What is an atom?

The word 'atom' comes from the Greek 'atomos' meaning 'indivisible' and an atom is the smallest unit (particle) still characterising a
chemical element. A chemical element can therefore be defined as a type of atom. Chemical substances include all the materials
we are familiar with in everyday experience as 'matter'. For example, steel, concrete, water and air are all chemicals or mixtures of
chemicals. However, most chemicals are
compounds, meaning they are made up of more than one element. Water for example is
given the chemical formula H20, meaning that water is made up of particles or molecules in which each molecule is made of two
hydrogen atoms and one oxygen atom joined or
bonded together to form the compound H2O or water.

Let's put this into perspective. A teaspoonful (5 ml) of water contains some 1.67 x 10^20 (167 000 000 000 000 000 000 or 167
million million million) molecules of water! Thus a water molecule is very tiny indeed! In fact one molecule of water weighs about 3 x
10^-11 nanograms (0.000 000 000 000 000 000 03 grams)! This is the smallest amount of water that one can have and still call it
water. Go any smaller and you have to break the single water molecule up into its component atoms of hydrogen and oxygen.
Break up the atoms of hydrogen and oxygen and you end up with sub-atomic particles and no matter remains with which we are
familiar with in everyday experience. Thus, one atom of oxygen is the smallest unit of material that we can have and still call it
oxygen. This is the difference between water and oxygen - the smallest unit of water, the H2O molecule, is still divisible into smaller
units that are made up of the smallest divisible units, or atoms, of other substances (oxygen and hydrogen) and these atoms
cannot be divided up further into any other familiar substance. In reality things are not quite so simple, since hydrogen and oxygen
gases at room temperature consist of molecules of two like atoms: H2 and O2 or H-H and O=O which shows that two bonds (a
double bond) actually hold the two oxygen atoms together. Water can similarly be written as: H-O-H, showing that the two
hydrogens are bonded to the oxygen but not directly to one another. However, some gases such as argon are non-molecular and
composed of single argon atoms: Ar. The diagram below shows a model of a water molecule.
Above left: a space-filled model of a water molecule, showing the two smaller hydrogen
atoms (in grey) bonded to the central oxygen atom (in red). Above right: a
ball-and-stick model of a water molecule. The sticks indicate the relative bond length.
Above: a third representation of the water
molecule, illustrating that the H-O-H bond
angle is 104.5 degrees.
Compound or element?

We can tell whether a substance is a compound, such as water, or an element such as argon, by its chemical formula. Each atom is
given a symbol and there are 118 known chemical types of atom (ignoring isotopes which are atoms of the same element that have
different masses). These 118 elements are given both a number and a name 9and a symbol that abbreviates the name). Elements 112
to 118 have not been named yet and so have only numbers. Only the elements numbering 94 (plutonium) and below are found in
nature on Earth, the rest (95-118) have been synthesised in laboratories or in nuclear reactors. The higher the number given to an
element the more complex is its atomic structure, since higher number elements are composed of more subatomic particles.

The Periodic Table

The diagram below is the periodic table of the elements. It shows all 118 elements, giving their atomic number and their symbol (and
usually also their mass, but this has been omitted for clarity here). For example, hydrogen (H) is element number 1 (the atomic number
of hydrogen is 1) whilst sodium (Na) is element number 11 and polonium (Po) is element number 84. The elements are grouped into
rows called periods, numbered (in red) from 1 to 7. The two additional rows labelled 'Lanthanides' and 'Actinides' come between
elements 57 (Lanthanum, La) and 72 (Hafnium, Hf) for the lanthanides, and between 89 (Actinium, Ac) and 104 (Rutherfordium, Rf) for
the actinides (as indicated by the double vertical lines) but are grouped separately because of their structural and chemical properties.

With the completion of each period, the chemical properties of the elements repeat and so the elements may be grouped into vertical
columns, with the members of each column having similar (or at least related) chemical properties. For example, group 1 consists of
hydrogen and the alkali metals lithium (Li), sodium (Na), potassium (K), rubidium (Rb), Caesium (Cs) and frankium (Fr). Hydrogen is
different since it is a gas whilst the others members of group 1 are metals, and so hydrogen is sometimes put in a block all by itself.
Putting hydrogen aside, lithium to frankium are elements with similar chemical and physical properties. These metals are soft (lithium
can be cut with difficulty a knife, whilst sodium and potassium are easy to cut with a knife and rubidium and caesium are rather like
metallic grey putty). All these metals react with water, fizzing and dissolving as hydrogen gas (from the water) is produced. However,
there are differences as one moves down a group. Lithium reacts slowly with water, but reactivity of the group 1 alkali metals increases
down the group. Sodium reacts rapidly with water, potassium reacts violently and catches fire (as the hydrogen produced burns in air to
form water again!) and frankium reacts explosively with water.  This is why these metals must be stored under oil in the laboratory.
In the simplest model, we can think of neutrons, electrons and protons as minute balls, with the protons being electrically positively charged
and the electrons negatively charged. Electricity in a wire is actually the flow of electrons carrying electric charge. Note that the amount of
electric charge carried by an electron is equal to that carried by a proton, but opposite in sign. A neutron carries no net electric charge.

The nucleus

The nucleus is at the centre of the atom and contains the protons and neutrons. The nucleus contains nearly all the mass of the atom
(because electrons have so little mass). Protons and neutrons are known collectively as nucleons. Despite containing almost all of the
atom's mass, the nucleus is tiny compared to the atom.
Above: the Periodic Table of the elements. The elements are arranged into horizontal rows
(periods) and vertical columns (groups) according to their chemical properties. The elements
are also divided into blocks (colour-coded) with the s-block in dark grey, the d-block in blue,
the p-block in green and the f-block (lanthanides and actinides) in light grey.
This atom has 9 protons and (N = A - Z = 19-9) 10 neutrons. The number of protons is equal to the atomic
number of the element (9 for fluorine). Note that A is the mass number and not the atomic number which is
Z! The atomic number tells us that this atom has 9 protons. The mass number tells us that this fluorine
isotope has 19 nucleons (protons + neutrons). Therefore, fluorine has 19 - 9 = 10 neutrons. (A = Z + N).
Atomic number and mass number must also balance in the examples we deal with – they are conserved. In the above equations, the
electric charge has been omitted, because these are
nuclear equations and we are not concerned with any electrons that may
orbit these nuclei.

Electron Shells

The electrons surround in the nucleus as electron shells. These shells are much larger (though much less massive) than the
nucleus. The electrons carry a negative electric charge and are held in place due to their attraction to the positively charged protons
in the nucleus (
opposite electric charges attract, like electric charges repel). Hydrogen is the simplest atom, consisting of a
single proton in its nucleus and one electron occupying a single shell.

If we give energy to this electron, it 'expands' and moves further from the nucleus (on average) as it has more energy to 'resist' the
electric attraction between it and the proton. If we give it enough energy then the electron may detach from the nucleus altogether
and fly off as a free electron, leaving the hydrogen nucleus exposed as a positively charged hydrogen ion H+.
An ion is an atom
that has lost or gained electrons, leaving it with a net positive or negative electric charge respectively
. An atom
ordinarily has a number of electrons equal to its number of protons (and hence to its atomic number) and so has no net negative
charge. Atoms prefer to be in this state since then they are less energetic (atoms are 'lazy'!) - to ionise an atom (remove an electron
to form a positively charged ion) we must give it energy. We can do this either by heating a collection of atoms (causing the atoms to
collide more violently, transferring energy to one another) or by bombarding the atoms with light (of the right wavelength or colour)
or by bombarding them with energetic particles, such as electrons. These electrons collide with the electrons in the atom and
energise them, possibly even ionising the atom.

Electrons are strange things and when attached to atoms they cannot have any arbitrary amount of energy, rather they can only
receive or emit energy in packets of a certain size, called quanta (singular quantum). Since the energy an electron has is
proportional to its distance from the nucleus, giving an electron one quantum of energy moves it certain distance further away from
the nucleus. Giving it more quanta moves it further away still. However, since half quanta do not exist, the electron can only be
positioned at one of a certain number of specific distances from the nucleus - its energy has to be one of a spectrum of discrete and
specific energies - its energy is quantised. This gives rise to the notion of energy levels. The more energy an electron has, the
further it is from the nucleus and the higher is its energy level. Eventually, if given enough energy, the electron is so far from the
nucleus that it escapes the attractive pull of the nucleus altogether and detaches from the atom which becomes ionised as the
electron takes away its negative charge with it.

For many electron atoms it is more useful to talk about electron shells rather than energy levels. Electron shells are labelled by
giving each one a principal quantum number, n. For the first shell n = 1, for the second shell n = 2, etc. The higher the value of n,
the further the shell is from the nucleus and so the greater is its energy. Each shell can hold more than one electron, but there is a

              first shell (K shell)              (n=1)        2 electrons
              second shell (L shell)        (n=2)        8 electrons
              third shell (M shell)            (n=3)        18 electrons
              fourth shell (N shell)          (n=4)        32 electrons

A shell that contains its maximum number of electrons is a
filled shell.

We can reconstruct an atom by starting with a naked nucleus, having no electrons and then adding back the electrons one at a time.
Electrons are arranged so that the
lowest energy shells are filled first. Each additional electron enters the lowest energy shell
that has space for it.

E.g. sodium has the electron shell configuration: 2.8.1
The table below lists the electron shell arrangements for the first 20 elements.
The outer shell electrons determine most of the chemistry of the atom (since this is the shell that is most
exposed to other atoms and so can react with them most easily). The diagram below illustrates a basic model of
a sodium atom.
What about potassium?

Potassium (K): Z = 19 and the electron shell configuration, but why?
To understand this we need to know that the electron shells are split into sub-shells. There are s, p, d and f sub-shells. Each sub-shell
can hold more than one electron, but there is a limit:

                      Sub-shell        Maximum number of electrons
                            s                                2
                            p                                6
                            d                                10
                            f                                14

Electron shells and sub-shells

The n = 1 shell contains 2 electrons in its s sub-shell only.
The n = 2 shell contains 2 electrons in its s sub-shell, 6 electrons in its p sub-shell and thus has a total of 8 electrons.
The n = 3 shell contains 2 electrons in its s sub-shell, 6 electrons in its p sub-shell and 10 electrons in its d sub-shell for a total of 18
The n = 4 shell  contains 2 electrons in its s sub-shell, 6 electrons in its p sub-shell, 10 electrons in its d sub-shell and 14 electrons in its
f sub-shell, for a total of 32 electrons.

Now this complicates the energy level diagram as the electrons within each subshell now have slightly different energies (though for a full
shell these energies average to give the shell energy designated by n). These energies are shown in the
energy-level diagram below:
Left: the energy levels corresponding to the first four
electron shells. The shells are designated by the
principle quantum number n. Thus, the n=1 electron
shell has less energy and is closer to the nucleus than
the n=2 electron shell. The number n must be a positive
whole number (1,2,3,4,5, ... ) as the energy levels are
quantised, meaning that an electron can only occupy
one of a specific spectrum of discrete energy levels or
electron shells. There are, however, an infinite number
of possible shells (though only the lowest will be filled as
the atom only has a finite number of electrons) and note
that the spacing between them gets less as the energy
increases, such that eventually a limit is reached for the
infinite shell. Imparting an electron with more energy
than this will expel it from the atom, ionising the atom,
and so is called the ionisation energy.
When we add electrons back to a naked fully ionised nucleus (with no electrons) then the electrons enter the lowest energy
sub-shell first. When this sub-shell is full, the next electrons enter the next lowest vacant sub-shell, and so on. This is the
aufbau principle. For elements up to and including nickel, Ni (Z = 28) the 4s sub-shell has lower energy than the 3d
sub-shell and so fills-up first!

Remember that an s sub-shell can only hold two electrons, a p sub-shell 6, a d sub-shell 10 and an f sub-shell 14 electrons.

This aufbau principle allows us to work out the
electron configuration of an atom, if we know how many electrons it
contains. For example, for potassium, K (element 19) with 19 electrons and for calcium, Ca (element 20) with 20 electrons,
the electron configurations are:

The s, p, d and f sub-shells are further divided into atomic orbitals. Each orbital holds a maximum of two electrons. An electron
in a given orbital can be found in a particular region of space around the nucleus.

Sub-shells always contain the following numbers of atomic orbitals:
An s sub-shell contains one s atomic orbital
A p sub-shell contains three p atomic orbitals
A d sub-shell contains five d atomic orbitals
An f sub-shell contains seven f atomic orbitals

Of course, not all orbitals are always full!! (Since atoms contain a finite number of electrons and electrons may jump to higher
energy levels if they gain sufficient energy. An electron that is so
excited to a higher energy level will eventually come back
down by losing its excess energy as a particle of light called a photon).

Electron Spin

Each atomic orbital can contain a maximum of two electrons (but these two electrons must have opposite or paired spins). You
can think of an electron spinning either in a clockwise direction (spin-up, or spin = +1/2) or anticlockwise (spin-down or spin = -1
/2) - though in reality it is not that simple!!! Electrons in the same orbital must have paired (opposite) spins:

The box represents the atomic orbital and the arrows the two spinning electrons. The up arrow is the electron with spin-up
(+1/2) and the down arrow the electron with spin-down (-1/2).

Thus, each electron has an ‘address’ - 4 bits of information that tells you where it is:

1.  The electron shell it is in (n)
2.  Its sub-shell (s, p, d or f)
3.  Its orbital within the sub-shell (e.g. px, py, pz)
4.  Its spin (up or down).

The full electron configuration of sodium is shown below:
We can continue in this way building up the periodic table by adding an electron at a time to the next available energy
level, this is the Aufbau principle. If we do this, we notice an interesting pattern! All the s-block elements (groups 1 and
2) are those in which the outermost electrons occupy s sub-shells. Furthermore, group 1 (the alkali metals) all have
only one s electron in this outermost sub-shell, so hydrogen (H) is simply 1s, lithium (Li) ends in 2s and sodium (Na) in
3s, where the principle quantum number n (the 1 of 1s or the 3 of 3s, etc.) gives us the period number to which the
element belongs! Similarly group 2 elements (the alkaline earth metals) such as calcium (Ca), barium (Ba) and
strontium outermost electrons that are filling the d sub-shells, the p-block are filling their p sub-shells and the f-block
their f sub-shells. It should be stressed that the periodic table was largely constructed before anything was known
about electron configurations, based on the chemical properties of the elements. This illustrates the fact that it is the
outermost electrons that determine the chemistry of the elements.
Some or all of the outermost electrons participate in forming bonds to other atoms, these electrons are called the valence
. For example, in water, H2O, each of the two H atoms donates its only 1s electron to form a bond with the oxygen.
The oxygen contributes its two unpaired 2p electrons (see the electron configuration of oxygen above). Thus four valence
electrons (one from each hydrogen and two from the oxygen) participate in forming a chemical bond between the oxygen and
each hydrogen, so four electrons produce two chemical bonds. That is, two electrons shared by the two bonding atoms (O
and H) form a single chemical bond. A chemical bond of this type, which involves a pair of shared valence electrons is called a
covalent bond.

The diagrams below show the shapes of some of the electron orbitals. Notice that the three p orbitals differ only in the
direction in which they point and are designated px, py and pz orbitals, as they are at right-angles to one another (orthogonal)
like the axes of a graph. Remember that each orbital can hold up to two electrons with opposite spins.
s-orbital (spherical)
px-orbital (dumbbell-shape)
py-orbital (dumbbell-shape)
pz-orbital (dumbbell-shape)
p-subshell (px + py + pz)
The shapes of d, f and higher (g,h, ..., etc.) orbitals are progressively more complex with more lobes. A 2s
differs from a is in being larger, but is otherwise the same external shape. Likewise 3s is larger than 2s and
4p larger than 3p, etc.
The Structure of the Atom

Although an atom is the smallest unit of an element, atoms can be split into smaller pieces, but ultimately the smallest of these pieces
resemble one of several types of sub-atomic particle, no matter what atom they come from.

The atom is made up of the following principle sub-atomic particles: electrons, protons and neutrons:
Sometimes the mass number is given as a superscript and the atomic number as a subscript, both
proceeding the element symbol. How many protons and neutrons does this atom of fluorine have?
The atomic number determines what element the atom is. All atoms with 9 protons have an atomic number 9 and are all fluorine
atoms. However, the number of neutrons can vary slightly in different isotopes of the same element.

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.
A 2p hydrogen atomic orbital.
Download a pdf on a mathematical
summary of Schrodinger's
Equation and the hydrogen atom.

Download a pdf on
Noble Gases,
including models of the helium
Explore the hydrogen atomic
spectrum interactively!

Go to the
physics intro.
Comment on this article!