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Models and atomic size

Module by: Free High School Science Texts Project. E-mail the author

The following video covers some of the properties of an atom.

Figure 1
Veritasium video on the atom - 1

We have now looked at many examples of the types of matter and materials that exist around us and we have investigated some of the ways that materials are classified. But what is it that makes up these materials? And what makes one material different from another? In order to understand this, we need to take a closer look at the building block of matter - the atom. Atoms are the basis of all the structures and organisms in the universe. The planets, sun, grass, trees, air we breathe and people are all made up of different combinations of atoms.

Models of the Atom

It is important to realise that a lot of what we know about the structure of atoms has been developed over a long period of time. This is often how scientific knowledge develops, with one person building on the ideas of someone else. We are going to look at how our modern understanding of the atom has evolved over time.

The idea of atoms was invented by two Greek philosophers, Democritus and Leucippus in the fifth century BC. The Greek word ατoμoνατoμoν

 
(atom) means indivisible because they believed that atoms could not be broken into smaller pieces.

Nowadays, we know that atoms are made up of a positively charged nucleus in the centre surrounded by negatively charged electrons. However, in the past, before the structure of the atom was properly understood, scientists came up with lots of different models or pictures to describe what atoms look like.

Definition 1: Model

A model is a representation of a system in the real world. Models help us to understand systems and their properties. For example, an atomic model represents what the structure of an atom could look like, based on what we know about how atoms behave. It is not necessarily a true picture of the exact structure of an atom.

The Plum Pudding Model

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
Figure 2 (CG10C3_001.png)

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.

Rutherford's 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.
Figure 3 (CG10C3_002.png)

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.

Figure 4: Rutherford's model of the atom
Figure 4 (CG10C3_003.png)

The Bohr Model

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.

Note: Interesting Fact :

Light has the properties of both a particle and a wave! Einstein discovered that light comes in energy packets which are called photons. When an electron in an atom changes energy levels, a photon of light is emitted. This photon has the same energy as the difference between the two electron energy levels.

The size of atoms

It is difficult sometimes to imagine the size of an atom, or its mass, because we cannot see an atom and also because we are not used to working with such small measurements.

How heavy is an atom?

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 x 10-2610-26kg, while the mass of an atom of hydrogen is about 1.67 x 10-2710-27kg. 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.

Table 1: The atomic mass number of some of the elements
Element Atomic mass (u)
Carbon (C) 12
Nitrogen (N) 14
Bromine (Br) 80
Magnesium (Mg) 24
Potassium (K) 39
Calcium (Ca) 40
Oxygen (O) 16

The actual value of 1 atomic mass unit is 1.67 x 10-2410-24g or 1.67 x 10-2710-27kg. This is a very tiny mass!

How big is an atom?

Tip:

pm stands for picometres. 1 pm = 10-12-12 m

Atomic diameter also varies depending on the element. On average, the diameter of an atom ranges from 100 pm (Helium) to 670 pm (Caesium). Using different units, 100 pm = 1 Angstrom, and 1 Angstrom = 10-1010-10 m. That is the same as saying that 1 Angstrom = 0,0000000010 m or that 100 pm = 0,0000000010 m! In other words, the diameter of an atom ranges from 0.0000000010 m to 0.0000000067 m. This is very small indeed.

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