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Introduction

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.

Project: Models of the atom

Our current understanding of the atom came about over a long period of time, with many different people playing a role. Conduct some research into the development of the different ideas of the atom and the people who contributed to it. Some suggested people to look at are: JJ Thomson, Ernest Rutherford, Marie Curie, JC Maxwell, Max Planck, Albert Einstein, Niels Bohr, Lucretius, LV de Broglie, CJ Davisson, LH Germer, Chadwick, Werner Heisenberg, Max Born, Erwin Schrodinger, John Dalton, Empedocles, Leucippus, Democritus, Epicurus, Zosimos, Maria the Jewess, Geber, Rhazes, Robert Boyle, Henry Cavendish, A Lavoisier and H Becquerel. You do not need to find information on all these people, but try to find information about as many of them as possible.

Make a list of the key contributions to a model of the atom that each of these people made and then make a timeline of this information. (You can use an online tool such as Dipity to make a timeline.) Try to get a feel for how it all eventually fit together into the modern understanding of the atom.

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. This model is sometimes known as the planetary model of the atom.

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.

Other models of the atom

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.

Models of the atom

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

Table 1
Column A Column B
Discovery of electrons and the plum pudding model Niels Bohr
Arrangement of electrons Marie Curie and her husband, Pierre
Atoms as the smallest building block of matter Ancient Greeks
Discovery of the nucleus JJ Thomson
Discovery of radiation Rutherford
Click here for the solution

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 1,99×10-26kg1,99×10-26kg, 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.

Table 2: The atomic mass number of some of the elements
Element Atomic mass (u)
Carbon (CC) 12
Nitrogen (NN) 14
Bromine (BrBr) 80
Magnesium (MgMg) 24
Potassium (KK) 39
Calcium (CaCa) 40
Oxygen (OO) 16

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!

How big is an atom?

Tip:

pm stands for picometres. 1pm=10-12 m1pm=10-12m

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.

Atomic structure

As a result of the work done by previous scientists on atomic models (that we discussed in "Models of the Atom"), scientists now have a good idea of what an atom looks like. This knowledge is important because it helps us to understand why materials have different properties and why some materials bond with others. Let us now take a closer look at the microscopic structure of the atom.

So far, we have discussed that atoms are made up of a positively charged nucleus surrounded by one or more negatively charged electrons. These electrons orbit the nucleus.

The Electron

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.

The Nucleus

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 Proton

Each proton carries one unit of positive electric charge. Since we know that atoms are electrically neutral, i.e. do not carry any extra charge, then the number of protons in an atom has to be the same as the number of electrons to balance out the positive and negative charge to zero. The total positive charge of a nucleus is equal to the number of protons in the nucleus. The proton is much heavier than the electron (10 000 times heavier!) and has a mass of 1,6726×10-27kg1,6726×10-27kg. When we talk about the atomic mass of an atom, we are mostly referring to the combined mass of the protons and neutrons, i.e. the nucleons.

The Neutron

The neutron is electrically neutral i.e. it carries no charge at all. Like the proton, it is much heavier than the electron and its mass is 1,6749×10-27kg1,6749×10-27kg (slightly heavier than the proton).

Note: Interesting Fact :

Rutherford predicted (in 1920) that another kind of particle must be present in the nucleus along with the proton. He predicted this because if there were only positively charged protons in the nucleus, then it should break into bits because of the repulsive forces between the like-charged protons! Also, if protons were the only particles in the nucleus, then a helium nucleus (atomic number 2) would have two protons and therefore only twice the mass of hydrogen. However, it is actually four times heavier than hydrogen. This suggested that there must be something else inside the nucleus as well as the protons. To make sure that the atom stays electrically neutral, this particle would have to be neutral itself. In 1932 James Chadwick discovered the neutron and measured its mass.

Table 3: Summary of the particles inside the atom
  proton neutron electron
Mass (kg) 1,6726×10-271,6726×10-27 1,6749×10-271,6749×10-27 9,11×10-319,11×10-31
Units of charge +1+1 00 -1-1
Charge (C) 1,6×10-191,6×10-19 00 -1,6×10-19-1,6×10-19
Note: Interesting Fact :

Unlike the electron which is thought to be a point particle and unable to be broken up into smaller pieces, the proton and neutron can be divided. Protons and neutrons are built up of smaller particles called quarks. The proton and neutron are made up of 3 quarks each.

Atomic number and atomic mass number

The chemical properties of an element are determined by the charge of its nucleus, i.e. by the number of protons. This number is called the atomic number and is denoted by the letter Z.

Definition 2: Atomic number (Z)

The number of protons in an atom

You can find the atomic number on the periodic table. The atomic number is an integer and ranges from 1 to about 118.

The mass of an atom depends on how many nucleons its nucleus contains. The number of nucleons, i.e. the total number of protons plus neutrons, is called the atomic mass number and is denoted by the letter A.

Definition 3: Atomic mass number (A)

The number of protons and neutrons in the nucleus of an atom

Standard notation shows the chemical symbol, the atomic mass number and the atomic number of an element as follows:

Figure 5
Figure 5 (atom_sym.png)

Note:

A nuclide is a distinct kind of atom or nucleus characterized by the number of protons and neutrons in the atom. To be absolutely correct, when we represent atoms like we do here, then we should call them nuclides.

For example, the iron nucleus which has 26 protons and 30 neutrons, is denoted as:

26 56 Fe 26 56 Fe
(1)

where the atomic number is Z=26Z=26

 
and the mass number A=56A=56. The number of neutrons is simply the difference N=A-ZN=A-Z.

Tip:

Don't confuse the notation we have used above with the way this information appears on the Periodic Table. On the Periodic Table, the atomic number usually appears in the top lefthand corner of the block or immediately above the element's symbol. The number below the element's symbol is its relative atomic mass. This is not exactly the same as the atomic mass number. This will be explained in "Isotopes". The example of iron is shown below.

Figure 6
Figure 6 (CG10C3_004.png)

You will notice in the example of iron that the atomic mass number is more or less the same as its atomic mass. Generally, an atom that contains n nucleons (protons and neutrons), will have a mass approximately equal to nnu. For example the mass of a C-12C-12 atom which has 6 protons, 6 neutrons and 6 electrons is 12u, where the protons and neutrons have about the same mass and the electron mass is negligible.

Exercise 1

Use standard notation to represent sodium and give the number of protons, neutrons and electrons in the element.

Solution

  1. Step 1. Write the element symbol: Sodium is given by NaNa
  2. Step 2. Write down the number of protons: Sodium has 11 protons, so we have: 11Na11Na
  3. Step 3. Write down the number of neutrons: Sodium has 12 neutrons.
  4. Step 4. Work out A: A=N+Z=12+11=23A=N+Z=12+11=23
  5. Step 5. Write the answer: In standard notation sodium is given by: 1123Na1123Na. The number of protons is 11, the number of neutrons is 12 and the number of electrons is 11.

The structure of the atom

  1. Explain the meaning of each of the following terms:
    1. nucleus
    2. electron
    3. atomic mass
    Click here for the solution
  2. Complete the following table: (Note: You will see that the atomic masses on the Periodic Table are not whole numbers. This will be explained later. For now, you can round off to the nearest whole number.)
    Table 4
    ElementAtomic massAtomic numberNumber of protonsNumber of electronsNumber of neutrons
    MgMg2412   
    OO  8  
      17   
    NiNi   28 
     40   20
    ZnZn     
         0
    CC12  6 
    Click here for the solution
  3. Use standard notation to represent the following elements:
    1. potassium
    2. copper
    3. chlorine
    Click here for the solution
  4. For the element 1735Cl1735Cl, give the number of ...
    1. protons
    2. neutrons
    3. electrons
    ... in the atom.
    Click here for the solution
  5. Which of the following atoms has 7 electrons?
    1. 25He25He
    2. 613C613C
    3. 37Li37Li
    4. 715N715N
    Click here for the solution
  6. In each of the following cases, give the number or the element symbol represented by 'X'.
    1. 1840X1840X
    2. 20xCa20xCa
    3. x31Px31P
    Click here for the solution
  7. Complete the following table:
    Table 5
     AZN
    92235U92235U   
    92238U92238U   
    In these two different forms of Uranium...
    1. What is the same?
    2. What is different?
    Uranium can occur in different forms, called isotopes. You will learn more about isotopes in "Isotopes".
    Click here for the solution

Isotopes

What is an isotope?

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.

Note: Interesting Fact :

In Greek, “same place” reads as ι`σoςι`σoςτo`πoςτo`πoς

 
(isos topos). This is why atoms which have the same number of protons, but different numbers of neutrons, are called isotopes. They are in the same place on the Periodic Table!

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

Exercise 2: Isotopes

For the element 92234U92234U (uranium), use standard notation to describe:

  1. the isotope with 2 fewer neutrons
  2. the isotope with 4 more neutrons
Solution
  1. Step 1. Go over the definition of isotope :

    We know that isotopes of any element have the same number of protons (same atomic number) in each atom, which means that they have the same chemical symbol. However, they have a different number of neutrons, and therefore a different mass number.

  2. Step 2. Rewrite the notation for the isotopes :

    Therefore, any isotope of uranium will have the symbol:

    U U
    (2)

    Also, since the number of protons in uranium isotopes is always the same, we can write down the atomic number:

    92 U 92 U
    (3)

    Now, if the isotope we want has 2 fewer neutrons than 92234U92234U, then we take the original mass number and subtract 2, which gives:

    92 232 U 92 232 U
    (4)

    Following the steps above, we can write the isotope with 4 more neutrons as:

    92 238 U 92 238 U
    (5)

Exercise 3: Isotopes

Which of the following are isotopes of 2040 Ca 2040 Ca ?

  • 19 40 K 19 40 K
  • 20 42 Ca 20 42 Ca
  • 18 40 Ar 18 40 Ar
Solution
  1. Step 1. Go over the definition of isotope: :

    We know that isotopes have the same atomic number but different mass numbers.

  2. Step 2. Determine which of the elements listed fits the definition of an isotope. :

    You need to look for the element that has the same atomic number but a different atomic mass number. The only element is 2042Ca2042Ca. What is different is that there are 2 more neutrons than in the original element.

Exercise 4: Isotopes

For the sulphur isotope 1633S1633S, give the number of...

  1. protons
  2. nucleons
  3. electrons
  4. neutrons
Solution
  1. Step 1. Determine the number of protons by looking at the atomic number, Z. :

    Z=16Z=16, therefore the number of protons is 16 (answer to (a)).

  2. Step 2. Determine the number of nucleons by looking at the atomic mass number, A. :

    A=33A=33, therefore the number of nucleons is 33 (answer to (b)).

  3. Step 3. Determine the number of electrons :

    The atom is neutral, and therefore the number of electrons is the same as the number of protons. The number of electrons is 16 (answer to (c)).

  4. Step 4. Calculate the number of neutrons :
    N = A - Z = 33 - 16 = 17 N = A - Z = 33 - 16 = 17
    (6)

    The number of neutrons is 17 (answer to (d)).

Isotopes

  1. Atom A has 5 protons and 5 neutrons, and atom B has 6 protons and 5 neutrons. These atoms are...
    1. allotropes
    2. isotopes
    3. isomers
    4. atoms of different elements
    Click here for the solution
  2. For the sulphur isotopes, 1632S1632S and 1634S1634S, give the number of...
    1. protons
    2. nucleons
    3. electrons
    4. neutrons
    Click here for the solution
  3. Which of the following are isotopes of 1735Cl1735Cl?
    1. 3517Cl3517Cl
    2. 1735Cl1735Cl
    3. 1737Cl1737Cl
    Click here for the solution
  4. Which of the following are isotopes of U-235U-235? (X represents an element symbol)
    1. 92238X92238X
    2. 90238X90238X
    3. 92235X92235X
    Click here for the solution

Relative atomic mass

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.

Note: Interesting fact:

The relative atomic mass of some elements depends on where on Earth the element is found. This is because the isotopes can be found in varying ratios depending on certain factors such as geological composition, etc. The International Union of Pure and Applied Chemistry (IUPAC) has decided to give the relative atomic mass of some elements as a range to better represent the varying isotope ratios on the Earth. For the calculations that you will do at high school, it is enough to simply use one number without worrying about these ranges.

Exercise 5: The relative atomic mass of an isotopic element

The element chlorine has two isotopes, chlorine-35 and chlorine-37. The abundance of these isotopes when they occur naturally is 75% chlorine-35 and 25% chlorine-37. Calculate the average relative atomic mass for chlorine.

Solution
  1. Step 1. Calculate the mass contribution of chlorine-35 to the average relative atomic mass :

    Contribution of Cl-35=(75100×35)=26,25uCl-35=(75100×35)=26,25u

  2. Step 2. Calculate the contribution of chlorine-37 to the average relative atomic mass :

    Contribution of Cl-37=(25100×37)=9,25uCl-37=(25100×37)=9,25u

  3. Step 3. Add the two values to arrive at the average relative atomic mass of chlorine :

    Relative atomic mass of chlorine=26,25u+9,25u=35,5uRelative atomic mass of chlorine=26,25u+9,25u=35,5u

    If you look on the periodic table, the average relative atomic mass for chlorine is 35,5u35,5u. You will notice that for many elements, the relative atomic mass that is shown is not a whole number. You should now understand that this number is the average relative atomic mass for those elements that have naturally occurring isotopes.

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

Figure 7
Figure 7 (isotope.png)
run demo

Isotopes

  1. Complete the table below:
    Table 6
    IsotopeZAProtonsNeutronsElectrons
    Carbon-12     
    Carbon-14     
    Chlorine-35     
    Chlorine-37     
    Click here for the solution
  2. If a sample contains 90% carbon-12 and 10% carbon-14, calculate the relative atomic mass of an atom in that sample.
     
    Click here for the solution
  3. If a sample contains 22,5% Cl-37 and 77,5% Cl-35, calculate the relative atomic mass of an atom in that sample.
     
    Click here for the solution

Group Discussion : The changing nature of scientific knowledge

Scientific knowledge is not static: it changes and evolves over time as scientists build on the ideas of others to come up with revised (and often improved) theories and ideas. In this chapter for example, we saw how peoples' understanding of atomic structure changed as more information was gathered about the atom. There are many more examples like this one in the field of science. For example, think about our knowledge of the solar system and the origin of the universe, or about the particle and wave nature of light.

Often, these changes in scientific thinking can be very controversial because they disturb what people have come to know and accept. It is important that we realise that what we know now about science may also change. An important part of being a scientist is to be a critical thinker. This means that you need to question information that you are given and decide whether it is accurate and whether it can be accepted as true. At the same time, you need to learn to be open to new ideas and not to become stuck in what you believe is right... there might just be something new waiting around the corner that you have not thought about!

In groups of 4-5, discuss the following questions:

  • Think about some other examples where scientific knowledge has changed because of new ideas and discoveries:
    • What were these new ideas?
    • Were they controversial? If so, why?
    • What role (if any) did technology play in developing these new ideas?
    • How have these ideas affected the way we understand the world?
  • Many people come up with their own ideas about how the world works. The same is true in science. So how do we, and other scientists, know what to believe and what not to? How do we know when new ideas are 'good' science or 'bad' science? In your groups, discuss some of the things that would need to be done to check whether a new idea or theory was worth listening to, or whether it was not.
  • Present your ideas to the rest of the class.

Electron configuration

The energy of electrons

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.

Energy quantisation and line emission spectra (Not in CAPS, included for completeness)

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.

Electron configuration

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.

  1. Lithium Lithium (Li) has an atomic number of 3, meaning that in a neutral atom, the number of electrons will also be 3. The first two electrons are found in the first energy level, while the third electron is found in the second energy level (Figure 8).
    Figure 8: The arrangement of electrons in a lithium atom.
    Figure 8 (CG10C3_005.png)
  2. Fluorine Fluorine (FF) has an atomic number of 9, meaning that a neutral atom also has 9 electrons. The first 2 electrons are found in the first energy level, while the other 7 are found in the second energy level (Figure 9).
    Figure 9: The arrangement of electrons in a fluorine atom.
    Figure 9 (CG10C3_006.png)
  3. Argon Argon has an atomic number of 18, meaning that a neutral atom also has 18 electrons. The first 2 electrons are found in the first energy level, the next 8 are found in the second energy level, and the last 8 are found in the third energy level (Figure 10).
    Figure 10: The arrangement of electrons in an argon atom.
    Figure 10 (CG10C3_007.png)

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.
Figure 11 (orbitals.png)

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.
Figure 12 (CG10C3_008.png)

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:

  • Each orbital can only hold two electrons. Electrons that occur together in an orbital are called an electron pair.
  • An electron will always try to enter an orbital with the lowest possible energy.
  • An electron will occupy an orbital on its own, rather than share an orbital with another electron. An electron would also rather occupy a lower energy orbital with another electron, before occupying a higher energy orbital. In other words, within one energy level, electrons will fill an 's' orbital before starting to fill 'p' orbitals.
  • The s subshell can hold 2 electrons
  • The p subshell can hold 6 electrons

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:

  1. Determine the number of electrons that the atom has.
  2. Fill the 's' orbital in the first energy level (the 1s1s orbital) with the first two electrons.
  3. Fill the 's' orbital in the second energy level (the 2s2s orbital) with the second two electrons.
  4. Put one electron in each of the three 'p' orbitals in the second energy level (the 2p2p orbitals) and then if there are still electrons remaining, go back and place a second electron in each of the 2p2p orbitals to complete the electron pairs.
  5. Carry on in this way through each of the successive energy levels until all the electrons have been drawn.

Tip:

When there are two electrons in an orbital, the electrons are called an electron pair. If the orbital only has one electron, this electron is said to be an unpaired electron. Electron pairs are shown with arrows pointing in opposite directions. You may hear people talking of the Pauli exclusion principle. This principle says that electrons have a property known as spin and two electrons in an orbital will not spin the same way. This is why we use arrows pointing in opposite directions. An arrow pointing up denotes an electron spinning one way and an arrow pointing downwards denotes an electron spinning the other way.

Note: Interesting fact:

Aufbau is the German word for 'building up'. Scientists used this term since this is exactly what we are doing when we work out electron configuration, we are building up the atoms structure.

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
Figure 13 (CG10C3_009.png)

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 14 (CG10C3_010.png)
Figure 15: An Aufbau diagram showing the electron configuration of argon
Figure 15 (CG10C3_011.png)

Exercise 6: Aufbau diagrams

Give the electron configuration for sodium (NaNa) and draw an aufbau diagram.

Solution

  1. Step 1. Write down the number of electrons: Sodium has 11 electrons.
  2. Step 2. Work out which orbitals to fill: We start by placing two electrons in the 1s1s orbital: 1s21s2. Now we have 9 electrons left to place in orbitals, so we put two in the 2s2s orbital: 2s22s2. There are now 7 electrons to place in orbitals so we place 6 of them in the 2p2p orbital: 2p62p6. The last electron goes into the 3s3s orbital: 3s13s1.
  3. Step 3. Write down the electron configuration: The electron configuration is: 1s22s22p63s11s22s22p63s1
  4. Step 4. Draw the Aufbau diagram: Using the electron configuration we get the following diagram:
    Figure 16
    Figure 16 (wexaufbau.png)

Core and valence electrons

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

The importance of understanding electron configuration

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.

Experiment: Flame tests

Aim:


To determine what colour a metal cation will cause a flame to be.

Apparatus:


Watch glass, bunsen burner, methanol, bamboo sticks, metal salts (e.g. NaClNaCl, CuCl2CuCl2, CaCl2CaCl2, KClKCl, etc. ) and metal powders (e.g. copper, magnesium, zinc, iron, etc.)

Method:


For each salt or powder do the following:

  1. Dip a clean bamboo stick into the methanol
  2. Dip the stick into the salt or powder
  3. Wave the stick through the flame from the bunsen burner. DO NOT hold the stick in the flame, but rather wave it back and forth through the flame.
  4. Observe what happens

Results:


Record your results in a table, listing the metal salt and the colour of the flame.

Conclusion:


You should have observed different colours for each of the metal salts and powders that you tested.

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.

Energy diagrams and electrons

  1. Draw Aufbau diagrams to show the electron configuration of each of the following elements:
    1. magnesium
    2. potassium
    3. sulphur
    4. neon
    5. nitrogen
  2. Use the Aufbau diagrams you drew to help you complete the following table:
    Table 7
    ElementNo. of energy levelsNo. of core electronsNo. of valence electronsElectron configuration (standard notation)
    MgMg    
    KK    
    SS    
    NeNe    
    NN    
  3. Rank the elements used above in order of increasing reactivity. Give reasons for the order you give. Click here for the answer

Group work : Building a model of an atom

Earlier in this chapter, we talked about different 'models' of the atom. In science, one of the uses of models is that they can help us to understand the structure of something that we can't see. In the case of the atom, models help us to build a picture in our heads of what the atom looks like.

Models are often simplified. The small toy cars that you may have played with as a child are models. They give you a good idea of what a real car looks like, but they are much smaller and much simpler. A model cannot always be absolutely accurate and it is important that we realise this so that we don't build up a false idea about something.

In groups of 4-5, you are going to build a model of an atom. Before you start, think about these questions:

  • What information do I know about the structure of the atom? (e.g. what parts make it up? how big is it?)
  • What materials can I use to represent these parts of the atom as accurately as I can?
  • How will I put all these different parts together in my model?

As a group, share your ideas and then plan how you will build your model. Once you have built your model, discuss the following questions:

  • Does our model give a good idea of what the atom actually looks like?
  • In what ways is our model inaccurate? For example, we know that electrons move around the atom's nucleus, but in your model, it might not have been possible for you to show this.
  • Are there any ways in which our model could be improved?

Now look at what other groups have done. Discuss the same questions for each of the models you see and record your answers.

The following simulation allows you to build an atom
run demo

Figure 17: Build an atom simulation
Figure 17 (atom1.png)

This is another simulation that allows you to build an atom. This simulation also provides a summary of what you have learnt so far.
Run demo

Figure 18: Build an atom simulation 2
Figure 18 (atom2.png)

Summary

  • Much of what we know today about the atom, has been the result of the work of a number of scientists who have added to each other's work to give us a good understanding of atomic structure.
  • Some of the important scientific contributors include J.J.Thomson (discovery of the electron, which led to the Plum Pudding Model of the atom), Ernest Rutherford (discovery that positive charge is concentrated in the centre of the atom) and Niels Bohr (the arrangement of electrons around the nucleus in energy levels).
  • Because of the very small mass of atoms, their mass is measured in atomic mass units (u). 1u= 1,67×10-24g1u=1,67×10-24g.
  • An atom is made up of a central nucleus (containing protons and neutrons), surrounded by electrons.
  • The atomic number (Z) is the number of protons in an atom.
  • The atomic mass number (A) is the number of protons and neutrons in the nucleus of an atom.
  • The standard notation that is used to write an element, is ZAXZAX, where X is the element symbol, A is the atomic mass number and Z is the atomic number.
  • 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. This means that not all atoms of an element will have the same atomic mass.
  • The relative atomic mass of an element is the average mass of one atom of all the naturally occurring isotopes of a particular chemical element, expressed in atomic mass units. The relative atomic mass is written under the elements' symbol on the Periodic Table.
  • The energy of electrons in an atom is quantised. Electrons occur in specific energy levels around an atom's nucleus.
  • Within each energy level, an electron may move within a particular shape of orbital. An orbital defines the space in which an electron is most likely to be found. There are different orbital shapes, including s, p, d and f orbitals.
  • Energy diagrams such as Aufbau diagrams are used to show the electron configuration of atoms.
  • The electrons in the outermost energy level are called valence electrons.
  • The electrons that are not valence electrons are called core electrons.
  • Atoms whose outermost energy level is full, are less chemically reactive and therefore more stable, than those atoms whose outer energy level is not full.

Figure 19

End of chapter exercises

  1. Write down only the word/term for each of the following descriptions.
    1. The sum of the number of protons and neutrons in an atom
    2. The defined space around an atom's nucleus, where an electron is most likely to be found
    Click here for the solution
  2. For each of the following, say whether the statement is True or False. If it is False, re-write the statement correctly.
    1. 1020Ne1020Ne and 1022Ne1022Ne each have 10 protons, 12 electrons and 12 neutrons.
    2. The atomic mass of any atom of a particular element is always the same.
    3. It is safer to use helium gas rather than hydrogen gas in balloons.
    4. Group 1 elements readily form negative ions.
    Click here for the solution
  3. Multiple choice questions: In each of the following, choose the one correct answer.
    1. The three basic components of an atom are:
      1. protons, neutrons, and ions
      2. protons, neutrons, and electrons
      3. protons, neutrinos, and ions
      4. protium, deuterium, and tritium
      Click here for the solution
    2. The charge of an atom is...
      1. positive
      2. neutral
      3. negative
      Click here for the solution
    3. If Rutherford had used neutrons instead of alpha particles in his scattering experiment, the neutrons would...
      1. not deflect because they have no charge
      2. have deflected more often
      3. have been attracted to the nucleus easily
      4. have given the same results
      Click here for the solution
    4. Consider the isotope 92234U92234U. Which of the following statements is true?
      1. The element is an isotope of 94234Pu94234Pu
      2. The element contains 234 neutrons
      3. The element has the same electron configuration as 92238U92238U
      4. The element has an atomic mass number of 92
      Click here for the solution
    5. The electron configuration of an atom of chlorine can be represented using the following notation:
      1. 1s22s83s71s22s83s7
      2. 1s22s22p63s23p51s22s22p63s23p5
      3. 1s22s22p63s23p61s22s22p63s23p6
      4. 1s22s22p51s22s22p5
      Click here for the solution
  4. Give the standard notation for the following elements:
    1. beryllium
    2. carbon-12
    3. titanium-48
    4. fluorine
    Click here for the solution
  5. Give the electron configurations and aufbau diagrams for the following elements:
    1. aluminium
    2. phosphorus
    3. carbon
    Click here for the solution
  6. Use standard notation to represent the following elements:
    1. argon
    2. calcium
    3. silver-107
    4. bromine-79
    Click here for the solution
  7. For each of the following elements give the number of protons, neutrons and electrons in the element:
    1. 78195Pt78195Pt
    2. 1840Ar1840Ar
    3. 2759Co2759Co
    4. 37Li37Li
    5. 511B511B
    Click here for the solution
  8. For each of the following elements give the element or number represented by 'x':
    1. 45103X45103X
    2. x35Clx35Cl
    3. 4xBe4xBe
    Click here for the solution
  9. Which of the following are isotopes of 1224Mg1224Mg:
    1. 2512Mg2512Mg
    2. 1226Mg1226Mg
    3. 1324Al1324Al
    Click here for the solution
  10. If a sample contains 69% of copper-63 and 31% of copper-65, calculate the relative atomic mass of an atom in that sample.
    Click here for the solution
  11. Complete the following table:
    Table 8
    Element Electron configuration Core electrons Valence electrons
    Boron (B)      
    Calcium (Ca)      
    Silicon (Si)      
    Lithium (Li)      
    Neon (Ne)      
    Click here for the solution
  12. Draw aufbau diagrams for the following elements:
    1. beryllium
    2. sulphur
    3. argon
    Click here for the solution

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