After the electron was discovered by J.J. Thomson in 1897, people realised that atoms were made up of even smaller particles than they had previously thought. However, the atomic nucleus had not been discovered yet and so the 'plum pudding model' was put forward in 1904. In this model, the atom is made up of negative electrons that float in a soup of positive charge, much like plums in a pudding or raisins in a fruit cake (Figure 2). In 1906, Thomson was awarded the Nobel Prize for his work in this field. However, even with the Plum Pudding Model, there was still no understanding of how these electrons in the atom were arranged.

Figure 2: A schematic diagram to show what the atom looks like according to the Plum Pudding model

The discovery of radiation was the next step along the path to building an accurate picture of atomic structure. In the early twentieth century, Marie Curie and her husband Pierre, discovered that some elements (the radioactive elements) emit particles, which are able to pass through matter in a similar way to X-rays (read more about this in Grade 11). It was Ernest Rutherford who, in 1911, used this discovery to revise the model of the atom.

Radioactive elements emit different types of particles. Some of these are positively charged alpha (αα) particles. Rutherford carried out a series of experiments where he bombarded sheets of gold foil with these particles, to try to get a better understanding of where the positive charge in the atom was. A simplified diagram of his experiment is shown in Figure 3.

Figure 3: Rutherford's gold foil experiment. Figure (a) shows the path of the αα particles after they hit the gold sheet. Figure (b) shows the arrangement of atoms in the gold sheets and the path of the αα particles in relation to this.

Rutherford set up his experiment so that a beam of alpha particles was directed at the gold sheets. Behind the gold sheets was a screen made of zinc sulphide. This screen allowed Rutherford to see where the alpha particles were landing. Rutherford knew that the electrons in the gold atoms would not really affect the path of the alpha particles, because the mass of an electron is so much smaller than that of a proton. He reasoned that the positively charged protons would be the ones to repel the positively charged alpha particles and alter their path.

What he discovered was that most of the alpha particles passed through the foil undisturbed and could be detected on the screen directly behind the foil (A). Some of the particles ended up being slightly deflected onto other parts of the screen (B). But what was even more interesting was that some of the particles were deflected straight back in the direction from where they had come (C)! These were the particles that had been repelled by the positive protons in the gold atoms. If the Plum Pudding model of the atom were true then Rutherford would have expected much more repulsion, since the positive charge according to that model is distributed throughout the atom. But this was not the case. The fact that most particles passed straight through suggested that the positive charge was concentrated in one part of the atom only.

Rutherford's work led to a change in ideas around the atom. His new model described the atom as a tiny, dense, positively charged core called a nucleus surrounded by lighter, negatively charged electrons. Another way of thinking about this model was that the atom was seen to be like a mini solar system where the electrons orbit the nucleus like planets orbiting around the sun. A simplified picture of this is shown in Figure 4. This model is sometimes known as the planetary model of the atom.

Figure 4: Rutherford's model of the atom

There were, however, some problems with this model: for example it could not explain the very interesting observation that atoms only emit light at certain wavelengths or frequencies. Niels Bohr solved this problem by proposing that the electrons could only orbit the nucleus in certain special orbits at different energy levels around the nucleus. The exact energies of the orbitals in each energy level depends on the type of atom. Helium for example, has different energy levels to Carbon. If an electron jumps down from a higher energy level to a lower energy level, then light is emitted from the atom. The energy of the light emitted is the same as the gap in the energy between the two energy levels. You can read more about this in "Energy quantisation and electron configuration". The distance between the nucleus and the electron in the lowest energy level of a hydrogen atom is known as the Bohr radius.

Although the most common model of the atom is the Bohr model, scientists have not stopped thinking about other ways to describe atoms. One of the most important contributions to atomic theory (the field of science that looks at atoms) was the development of quantum theory. Schrodinger, Heisenberg, Born and many others have had a role in developing quantum theory. The description of an atom by quantum theory is very complex and is only covered at university level.

Match the information in column A, with the key discoverer in column B.

It is possible to determine the mass of a single atom in kilograms. But to do this, you would need very modern mass spectrometers and the values you would get would be very clumsy and difficult to use. The mass of a carbon atom, for example, is about 1,99×10-26kg1,99×10-26kg, while the mass of an atom of hydrogen is about 1,67×10-27kg1,67×10-27kg. Looking at these very small numbers makes it difficult to compare how much bigger the mass of one atom is when compared to another.

To make the situation simpler, scientists use a different unit of mass when they are describing the mass of an atom. This unit is called the atomic mass unit (amu). We can abbreviate (shorten) this unit to just 'u'. Scientists use the carbon standard to determine amu. The carbon standard assigns carbon an atomic mass of 12 u. Using the carbon standard the mass of an atom of hydrogen will be 1 u. You can check this by dividing the mass of a carbon atom in kilograms (see above) by the mass of a hydrogen atom in kilograms (you will need to use a calculator for this!). If you do this calculation, you will see that the mass of a carbon atom is twelve times greater than the mass of a hydrogen atom. When we use atomic mass units instead of kilograms, it becomes easier to see this. Atomic mass units are therefore not giving us the actual mass of an atom, but rather its mass relative to the mass of one (carefully chosen) atom in the Periodic Table. Although carbon is the usual element to compare other elements to, oxygen and hydrogen have also been used. The important thing to remember here is that the atomic mass unit is relative to one (carefully chosen) element. The atomic masses of some elements are shown in the table below.

The actual value of 1 atomic mass unit is 1,67×10-24g1,67×10-24g or 1,67×10-27kg1,67×10-27kg. This is a very tiny mass!

Atomic radius also varies depending on the element. On average, the radius of an atom ranges from 32pm32pm (Helium) to 225pm225pm (Caesium). Using different units, 100pm=1Angstrom100pm=1Angstrom, and 1Angstrom=10-10m1Angstrom=10-10m. That is the same as saying that 1Angstrom=0,0000000010 m1Angstrom=0,0000000010m or that 100pm=0,0000000010 m100pm=0,0000000010m! In other words, the diameter of an atom ranges from 0,0000000010m0,0000000010m to 0,0000000067m0,0000000067m. This is very small indeed.

The atomic radii given above are for the whole atom (nucleus and electrons). The nucleus itself is even smaller than this by a factor of about 23 000 in uranium and 145 000 in hydrogen. If the nucleus were the size of a golf ball, then the nearest electrons would be about one kilometer away! This should give help you realise that the atom is mostly made up of empty space.

The electron is a very light particle. It has a mass of 9,11×10-31kg9,11×10-31kg. Scientists believe that the electron can be treated as a point particle or elementary particle meaning that it can't be broken down into anything smaller. The electron also carries one unit of negative electric charge which is the same as 1,6×10-19C1,6×10-19C (Coulombs).

The electrons determine the charge on an atom. If the number of electrons is the same as the number of protons then the atom will be neutral. If the number of electrons is greater than the number of protons then the atom will be negatively charged. If the number of electrons is less than the number of protons then the atom will be positively charged. Atoms that are not neutral are called ions. Ions will be covered in more detail in a later chapter. For now all you need to know is that for each electron you remove from an atom you loose -1-1 of charge and for each electron that you add to an atom you gain +1+1 of charge. For example, the charge on an atom of sodium after removing one electron is -1-1.

Unlike the electron, the nucleus can be broken up into smaller building blocks called protons and neutrons. Together, the protons and neutrons are called nucleons.

The chemical properties of an element depend on the number of protons and electrons inside the atom. So if a neutron or two is added or removed from the nucleus, then the chemical properties will not change. This means that such an atom would remain in the same place in the Periodic Table. For example, no matter how many neutrons we add or subtract from a nucleus with 6 protons, that element will always be called carbon and have the element symbol CC (see the Table of Elements). Atoms which have the same number of protons, but a different number of neutrons, are called isotopes.

Definition 4: Isotope

The isotope of a particular element is made up of atoms which have the same number of protons as the atoms in the original element, but a different number of neutrons.

The different isotopes of an element have the same atomic number ZZ but different mass numbers AA because they have a different number of neutrons NN. The chemical properties of the different isotopes of an element are the same, but they might vary in how stable their nucleus is. Note that we can also write elements as X - AX - A where the X is the element symbol and the A is the atomic mass of that element. For example, C-12C-12 has an atomic mass of 12 and Cl-35Cl-35 has an atomic mass of 35 u, while Cl-37Cl-37 has an atomic mass of 37 u.

The following worked examples will help you to understand the concept of an isotope better.

It is important to realise that the atomic mass of isotopes of the same element will be different because they have a different number of nucleons. Chlorine, for example, has two common isotopes which are chlorine-35 and chlorine-37. Chlorine-35 has an atomic mass of 35 u, while chlorine-37 has an atomic mass of 37 u. In the world around us, both of these isotopes occur naturally. It doesn't make sense to say that the element chlorine has an atomic mass of 35 u, or that it has an atomic mass of 37 u. Neither of these are absolutely true since the mass varies depending on the form in which the element occurs. We need to look at how much more common one is than the other in order to calculate the relative atomic mass for the element chlorine. This is the number that you find on the Periodic Table.

Definition 5: Relative atomic mass

Relative atomic mass is the average mass of one atom of all the naturally occurring isotopes of a particular chemical element, expressed in atomic mass units.

This simulation allows you to see how isotopes and relative atomic mass are inter related.

Figure 7

You will remember from our earlier discussions that an atom is made up of a central nucleus, which contains protons and neutrons and that this nucleus is surrounded by electrons. Although these electrons all have the same charge and the same mass, each electron in an atom has a different amount of energy. Electrons that have the lowest energy are found closest to the nucleus where the attractive force of the positively charged nucleus is the greatest. Those electrons that have higher energy, and which are able to overcome the attractive force of the nucleus, are found further away.

If the energy of an atom is increased (for example when a substance is heated), the energy of the electrons inside the atom can be increased (when an electron has a higher energy than normal it is said to be "excited"). For the excited electron to go back to its original energy (called the ground state), it needs to release energy. It releases energy by emitting light. If one heats up different elements, one will see that for each element, light is emitted only at certain frequencies (or wavelengths). Instead of a smooth continuum of frequencies, we see lines (called emission lines) at particular frequencies. These frequencies correspond to the energy of the emitted light. If electrons could be excited to any energy and lose any amount of energy, there would be a continuous spread of light frequencies emitted. However, the sharp lines we see mean that there are only certain particular energies that an electron can be excited to, or can lose, for each element.

You can think of this like going up a flight of steps: you can't lift your foot by any amount to go from the ground to the first step. If you lift your foot too low you'll bump into the step and be stuck on the ground level. You have to lift your foot just the right amount (the height of the step) to go to the next step, and so on. The same goes for electrons and the amount of energy they can have. This is called quantisation of energy because there are only certain quantities of energy that an electron can have in an atom. Like steps, we can think of these quantities as energy levels in the atom. The energy of the light released when an electron drops down from a higher energy level to a lower energy level is the same as the difference in energy between the two levels.

We will start with a very simple view of the arrangement or configuration of electrons around an atom. This view simply states that electrons are arranged in energy levels (or shells) around the nucleus of an atom. These energy levels are numbered 1, 2, 3, etc. Electrons that are in the first energy level (energy level 1) are closest to the nucleus and will have the lowest energy. Electrons further away from the nucleus will have a higher energy.

In the following examples, the energy levels are shown as concentric circles around the central nucleus. The important thing to know for these diagrams is that the first energy level can hold 2 electrons, the second energy level can hold 8 electrons and the third energy level can hold 8 electrons.

But the situation is slightly more complicated than this. Within each energy level, the electrons move in orbitals. An orbital defines the spaces or regions where electrons move.

Definition 6: Atomic orbital

An atomic orbital is the region in which an electron may be found around a single atom.

There are different orbital shapes, but we will be mainly dealing with only two. These are the 's' and 'p' orbitals (there are also 'd' and 'f' orbitals). The 's' orbitals are spherical and the 'p' orbitals are dumbbell shaped.

Figure 11: The shapes of orbitals. a) shows an 's' orbital, b) shows a single 'p' orbital and c) shows the three 'p' orbitals.

The first energy level contains only one 's' orbital, the second energy level contains one 's' orbital and three 'p' orbitals and the third energy level contains one 's' orbital and three 'p' orbitals (as well as 5 'd' orbitals). Within each energy level, the 's' orbital is at a lower energy than the 'p' orbitals. This arrangement is shown in Figure 12.

Figure 12: The positions of the first ten orbitals of an atom on an energy diagram. Note that each block is able to hold two electrons.

This diagram also helps us when we are working out the electron configuration of an element. The electron configuration of an element is the arrangement of the electrons in the shells and subshells. There are a few guidelines for working out the electron configuration. These are:

In the examples you will cover, you will mainly be filling the s and p subshells. Occasionally you may get an example that has the d subshell. The f subshell is more complex and is not covered at this level.

The way that electrons are arranged in an atom is called its electron configuration.

Definition 7: Electron configuration

Electron configuration is the arrangement of electrons in an atom, molecule or other physical structure.

An element's electron configuration can be represented using Aufbau diagrams or energy level diagrams. An Aufbau diagram uses arrows to represent electrons. You can use the following steps to help you to draw an Aufbau diagram:

Sometimes people refer to Hund's rule for electron configuration. This rule simply says that electrons would rather be in a subshell on it's own then share a subshell. This is why, when you are filling the subshells you put one electron in each subshell and only if there are extra electrons do you go back and fill the subshell, before moving onto the next energy level.

An Aufbau diagram for the element Lithium is shown in Figure 13.

Figure 13: The electron configuration of Lithium, shown on an Aufbau diagram

A special type of notation is used to show an atom's electron configuration. The notation describes the energy levels, orbitals and the number of electrons in each. For example, the electron configuration of lithium is 1s22s11s22s1. The number and letter describe the energy level and orbital and the number above the orbital shows how many electrons are in that orbital.

Aufbau diagrams for the elements fluorine and argon are shown in Figure 14 and Figure 15 respectively. Using standard notation, the electron configuration of fluorine is 1s22s22p51s22s22p5 and the electron configuration of argon is 1s22s22p61s22s22p6.

Figure 14: An Aufbau diagram showing the electron configuration of fluorine

Figure 15: An Aufbau diagram showing the electron configuration of argon

Electrons in the outermost energy level of an atom are called valence electrons. The electrons that are in the energy shells closer to the nucleus are called core electrons. Core electrons are all the electrons in an atom, excluding the valence electrons. An element that has its valence energy level full is more stable and less likely to react than other elements with a valence energy level that is not full.

Definition 8: Valence electrons

The electrons in the outer energy level of an atom

Definition 9: Core electrons

All the electrons in an atom, excluding the valence electrons

By this stage, you may well be wondering why it is important for you to understand how electrons are arranged around the nucleus of an atom. Remember that during chemical reactions, when atoms come into contact with one another, it is the electrons of these atoms that will interact first. More specifically, it is the valence electrons of the atoms that will determine how they react with one another.

To take this a step further, an atom is at its most stable (and therefore unreactive) when all its orbitals are full. On the other hand, an atom is least stable (and therefore most reactive) when its valence electron orbitals are not full. This will make more sense when we go on to look at chemical bonding in a later chapter. To put it simply, the valence electrons are largely responsible for an element's chemical behaviour and elements that have the same number of valence electrons often have similar chemical properties.

One final point to note about electron configurations is stability. Which configurations are stable and which are not? Very simply, the most stable configurations are the ones that have full energy levels. These configurations occur in the noble gases. The noble gases are very stable elements that do not react easily (if at all) with any other elements. This is due to the full energy levels. All elements would like to reach the most stable electron configurations, i.e. all elements want to be noble gases. This principle of stability is sometimes referred to as the octet rule. An octet is a set of 8, and the number of electrons in a full energy level is 8.

The above experiment on flame tests relates to the line emission spectra of the metals. These line emission spectra are a direct result of the arrangement of the electrons in metals.