron.
These neutral species are called free radicals, and the presence of the unpaired electron makes them unstable, short-lived and highly reactive. A radical will seize an electron from almost anything it meets, in order to pair up the one it has.
Breaking a bond costs energy. Homolytic cleavage requires enough energy to overcome the bond dissociation energy, and this is usually supplied by heat or by high-energy radiation such as ultraviolet light. Bond breaking brought about by light is called photolysis.
Ultraviolet radiation splits the covalent bond in a chlorine molecule, giving two chlorine radicals:
chlorine molecule → two chlorine free radicals
\[\mathrm{Cl_2} \xrightarrow{\ \mathrm{UV}\ } \mathrm{Cl^{\bullet}} + \mathrm{^{\bullet}Cl}\]
Notice that neither product carries a charge. The chlorine molecule had 17 protons and 17 electrons in each half, and each half keeps exactly that — one atom, one spare electron, no charge.
4.18 Summary
- The idea of an ultimate particle of matter was first put forward in India by Kanada, who called it paramanu; the Greek philosopher Democritus called it atomos, meaning indivisible.
- John Dalton gave the first scientific atomic theory in 1808, picturing atoms as small, indivisible, indestructible particles, alike within an element and combining in simple ratios.
- An element is a pure substance that can neither be formed from, nor decomposed into, simpler substances by ordinary physical or chemical methods; it is made of only one type of atom.
- An atom is the smallest particle of an element that exhibits all the properties of that element. It may or may not exist independently, but it takes part in every chemical reaction.
- Gases conduct electricity at a very low pressure of about 0.01 mm of mercury and a high potential of about 10 000 volts, emitting cathode rays.
- Cathode rays travel in straight lines, cast shadows, turn a paddle wheel, produce a greenish yellow fluorescence, produce X-rays on striking a metal target, ionise gases, and are deflected towards a positive plate — showing that they consist of negatively charged particles.
- The charge-to-mass ratio of these particles is \(1.76\times10^{11}\ \text{C kg}^{-1}\) whatever the gas or the cathode metal, which shows that the same particle is present in all matter. Thomson concluded that cathode rays are streams of electrons, and that electrons are a part of every atom.
- An electron carries one unit of negative charge, \(-1.602\times10^{-19}\ \text{C}\), and has a mass of \(1/1837\) of that of a hydrogen atom. It is written as \(_{-1}e^{0}\). The name was suggested by G. J. Stoney in 1874.
- Goldstein produced anode rays, or canal rays, using a perforated cathode. They are positively charged, and their charge-to-mass ratio depends on the gas, being greatest for hydrogen; they arise from gas atoms that have lost electrons, not from the anode metal.
- A proton carries one unit of positive charge, \(1.602\times10^{-19}\ \text{C}\), and has a mass of 1 amu, that is \(1.672\times10^{-24}\ \text{g}\). It is written as \(_{+1}\mathrm{P^{1}}\), and is formed when a hydrogen atom loses its electron: \(\mathrm{H} - e^{-} \rightarrow \mathrm{H^{+}}\).
- Thomson’s plum pudding model of 1904 pictured the atom as a sphere of uniform positive charge with electrons embedded in it, and the mass spread evenly. It could not account for a nucleus, for the stability of the atom, for the arrangement of electrons, or for Rutherford’s results.
- In Rutherford’s gold foil experiment of 1911, most alpha particles passed straight through the foil, some were deflected slightly, and about one in ten thousand was deflected through a large angle or bounced back.
- From these observations Rutherford concluded that the atom is largely empty space, with a small, dense, positively charged nucleus at the centre in which the whole mass is concentrated, and electrons revolving around it at large distances. The atom is neutral because the number of protons equals the number of electrons.
- Rutherford’s model failed to explain the stability of the atom, since a revolving charged particle should radiate energy and spiral into the nucleus, and it did not explain how the electrons are distributed.
- Bohr’s model of 1913 states that electrons are confined to fixed orbits called shells or energy levels, that an electron neither gains nor loses energy while revolving in an orbit, and that it jumps to a higher orbit on gaining energy and falls back on losing it. The shells are labelled K, L, M, N outwards from the nucleus, K having the least energy.
- Chadwick discovered the neutron in 1932 by bombarding light nuclei such as beryllium with alpha particles: \(^{9}_{4}\mathrm{Be} + ^{4}_{2}\mathrm{He} \rightarrow ^{12}_{6}\mathrm{C} + ^{1}_{0}n\). A neutron carries no charge and has a mass of 1 amu, that is \(1.676\times10^{-24}\ \text{g}\).
- Electrons lie outside the nucleus and have negligible mass; protons and neutrons lie inside it and are together called nucleons, and they account for the whole mass of the atom.
- The atomic number, \(Z\), is the number of protons in the nucleus, which equals the number of electrons in the neutral atom and the units of positive charge on the nucleus. It is unique to each element.
- The mass number, \(A\), is the total number of protons and neutrons in the nucleus. An atom is written as \(^{A}_{Z}\mathrm{X}\), and the number of neutrons is found from \[n = A - Z\]
- The distribution of electrons in the shells is the electronic configuration. By the Bohr–Bury scheme, the maximum number of electrons in a shell is \(2n^{2}\); shells are filled stepwise from the inside outwards; and the outermost shell can hold no more than 8 electrons whatever its capacity.
- The outermost shell is the valence shell, and the electrons in it are the valence electrons. The chemical properties of an element are decided by its valence electrons.
- Elements with 1, 2 or 3 valence electrons are metals; they lose electrons to form cations, which are smaller than the atom, and they are good reducing agents. Elements with 4, 5, 6 or 7 valence electrons are non-metals; they gain electrons to form anions, which are larger than the atom, and they are good oxidising agents. When an atom becomes an ion, only the number of electrons changes.
- Valency is the combining capacity of an element, equal to the number of electrons that take part in a chemical reaction: 1, 2 or 3 for elements with 1, 2 or 3 valence electrons, and \(8 -\) (valence electrons) for those with 4, 5, 6 or 7.
- The inert or noble gases — helium, neon, argon, krypton, xenon and radon — have complete outermost shells, an octet of 8 electrons or, in helium, a duplet of 2. They cannot gain or lose electrons, so their valency is zero and they are chemically inactive.
- Kossel and Lewis put forward the octet theory of valency in 1916: atoms of all elements other than the inert gases combine because their valence shells are incomplete, and they attain the configuration of the nearest inert gas by gaining, losing or sharing electrons. This redistribution produces a force of attraction between the atoms, called a chemical bond.
- Isotopes are atoms of the same element having the same atomic number but different mass numbers, differing only in the number of neutrons. Hydrogen has protium, deuterium and tritium; carbon has C-12, C-13 and C-14; chlorine has \(^{35}_{17}\mathrm{Cl}\) and \(^{37}_{17}\mathrm{Cl}\).
- Isotopes have the same electronic configuration and therefore identical chemical properties, but differ in physical properties such as density and boiling point. Unstable isotopes such as tritium and carbon-14 are radioactive.
- The atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes, which is why it may be fractional. Chlorine, being 75 per cent \(^{35}_{17}\mathrm{Cl}\) and 25 per cent \(^{37}_{17}\mathrm{Cl}\), has an atomic mass of \(\dfrac{35 \times 3 + 37 \times 1}{4} = 35.5\)
- Isotopes are used in radiotherapy (\(^{60}_{27}\mathrm{Co}\)), in dating material (\(^{14}_{6}\mathrm{C}\)), as nuclear fuel (\(^{235}_{92}\mathrm{U}\)), in treating thyroid disorders (\(^{131}_{53}\mathrm{I}\)) and in detecting leaks in underground pipelines.
- Isobars are atoms of different elements having the same mass number but different atomic numbers, such as \(^{40}_{18}\mathrm{Ar}\) and \(^{40}_{20}\mathrm{Ca}\).
- An electrovalent, or ionic, bond is formed by the transfer of electrons from one atom to another. It forms between a metal, which is electropositive, and a non-metal, which is electronegative, and the ions produced are held together by electrostatic attraction.
- In sodium chloride one electron passes from sodium (2, 8, 1) to chlorine (2, 8, 7), giving \(\mathrm{Na^{+}}\) (2, 8) and \(\mathrm{Cl^{-}}\) (2, 8, 8). In magnesium chloride each magnesium atom gives one electron to each of two chlorine atoms, giving \(\mathrm{MgCl_2}\). In calcium oxide one calcium atom gives both its valence electrons to one oxygen atom, giving \(\mathrm{CaO}\).
- The formation of an electrovalent bond is a redox reaction: the electropositive atom loses electrons and is oxidised, the electronegative atom gains them and is reduced. \[2\mathrm{Na} \rightarrow 2\mathrm{Na^{+}} + 2e^{-} \qquad \mathrm{Cl_2} + 2e^{-} \rightarrow 2\mathrm{Cl^{-}}\] An oxidising agent is an acceptor of electrons and a reducing agent a donor of electrons.
- Electrovalent compounds are hard, brittle, crystalline solids made of ions. They are non-volatile with high melting and boiling points, and they do not conduct electricity as solids but conduct well when fused or in aqueous solution.
- A covalent, or molecular, bond is formed by the mutual sharing of one or more pairs of electrons between two combining atoms, each contributing an equal number of electrons. It forms between non-metals.
- Sharing one pair gives a single covalent bond, as in \(\mathrm{H_2}\), \(\mathrm{Cl_2}\), \(\mathrm{HCl}\), \(\mathrm{H_2O}\), \(\mathrm{NH_3}\), \(\mathrm{CCl_4}\) and \(\mathrm{CH_4}\); sharing two pairs gives a double bond, as in \(\mathrm{O_2}\); sharing three pairs gives a triple bond, as in \(\mathrm{N_2}\). Ethene has one double and four single bonds; ethyne has one triple and two single bonds.
- Covalency is the number of electrons of an atom that take part in forming shared pairs — 1 for hydrogen, 2 for oxygen, 3 for nitrogen and 4 for carbon.
- A covalent molecule is non-polar when the shared pair is equally distributed between identical atoms, as in \(\mathrm{H_2}\), \(\mathrm{Cl_2}\) and \(\mathrm{O_2}\), and polar when the shared pair lies nearer one atom, as in hydrogen chloride, so that fractional charges develop and the molecule ionises in water.
- Covalent compounds consist of molecules with weak forces between them. They are gases, liquids or soft solids, volatile with low melting and boiling points, non-conductors of electricity, insoluble in water but soluble in organic solvents; polar ones ionise in solution and act as electrolytes.
- When a covalent bond breaks symmetrically, each atom retains one electron of the shared pair. This is homolytic bond fission, and it produces free radicals — neutral species with an unpaired electron, which are unstable, short-lived and highly reactive. Ultraviolet light splits chlorine in this way: \[\mathrm{Cl_2} \xrightarrow{\ \mathrm{UV}\ } \mathrm{Cl^{\bullet}} + \mathrm{^{\bullet}Cl}\]
Flowchart
FC 4.1 — How An Atom Attains A Stable Configuration Atom with incomplete valence shell → tends to reach the configuration of the nearest inert gas → three routes. Route 1: metal atom loses valence electrons → cation → attracted to anion → electrovalent (ionic) bond → ionic compound, e.g. NaCl, MgCl₂, CaO. Route 2: non-metal atom gains electrons from a metal → anion → joins route 1. Route 3: two non-metal atoms share electron pairs → covalent bond, single, double or triple → covalent molecule, e.g. H₂, O₂, N₂, CH₄. Separate branch: atom with a complete octet or duplet (inert gas) → no bond formed → valency zero.
4.19 Key Terms
- Alpha particle — a doubly charged helium ion, \(\mathrm{He^{2+}}\), made of two protons and two neutrons, emitted by radioactive substances.
- Anion — a negatively charged ion, formed when an atom gains one or more electrons.
- Anode — the electrode joined to the positive terminal of the supply.
- Anode rays — positively charged rays travelling away from the anode in a discharge tube, made of gas atoms that have lost electrons; also called canal rays or positive rays.
- Atom — the smallest particle of an element that shows all the properties of that element.
- Atomic number (Z) — the number of protons in the nucleus of an atom, equal to the number of electrons in the neutral atom.
- Canal rays — anode rays, so called because they pass through the holes, or canals, of a perforated cathode.
- Cathode — the electrode joined to the negative terminal of the supply.
- Cathode rays — streams of electrons produced when a high voltage is passed through a gas at very low pressure in a discharge tube.
- Cation — a positively charged ion, formed when an atom loses one or more electrons.
- Chemical bond — the force of attraction that binds two atoms together as a unit called a molecule.
- Covalency — the number of electrons of an atom that take part in forming shared pairs.
- Covalent bond — a bond formed between two atoms by the mutual sharing of one or more pairs of electrons.
- Discharge tube — a hard glass tube fitted with two electrodes, in which a gas at very low pressure is subjected to a high voltage.
- Double bond — a covalent bond formed by the sharing of two pairs of electrons between two atoms.
- Duplet — an arrangement of 2 electrons in the outermost shell, the stable arrangement for an atom with only one shell.
- Duplet rule — the rule that an atom with a single shell attains stability with 2 electrons in it, matching helium.
- Electron — a sub-atomic particle carrying one unit of negative charge, with a mass of \(1/1837\) of that of a hydrogen atom.
- Electronegative element — a non-metallic element whose atom readily accepts electrons to form a negative ion.
- Electronic configuration — the distribution of the electrons of an atom among its shells, written in order from the innermost shell outwards.
- Electropositive element — a metallic element whose atom readily loses electrons to form a positive ion.
- Electrovalency — the number of electrons an atom loses or gains in forming an electrovalent bond.
- Electrovalent bond — a bond formed by the transfer of electrons from one atom to another; also called an ionic bond.
- Element — a pure substance made of only one type of atom, which cannot be formed from or broken into simpler substances by ordinary means.
- Energy level — a fixed orbit associated with a definite amount of energy, in which an electron revolves without gaining or losing energy.
- Free radical — a neutral atom or group of atoms carrying an unpaired electron, formed when a covalent bond breaks symmetrically.
- Homolytic bond fission — the symmetrical breaking of a covalent bond, in which each atom retains one electron of the shared pair.
- Ion — an electrically charged atom or group of atoms, formed by the loss or gain of electrons.
- Isobars — atoms of different elements having the same mass number but different atomic numbers.
- Isotopes — atoms of the same element having the same atomic number but different mass numbers.
- Mass number (A) — the total number of protons and neutrons in the nucleus of an atom.
- Metals — elements whose atoms have 1, 2 or 3 valence electrons and lose them to form cations.
- Neutron — a sub-atomic particle in the nucleus, carrying no charge and having a mass of about 1 amu.
- Non-metals — elements whose atoms have 4, 5, 6 or 7 valence electrons and gain or share electrons.
- Non-polar covalent compound — a covalent compound in which the shared pair is equally distributed between the two atoms, so no charge separation occurs.
- Nucleons — the particles of the nucleus, that is the protons and the neutrons taken together.
- Nucleus — the small, dense, positively charged centre of an atom in which its whole mass is concentrated.
- Octet — an arrangement of 8 electrons in the outermost shell, the stable arrangement for most atoms.
- Octet rule — the rule that an atom tends to attain 8 electrons in its valence shell, the configuration of the nearest inert gas.
- Octet theory of valency — the theory of Kossel and Lewis that atoms combine because their valence shells are incomplete, and combine so as to complete them.
- Orbit — an imaginary path around the nucleus along which an electron revolves; also called a shell.
- Oxidation — the loss of electrons by a substance.
- Photolysis — the breaking of a bond by high-energy radiation such as ultraviolet light.
- Polar covalent compound — a covalent compound in which the shared pair lies nearer one atom, so that fractional charges develop.
- Proton — a sub-atomic particle in the nucleus, carrying one unit of positive charge and having a mass of 1 amu.
- Radioactivity — the spontaneous decay of unstable radioisotopes, in which the nucleus breaks up and emits radiation.
- Radioisotope — an isotope whose nucleus is unstable and breaks up spontaneously, emitting radiation.
- Redox reaction — a reaction in which oxidation and reduction take place at the same time.
- Reduction — the gain of electrons by a substance.
- Shell — a fixed circular path around the nucleus in which electrons revolve; labelled K, L, M, N outwards.
- Single covalent bond — a covalent bond formed by the sharing of one pair of electrons, each atom contributing one electron.
- Triple bond — a covalent bond formed by the sharing of three pairs of electrons between two atoms.
- Valence electrons — the electrons present in the outermost shell of an atom.
- Valence shell — the outermost shell of an atom.
- Valency — the combining capacity of an element, equal to the number of electrons of its atom that take part in a chemical reaction.
4.20 Points To Ponder
Rutherford’s model was overturned within two years, and Dalton’s within a century, yet both are still taught. What does a scientific model have to do to be worth keeping, if being correct in every detail is not the test?
The atom is almost entirely empty space, and so is everything made of atoms. Why, then, can you not push your hand through a table? What is actually stopping it, if there is so little matter in the way?
Isotopes of an element are chemically identical but physically different. If you were handed a mixture of two isotopes and told to separate them, which kind of property would you have to work with, and why would ordinary chemical methods be useless?
An atom’s chemistry is decided entirely by the electrons in its outermost shell — a handful of particles out of the whole atom. Is it fair to say that the nucleus, which holds almost all the mass, plays no part in chemistry at all? Consider what the nucleus does determine.
The octet rule explains why sodium and chlorine react. But an atom does not know that argon exists, and does not aim at anything. What is really happening when we say an atom ‘tends to attain’ a stable configuration?
Sodium chloride conducts electricity when molten but not when solid, while sugar conducts in neither state. What could you conclude about an unknown white solid that dissolved in water and made the solution conduct — and what would you still not know about it?
4.21 Objective Questions
Choose the correct option: [1 each]
- The particle that is not present in the nucleus of an atom is the
- proton
- neutron
- electron
- nucleon
- The maximum number of electrons that the L shell can hold is
- 2
- 8
- 18
- 32
- An atom with the electronic configuration 2, 8, 7 will form
- a cation carrying a charge of \(+1\)
- an anion carrying a charge of \(-1\)
- a cation carrying a charge of \(+7\)
- no ion at all
- The charge-to-mass ratio of anode rays is greatest when the gas in the discharge tube is
- oxygen
- nitrogen
- hydrogen
- chlorine
- The number of shared pairs of electrons in a molecule of nitrogen is
- 1
- 2
- 3
- 4
- Which of the following pairs are isobars?
- \(^{35}_{17}\mathrm{Cl}\) and \(^{37}_{17}\mathrm{Cl}\)
- \(^{40}_{18}\mathrm{Ar}\) and \(^{40}_{20}\mathrm{Ca}\)
- \(^{12}_{6}\mathrm{C}\) and \(^{13}_{6}\mathrm{C}\)
- protium and deuterium
- An element X has the configuration 2, 8, 8, 2 and an element Y has the configuration 2, 6. The formula of the compound formed between them is
- \(\mathrm{XY}\)
- \(\mathrm{XY_2}\)
- \(\mathrm{X_2Y}\)
- \(\mathrm{X_2Y_3}\)
- Which of the following conducts electricity in the molten state?
- Methane
- Carbon tetrachloride
- Magnesium chloride
- Hydrogen
- A free radical is
- positively charged
- negatively charged
- neutral, with an unpaired electron
- neutral, with all its electrons paired
- The scientist who showed that it is the number of protons, and not the atomic mass, that identifies an element was
- Niels Bohr
- Henry Moseley
- James Chadwick
- E. Goldstein
- The particle that is not present in the nucleus of an atom is the
Each of the following consists of an Assertion (A) and a Reason (R). Choose from: (a) both A and R are true and R is the correct explanation of A; (b) both A and R are true but R is not the correct explanation of A; (c) A is true but R is false; (d) A is false but R is true. [1 each]
Assertion (A): An atom as a whole carries no charge. Reason (R): The number of protons in an atom is equal to the number of electrons.
Assertion (A): The M shell of a potassium atom holds only 8 electrons, although its capacity is 18. Reason (R): Whichever shell is outermost at the time cannot hold more than 8 electrons.
Assertion (A): Covalent compounds have low boiling points. Reason (R): The covalent bond inside a molecule is a weak bond.
Assertion (A): The two isotopes of chlorine undergo exactly the same chemical reactions. Reason (R): The two isotopes have the same number of neutrons.
Assertion (A): An anion is larger than the atom from which it is formed. Reason (R): The number of protons in the nucleus falls when an anion is formed.
Match each atomic number in the first list with the correct description in the second list. [4]
| Atomic number | Description |
|---|---|
| a. 6 | i. Forms a cation carrying a charge of \(+2\) |
| b. 10 | ii. Forms four covalent bonds |
| c. 12 | iii. Has a valency of zero |
| d. 17 | iv. Forms an anion carrying a charge of \(-1\) |
Fill in the blanks: [1 each]
- The rays that travel from the anode side towards the cathode in a discharge tube are called ………… rays.
- The number of neutrons in an atom is obtained by subtracting the ………… from the ………… .
- The bond between the two atoms in a molecule of oxygen is a ………… covalent bond.
- Atoms of different elements having the same mass number are called ………… .
- The symmetrical breaking of a covalent bond is called ………… bond fission.
State whether each of the following is true or false: [1 each]
- An electron and a proton carry charges of equal magnitude.
- The mass number of an atom is always a whole number.
- Cathode rays consist of positively charged particles.
- Covalent compounds conduct electricity in the molten state.
- Isotopes of an element have the same electronic configuration.
Name the following: [1 each]
- The isotope of hydrogen whose nucleus contains two neutrons.
- The sub-atomic particle discovered by James Chadwick.
- The shell of an atom that has the least energy.
- The type of bond formed between the two atoms in a chlorine molecule.
- The neutral species produced when a covalent bond breaks symmetrically.
Give one word or phrase for each of the following: [1 each]
- The distribution of the electrons of an atom among its shells.
- The combining capacity of an element.
- The spontaneous decay of an unstable nucleus with the emission of radiation.
- The outermost shell of an atom.
- The protons and neutrons of an atom taken together.
Pick the odd one out and give a reason for your choice: [2 each]
- Helium, neon, argon, chlorine.
- Sodium chloride, magnesium chloride, calcium oxide, methane.
- Sodium ion, chloride ion, chlorine radical, oxide ion.
4.22 Check Your Concepts
Define an isotope, and give one example other than that of hydrogen. [2]
State any two properties of cathode rays. [2]
Name the three sub-atomic particles of an atom and state the charge carried by each. [2]
Give a reason why an anion is larger than the atom from which it is formed. [2]
State what is meant by the valence shell of an atom, and say why it decides the chemistry of the element. [2]
Explain why the atomic mass of chlorine is 35.5, although no chlorine atom has that mass. [3]
Differentiate between an electrovalent bond and a covalent bond, stating clearly the basis of the difference. [3]
Why does argon form no compounds, while chlorine forms a great many? [3]
What would happen to an atom if a revolving electron did radiate energy, as the classical laws predict? State how Bohr’s model avoids this. [3]
Give reasons for the following: [4]
- Sodium chloride does not conduct electricity in the solid state but does so when molten.
- Covalent compounds are usually volatile.
Differentiate between isotopes and isobars, giving four points of difference. [4]
Explain, using the electronic configurations of the atoms concerned, how magnesium chloride is formed, and why its formula is \(\mathrm{MgCl_2}\) and not \(\mathrm{MgCl}\). [4]
Describe Rutherford’s gold foil experiment. State the three observations he made, and the conclusion he drew from each. [5]
State the rules of the Bohr–Bury scheme for filling the shells of an atom, and use them to work out the electronic configuration of an element of atomic number 20, explaining each step. [5]
4.23 Equations, Conversions & Identification
- Write equations, showing the electrons, for the changes in which: [2]
- a sodium atom becomes a sodium ion;
- an oxygen atom becomes an oxide ion.
- Complete and balance the following, and name the products: [2]
\[\mathrm{Cl_2} \xrightarrow{\ \mathrm{UV}\ } \ \ldots\ldots + \ldots\ldots\]
- Complete the following nuclear change and name the particle represented by the missing term: [2]
\[^{9}_{4}\mathrm{Be} + ^{4}_{2}\mathrm{He} \rightarrow ^{12}_{6}\mathrm{C} + \ldots\ldots\]
Write the balanced equation for the reaction between sodium and chlorine, and write the two half reactions, stating which is oxidation and which is reduction. [3]
A white crystalline solid melts only at a very high temperature. It does not conduct electricity in the solid state, but its molten form conducts well. [3]
- Identify the type of compound.
- Name the particles present in it.
- Explain why conduction begins only on melting.
An element A has the electronic configuration 2, 8, 1 and an element B has 2, 8, 6. [3]
- Identify which is the metal and which the non-metal, giving a reason.
- State the type of bond formed between them.
- Write the formula of the compound formed.
Look at Fig. 4.3-1. [3]
- Name the electrode from which the rays start, and state the sign of the terminal it is joined to.
- State the purpose of the side tube.
- State the two conditions that must be met before the gas will conduct.
Look at Fig. 4.15-2. [3]
- State what the curved arrows in the diagram represent.
- Explain why two chlorine atoms are shown for every magnesium atom.
- Write the configurations of the two ions that are produced.
How would you obtain the following? State in each case what has to be supplied or removed. [3]
- Chlorine radicals from a chlorine molecule.
- A calcium ion from a calcium atom.
- A proton from a hydrogen atom.
4.24 Practice With Numericals
An atom of an element contains 17 protons and 20 neutrons. Calculate its atomic number and its mass number, and write its electronic configuration. [3]
Boron occurs naturally as two isotopes, of mass numbers 10 and 11, present in the ratio 1 : 4. Calculate the atomic mass of boron. [3]
A sample of an element contains isotopes of mass numbers 20 and 22, in the proportions 90 per cent and 10 per cent respectively. Calculate the average atomic mass of the element. [3]
The unipositive ion of an element contains 10 electrons and 12 neutrons. Calculate the atomic number and the mass number of the element. [3]
An atom of an element contains 12 protons and 12 neutrons. [4]
- State its mass number.
- Write its electronic configuration.
- State its valency.
- Write the symbol of the ion it forms, with the charge.
An element X occurs as two isotopes of mass numbers 69 and 71, and its average atomic mass is 69.5. Calculate the percentage of each isotope present in the sample. [4]