Atomic Structure and Bonding
Fundamental Concepts
Atoms are composed of electrons, protons, and neutrons. Electron and protons are negative and
positive charges of the same magnitude, 1.6 × 10-19 Coulombs.
The mass of the electron is negligible with respect to those of the proton and the neutron, which
form the nucleus of the atom. The unit of mass is an atomic mass unit (amu) = 1.66 × 10 -27 kg, and
equals 1/12 the mass of a carbon atom. The Carbon nucleus has Z=6, and A=6, where Z is the
number of protons, and A the number of neutrons. Neutrons and protons have very similar masses,
roughly equal to 1 amu. A neutral atom has the same number of electrons and protons, Z.
A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the
atoms. Thus, a mole of carbon has a mass of 12 grams. The number of atoms in a mole is called the
Avogadro number, Nav = 6.023 × 1023. Note that Nav = 1 gram/1 amu.
Calculating n, the number of atoms per cm 3 in a piece of material of density (g/cm3).
n = Nav × / M
where M is the atomic mass in amu (grams per mol). Thus, for graphite (carbon) with a density =
1.8 g/cm3, M =12, we get 6 × 1023 atoms/mol × 1.8 g/cm3 / 12 g/mol) = 9 × 1022 C/cm3.
For a molecular solid like ice, one uses the molecular mass, M(H2O) = 18. With a density of 1
g/cm3, one obtains n = 3.3 × 1022 H2O/cm3. Note that since the water molecule contains 3 atoms,
this is equivalent to 9.9 × 1022 atoms/cm3.
Most solids have atomic densities around 6 × 1022 atoms/cm3. The cube root of that number gives
the number of atoms per centimeter, about 39 million. The mean distance between atoms is the
inverse of that, or 0.25 nm. This is an important number that gives the scale of atomic structures in
solids.
Electrons in Atoms
The forces in the atom are repulsions between electrons and attraction between electrons and
protons. The neutrons play no significant role. Thus, Z is what characterizes the atom.
The electrons form a cloud around the neutron, of radius of 0.05 – 2 nanometers. Electrons do not
move in circular orbits, as in popular drawings, but in 'fuzzy' orbits. We cannot tell how it moves,
but only say what is the probability of finding it at some distance from the nucleus. According to
quantum mechanics, only certain orbits are allowed (thus, the idea of a mini planetary system is not
correct). The orbits are identified by a principal quantum number n, which can be related to the size,
n = 0 is the smallest; n = 1, 2 .. are larger. (They are "quantized" or discrete, being specified by
integers). The angular momentum l is quantized, and so is the projection in a specific direction m.
The structure of the atom is determined by the Pauli exclusion principle, only two electrons can be
placed in an orbit with a given n, l, m – one for each spin. Table 2.1 in the textbook gives the
number of electrons in each shell (given by n) and subshells (given by l).
The Periodic Table
Elements are categorized by placing them in the periodic table. Elements in a column share similar
properties. The noble gases have closed shells, and so they do not gain or lose electrons near another
atom. Alkalis can easily lose an electron and become a closed shell; halogens can easily gain one to
form a negative ion, again with a closed shell. The propensity to form closed shells occurs in
molecules, when they share electrons to close a molecular shell. Examples are H2, N2, and NaCl.
The ability to gain or lose electrons is termed electronegativity or electropositivity, an important
factor in ionic bonds.
Bonding Forces and Energies
The Coulomb forces are simple: attractive between electrons and nuclei, repulsive between electrons
and between nuclei. The force between atoms is given by a sum of all the individual forces, and the
fact that the electrons are located outside the atom and the nucleus in the center.
When two atoms come very close, the force between them is always repulsive, because the electrons
stay outside and the nuclei repel each other. Unless both atoms are ions of the same charge (e.g.,
both negative) the forces between atoms is always attractive at large internuclear distances r. Since
the force is repulsive at small r, and attractive at small r, there is a distance at which the force is zero.
This is the equilibrium distance at which the atoms prefer to stay.
The interaction energy is the potential energy between the atoms. It is negative if the atoms are bound
and positive if they can move away from each other. The interaction energy is the integral of the
force over the separation distance, so these two quantities are directly related. The interaction energy
is a minimum at the equilibrium position. This value of the energy is called the bond energy, and is the
energy needed to separate completely to infinity (the work that needs to be done to overcome the
attractive force.) The strongest the bond energy, the hardest is to move the atoms, for instance the
hardest it is to melt the solid, or to evaporate its atoms.
Primary Interatomic Bonds
Ionic Bonding
This is the bond when one of the atoms is negative (has an extra electron) and another is positive
(has lost an electron). Then there is a strong, direct Coulomb attraction. An example is NaCl. In the
molecule, there are more electrons around Cl, forming Cl- and less around Na, forming Na+. Ionic
bonds are the strongest bonds. In real solids, ionic bonding is usually combined with covalent
bonding. In this case, the fractional ionic bonding is defined as %ionic = 100 × [1 – exp(-0.25 (XA –
XB)2], where XA and XB are the electronegativities of the two atoms, A and B, forming the molecule.
Covalent Bonding
In covalent bonding, electrons are shared between the molecules, to saturate the valency. The
simplest example is the H2 molecule, where the electrons spend more time in between the nuclei
than outside, thus producing bonding.
Metallic Bonding
In metals, the atoms are ionized, loosing some electrons from the valence band. Those electrons
form a electron sea, which binds the charged nuclei in place, in a similar way that the electrons in
between the H atoms in the H2 molecule bind the protons.
Secondary Bonding (Van der Waals)
Fluctuating Induced Dipole Bonds
Since the electrons may be on one side of the atom or the other, a dipole is formed: the + nucleus at
the center, and the electron outside. Since the electron moves, the dipole fluctuates. This fluctuation
in atom A produces a fluctuating electric field that is felt by the electrons of an adjacent atom, B.
Atom B then polarizes so that its outer electrons are on the side of the atom closest to the + side (or
opposite to the – side) of the dipole in A. This bond is called van der Waals bonding.
Polar Molecule-Induced Dipole Bonds
A polar molecule like H2O (Hs are partially +, O is partially – ), will induce a dipole in a nearby
atom, leading to bonding.
Permanent Dipole Bonds
This is the case of the hydrogen bond in ice. The H end of the molecule is positively charged and
can bond to the negative side of another dipolar molecule, like the O side of the H 2O dipole.
Molecules
If molecules formed a closed shell due to covalent bonding (like H 2, N2) then the interaction
between molecules is weak, of the van der Waals type. Thus, molecular solids usually have very low
melting points.
ATOMIC PROPERTIES MENU
The sub-atomic particles
Protons, neutrons and electrons.
proton
neutron
electron
relative mass
1
1
1/1836
relative charge
+1
0
-1
Beyond A'level: Protons and neutrons don't in fact have exactly the same mass - neither of them has a mass of exactly 1
on the carbon-12 scale (the scale on which the relative masses of atoms are measured). On the carbon-12 scale, a proton
has a mass of 1.0073, and a neutron a mass of 1.0087.
The behaviour of protons, neutrons and electrons in electric fields
What happens if a beam of each of these particles is passed between two electrically charged plates one positive and one negative? Opposites will attract.
Protons are positively charged and so would be deflected on a curving path towards the negative
plate.
Electrons are negatively charged and so would be deflected on a curving path towards the positive
plate.
Neutrons don't have a charge, and so would continue on in a straight line.
Exactly what happens depends on whether the beams of particles enter the electric field with the
various particles having the same speeds or the same energies
If the particles have the same energy
If beams of the three sorts of particles, all with the same energy, are passed between two electrically
charged plates:


Protons are deflected on a curved path towards the negative plate.
Electrons are deflected on a curved path towards the positive plate.
The amount of deflection is exactly the same in the electron beam as the proton beam if the
energies are the same - but, of course, it is in the opposite direction.

Neutrons continue in a straight line.
If the electric field was strong enough, then the electron and proton beams might curve enough to
hit their respective plates.
If the particles have the same speeds
If beams of the three sorts of particles, all with the same speed, are passed between two electrically
charged plates:


Protons are deflected on a curved path towards the negative plate.
Electrons are deflected on a curved path towards the positive plate.
If the electrons and protons are travelling with the same speed, then the lighter electrons are
deflected far more strongly than the heavier protons.

Neutrons continue in a straight line.
Note: This is potentially very confusing! Most chemistry sources that talk about this give either one or the other of
these two diagrams without any comment at all - they don't specifically say that they are using constant energy or
constant speed beams. But it matters!
If this is on your syllabus, it is important that you should know which version your examiners are going to expect, and
they probably won't tell you in the syllabus. You should look in detail at past questions, mark schemes and examiner's
reports which you can get from your examiners if you are doing a UK-based syllabus. Information about how to do this
is on the syllabuses page.
If in doubt, I suggest you use the second (constant speed) version. This actually produces more useful information about
both masses and charges than the constant energy version.
The nucleus
The nucleus is at the centre of the atom and contains the protons and neutrons. Protons and
neutrons are collectively known as nucleons.
Virtually all the mass of the atom is concentrated in the nucleus, because the electrons weigh so
little.
Working out the numbers of protons and neutrons
No of protons = ATOMIC NUMBER of the atom
The atomic number is also given the more descriptive name of proton number.
No of protons + no of neutrons = MASS NUMBER of the atom
The mass number is also called the nucleon number.
This information can be given simply in the form:
How many protons and neutrons has this atom got?
The atomic number counts the number of protons (9); the mass number counts protons + neutrons
(19). If there are 9 protons, there must be 10 neutrons for the total to add up to 19.
The atomic number is tied to the position of the element in the Periodic Table and therefore the
number of protons defines what sort of element you are talking about. So if an atom has 8 protons
(atomic number = 8), it must be oxygen. If an atom has 12 protons (atomic number = 12), it must
be magnesium.
Similarly, every chlorine atom (atomic number = 17) has 17 protons; every uranium atom (atomic
number = 92) has 92 protons.
Isotopes
The number of neutrons in an atom can vary within small limits. For example, there are three kinds
of carbon atom 12C, 13C and 14C. They all have the same number of protons, but the number of
neutrons varies.
carbon-12
carbon-13
carbon-14
protons
6
6
6
neutrons
6
7
8
mass number
12
13
14
These different atoms of carbon are called isotopes. The fact that they have varying numbers of
neutrons makes no difference whatsoever to the chemical reactions of the carbon.
Isotopes are atoms which have the same atomic number but different mass numbers. They have the
same number of protons but different numbers of neutrons.
The electrons
Working out the number of electrons
Atoms are electrically neutral, and the positiveness of the protons is balanced by the negativeness of
the electrons. It follows that in a neutral atom:
no of electrons = no of protons
So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; if a chlorine
atom (atomic number = 17) has 17 protons, it must also have 17 electrons.
The arrangement of the electrons
The electrons are found at considerable distances from the nucleus in a series of levels called energy
levels. Each energy level can only hold a certain number of electrons. The first level (nearest the
nucleus) will only hold 2 electrons, the second holds 8, and the third also seems to be full when it
has 8 electrons. At GCSE you stop there because the pattern gets more complicated after that.
These levels can be thought of as getting progressively further from the nucleus. Electrons will
always go into the lowest possible energy level (nearest the nucleus) - provided there is space.
To work out the electronic arrangement of an atom



Look up the atomic number in the Periodic Table - making sure that you choose the right
number if two numbers are given. The atomic number will always be the smaller one.
This tells you the number of protons, and hence the number of electrons.
Arrange the electrons in levels, always filling up an inner level before you go to an outer one.
e.g. to find the electronic arrangement in chlorine



The Periodic Table gives you the atomic number of 17.
Therefore there are 17 protons and 17 electrons.
The arrangement of the electrons will be 2, 8, 7 (i.e. 2 in the first level, 8 in the second, and 7
in the third).
The electronic arrangements of the first 20 elements
After this the pattern alters as you enter the transition series in the Periodic Table.
Two important generalisations
If you look at the patterns in this table:

The number of electrons in the outer level is the same as the group number. (Except with
helium which has only 2 electrons. The noble gases are also usually called group 0 - not
group 8.) This pattern extends throughout the Periodic Table for the main groups (i.e. not
including the transition elements).
So if you know that barium is in group 2, it has 2 electrons in its outer level; iodine (group 7)
has 7 electrons in its outer level; lead (group 4) has 4 electrons in its outer level.

Noble gases have full outer levels. This generalisation will need modifying for A'level
purposes.
Dots-and-crosses diagrams
In any introductory chemistry course you will have come across the electronic structures of
hydrogen and carbon, for example, drawn as:
Note: There are many places where you could still make use of this model of the atom at A'level. It is, however, a
simplification and can be misleading. It gives the impression that the electrons are circling the nucleus in orbits like
planets around the sun. As you will find when you look at the A'level view of the atom, it is impossible to know exactly
how they are actually moving.
The circles show energy levels - representing increasing distances from the nucleus. You could
straighten the circles out and draw the electronic structure as a simple energy diagram.
Carbon, for example, would look like this:
Thinking of the arrangement of the electrons in this way makes a useful bridge to the A'level view.
IONIC (ELECTROVALENT) BONDING
This page explains what ionic (electrovalent) bonding is. It starts with a simple picture of the
formation of ions, and then modifies it slightly for A'level purposes.
Ionic bonding in sodium chloride
Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that
electron it would become more stable.
Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain an
electron from somewhere it too would become more stable.
The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both become more
stable.
The sodium has lost an electron, so it no longer has equal numbers of electrons and protons.
Because it has one more proton than electron, it has a charge of 1+. If electrons are lost from an
atom, positive ions are formed.
Positive ions are sometimes called cations.
The chlorine has gained an electron, so it now has one more electron than proton. It therefore has a
charge of 1-. If electrons are gained by an atom, negative ions are formed.
A negative ion is sometimes called an anion.
The nature of the bond
The sodium ions and chloride ions are held together by the strong electrostatic attractions between
the positive and negative charges.
The formula of sodium chloride
You need one sodium atom to provide the extra electron for one chlorine atom, so they combine
together 1:1. The formula is therefore NaCl.
Some other examples of ionic bonding
Magnesium oxide
Again, noble gas structures are formed, and the magnesium oxide is held together by very strong
attractions between the ions. The ionic bonding is stronger than in sodium chloride because this
time you have 2+ ions attracting 2- ions. The greater the charge, the greater the attraction.
The formula of magnesium oxide is MgO.
calcium chloride
This time you need two chlorines to use up the two outer electrons in the calcium. The formula of
calcium chloride is therefore CaCl2.
potassium oxide
Again, noble gas structures are formed. It takes two potassiums to supply the electrons the oxygen
needs. The formula of potassium oxide is K2O.
THE A'LEVEL VIEW OF IONIC BONDING


Electrons are transferred from one atom to another resulting in the formation of positive
and negative ions.
The electrostatic attractions between the positive and negative ions hold the compound
together.
So what's new? At heart - nothing. What needs modifying is the view that there is something magic
about noble gas structures. There are far more ions which don't have noble gas structures than there
are which do.
Some common ions which don't have noble gas structures
You may have come across some of the following ions in a basic course like GCSE. They are all
perfectly stable , but not one of them has a noble gas structure.
Fe3+
Cu2+
Zn2+
Ag+
Pb2+
[Ar]3d5
[Ar]3d9
[Ar]3d10
[Kr]4d10
[Xe]4f145d106s2
Noble gases (apart from helium) have an outer electronic structure ns2np6.
What determines what the charge is on an ion?
Elements combine to make the compound which is as stable as possible - the one in which the
greatest amount of energy is evolved in its making. The more charges a positive ion has, the greater
the attraction towards its accompanying negative ion. The greater the attraction, the more energy is
released when the ions come together.
That means that elements forming positive ions will tend to give away as many electrons as possible.
But there's a down-side to this.
Energy is needed to remove electrons from atoms. This is called ionisation energy. The more
electrons you remove, the greater the total ionisation energy becomes. Eventually the total ionisation
energy needed becomes so great that the energy released when the attractions are set up between
positive and negative ions isn't large enough to cover it.
The element forms the ion which makes the compound most stable - the one in which most energy
is released over-all.
For example, why is calcium chloride CaCl2 rather than CaCl or CaCl3?
If one mole of CaCl (containing Ca+ ions) is made from its elements, it is possible to estimate that
about 171 kJ of heat is evolved.
However, making CaCl2 (containing Ca2+ ions) releases more heat. You get 795 kJ. That extra
amount of heat evolved makes the compound more stable, which is why you get CaCl 2 rather than
CaCl.
What about CaCl3 (containing Ca3+ ions)? To make one mole of this, you can estimate that you
would have to put in 1341 kJ. This makes this compound completely non-viable. Why is so much
heat needed to make CaCl3? It is because the third ionisation energy (the energy needed to remove
the third electron) is extremely high (4940 kJ mol-1) because the electron is being removed from the
3-level rather than the 4-level. Because it is much closer to the nucleus than the first two electrons
removed, it is going to be held much more strongly.
Note: It would pay you to read about ionisation energies if you really want to understand this.
You could also go to a standard text book and investigate Born-Haber Cycles.
A similar sort of argument applies to the negative ion. For example, oxygen forms an O 2- ion rather
than an O- ion or an O3- ion, because compounds containing the O2- ion turn out to be the most
energetically stable.
A simple view of covalent bonding
The importance of noble gas structures
At a simple level (like GCSE) a lot of importance is attached to the electronic structures of noble
gases like neon or argon which have eight electrons in their outer energy levels (or two in the case of
helium). These noble gas structures are thought of as being in some way a "desirable" thing for an
atom to have.
You may well have been left with the strong impression that when other atoms react, they try to
achieve noble gas structures.
As well as achieving noble gas structures by transferring electrons from one atom to another as in
ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to
give covalent bonds.
Some very simple covalent molecules
Chlorine
For example, two chlorine atoms could both achieve stable structures by sharing their single
unpaired electron as in the diagram.
The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is
simply to show where all the electrons come from. In reality there is no difference between them.
The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine
atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine
atoms.
Hydrogen
Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of
helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is
attracted to both nuclei.
Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure.
Covalent bonding at A'level
Cases where there isn't any difference from the simple view
If you stick closely to modern A'level syllabuses, there is little need to move far from the simple
(GCSE) view. The only thing which must be changed is the over-reliance on the concept of noble
gas structures. Most of the simple molecules you draw do in fact have all their atoms with noble gas
structures.
For example:
Even with a more complicated molecule like PCl3, there's no problem. In this case, only the outer
electrons are shown for simplicity. Each atom in this structure has inner layers of electrons of 2,8.
Again, everything present has a noble gas structure.
Cases where the simple view throws up problems
Boron trifluoride, BF3
A boron atom only has 3 electrons in its outer level, and there is no possibility of it reaching a noble
gas structure by simple sharing of electrons. Is this a problem? No. The boron has formed the
maximum number of bonds that it can in the circumstances, and this is a perfectly valid structure.
Energy is released whenever a covalent bond is formed. Because energy is being lost from the
system, it becomes more stable after every covalent bond is made. It follows, therefore, that an atom
will tend to make as many covalent bonds as possible. In the case of boron in BF3, three bonds is
the maximum possible because boron only has 3 electrons to share.
Note: You might perhaps wonder why boron doesn't form ionic bonds with fluorine instead. Boron doesn't form ions
because the total energy needed to remove three electrons to form a B3+ ion is simply too great to be recoverable when
attractions are set up between the boron and fluoride ions.
Phosphorus(V) chloride, PCl5
In the case of phosphorus 5 covalent bonds are possible - as in PCl5.
Phosphorus forms two chlorides - PCl3 and PCl5. When phosphorus burns in chlorine both are
formed - the majority product depending on how much chlorine is available. We've already looked at
the structure of PCl3.
The diagram of PCl5 (like the previous diagram of PCl3) shows only the outer electrons.
Notice that the phosphorus now has 5 pairs of electrons in the outer level - certainly not a noble gas
structure. You would have been content to draw PCl3 at GCSE, but PCl5 would have looked very
worrying.
Why does phosphorus sometimes break away from a noble gas structure and form five bonds? In
order to answer that question, we need to explore territory beyond the limits of current A'level
syllabuses. Don't be put off by this! It isn't particularly difficult, and is extremely useful if you are
going to understand the bonding in some important organic compounds.