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Atomic Nature of Matter

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, the sun, grass and trees, the 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 chargednucleus in the centre surrounded by negatively chargedelectrons. 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 1). 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 1: A schematic diagram to show what the atom looks like according to the Plum Pudding model
Figure 1 (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 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 2.

Figure 2: 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 2 (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 sulfide. 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 3.

Figure 3: Rutherford's model of the atom
Figure 3 (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.

How big is an atom?

It is difficult sometimes to imagine the size of an atom, or its mass, because we cannot see them, 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'. If we give carbon an atomic mass of 12 u, then 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 other atoms in the Periodic Table. The atomic masses of some elements are shown in Table 1 below.

Table 1: The atomic mass of a number of elements
Element Atomic mass (u)
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.

Atomic structure

As a result of the 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 x 10-3110-31 kg. 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 x 10-1910-19 C (Coulombs).

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 x 10-2710-27 kg. 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 x 10-2710-27 kg (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 2: Summary of the particles inside the atom
  proton neutron electron
Mass (kg) 1.6726 x 10-2710-27 1.6749 x 10-2710-27 9.11 x 10-3110-31
Units of charge +1 0 -1
Charge (C) 1.6 x 10-1910-19 0 -1.6 x 10-1910-19
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

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:

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

26 56 Fe 26 56 Fe
(1)

where the total nuclear charge 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 again below.

Figure 4
Figure 4 (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-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.

The structure of the atom

  1. Explain the meaning of each of the following terms:
    1. nucleus
    2. electron
    3. atomic mass
  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 3
    ElementAtomic massAtomic numberNumber of protonsNumber of electronsNumber of neutrons
    Mg2412   
    O  8  
      17   
    Ni   28 
     40   20
    Zn     
         0
    C12  6 
  3. Use standard notation to represent the following elements:
    1. potassium
    2. copper
    3. chlorine
  4. For the element 17351735Cl, give the number of ...
    1. protons
    2. neutrons
    3. electrons
    ... in the atom.
  5. Which of the following atoms has 7 electrons?
    1. 2525He
    2. 613613C
    3. 3737Li
    4. 715715N
  6. In each of the following cases, give the number or the element symbol represented by 'X'.
    1. 18401840X
    2. 20x20xCa
    3. x31x31P
  7. Complete the following table:
    Table 4
     AZN
    9223592235U   
    9223892238U   
    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".

Isotopes

What is an isotope?

If a few neutrons are added to or removed from a nucleus, the chemical properties of the atom will stay the same because its charge is still the same. Therefore, the chemical properties of an element depend on the number of protons inside the atom. This means that such an atom should 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 C (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 orginal 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 if an element is written for example as C-12, the '12' is the atomic mass of that atom. So, Cl-35 has an atomic mass of 35 u, while Cl-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 1: 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 2: 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 20422042Ca. What is different is that there are 2 more neutrons than in the original element.

Exercise 3: Isotopes

For the sulfur isotope 16331633S, 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 = 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 = 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
  2. For the sulfur isotopes, 16321632S and 16341634S, give the number of...
    1. protons
    2. nucleons
    3. electrons
    4. neutrons
  3. Which of the following are isotopes of 3535Cl?
    1. 35173517Cl
    2. 17351735Cl
    3. 17371737Cl
  4. Which of the following are isotopes of U-235? (X represents an element symbol)
    1. 9223892238X
    2. 9023890238X
    3. 9223592235X

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 then the number that will appear 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.

Exercise 4: 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 = (75/100 x 35) = 26,25 u

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

    Contribution of Cl-37 = (25/100 x 37) = 9,25 u

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

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

    If you look on the periodic table, the average relative atomic mass for chlorine is 35,5 u. 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.

Isotopes

You are given a sample that contains carbon-12 and carbon-14.

  1. Complete the table below:
    Table 5
    IsotopeZAProtonsNeutronsElectrons
    Carbon-12     
    Carbon-14     
    Chlorine-35     
    Chlorine-37     
    1. If the sample you have contains 90% carbon-12 and 10% carbon-14, calculate the relative atomic mass of an atom in that sample.
    2. In another sample, you have 22.5% Cl-37 and 77.5% Cl-35. Calculate the relative atomic mass of an atom in that sample.

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.

The Arrangement of Atoms in the Periodic Table

The periodic table of the elements is a tabular method of showing the chemical elements. Most of the work that was done to arrive at the periodic table that we know, can be attributed to a man called Dmitri Mendeleev in 1869. Mendeleev was a Russian chemist who designed the table in such a way that recurring ("periodic") trends in the properties of the elements could be shown. Using the trends he observed, he even left gaps for those elements that he thought were 'missing'. He even predicted the properties that he thought the missing elements would have when they were discovered. Many of these elements were indeed discovered and Mendeleev's predictions were proved to be correct.

To show the recurring properties that he had observed, Mendeleev began new rows in his table so that elements with similar properties were in the same vertical columns, called groups. Each row was referred to as a period. One important feature to note in the periodic table is that all the non-metals are to the right of the zig-zag line drawn under the element boron. The rest of the elements are metals, with the exception of hydrogen which occurs in the first block of the table despite being a non-metal.

Figure 5: A simplified diagram showing part of the Periodic Table
Figure 5 (CG10C3_014.png)

Groups in the periodic table

A group is a vertical column in the periodic table, and is considered to be the most important way of classifying the elements. If you look at a periodic table, you will see the groups numbered at the top of each column. The groups are numbered from left to right as follows: 1, 2, then an open space which contains the transition elements, followed by groups 3 to 8. These numbers are normally represented using Roman numerals. In some periodic tables, all the groups are numbered from left to right from number 1 to number 18. In some groups, the elements display very similar chemical properties, and the groups are even given separate names to identify them.

The characteristics of each group are mostly determined by the electron configuration of the atoms of the element.

  • Group 1: These elements are known as the alkali metals and they are very reactive.
  • Group 2: These elements are known as the alkali earth metals. Each element only has two valence electrons and so in chemical reactions, the group 2 elements tend to lose these electrons so that the energy shells are complete. These elements are less reactive than those in group 1 because it is more difficult to lose two electrons than it is to lose one.
  • Group 3 elements have three valence electrons.

    Tip:

    The number of valence electrons of an element corresponds to its group number on the periodic table.
  • Group 7: These elements are known as the halogens. Each element is missing just one electron from its outer energy shell. These elements tend to gain electrons to fill this shell, rather than losing them.
  • Group 8: These elements are the noble gases. All of the energy shells of the halogens are full, and so these elements are very unreactive.
    Figure 6: Electron diagrams for two of the noble gases, helium (He) and neon (Ne).
    Figure 6 (CG10C3_016.png)
  • Transition metals: The differences between groups in the transition metals are not usually dramatic.

It is worth noting that in each of the groups described above, the atomic diameter of the elements increases as you move down the group. This is because, while the number of valence electrons is the same in each element, the number of core electrons increases as one moves down the group.

Investigation : The properties of elements

Refer to Figure 7.

  1. Use a Periodic Table to help you to complete the last two diagrams for sodium (Na) and potassium (K).
  2. What do you notice about the number of electrons in the valence energy level in each case?
  3. Explain why elements from group 1 are more reactive than elements from group 2 on the periodic table (Hint: Think back to 'ionisation energy').
Figure 7: Electron diagrams for some of the Group 1 elements, with sodium and potasium incomplete; to be completed as an excersise.
Figure 7 (CG10C3_015.png)

Periods in the periodic table

A period is a horizontal row in the periodic table of the elements. Some of the trends that can be observed within a period are highlighted below:

  • As you move from one group to the next within a period, the number of valence electrons increases by one each time.
  • Within a single period, all the valence electrons occur in the same energy shell. If the period increases, so does the energy shell in which the valence electrons occur.
  • In general, the diameter of atoms decreases as one moves from left to right across a period. Consider the attractive force between the positively charged nucleus and the negatively charged electrons in an atom. As you move across a period, the number of protons in each atom increases. The number of electrons also increases, but these electrons will still be in the same energy shell. As the number of protons increases, the force of attraction between the nucleus and the electrons will increase and the atomic diameter will decrease.
  • Ionisation energy increases as one moves from left to right across a period. As the valence electron shell moves closer to being full, it becomes more difficult to remove electrons. The opposite is true when you move down a group in the table because more energy shells are being added. The electrons that are closer to the nucleus 'shield' the outer electrons from the attractive force of the positive nucleus. Because these electrons are not being held to the nucleus as strongly, it is easier for them to be removed and the ionisation energy decreases.
  • In general, the reactivity of the elements decreases from left to right across a period.

Trends in ionisation energy

Refer to the data table below which gives the ionisation energy (in kJ.mol-1-1) and atomic number (Z) for a number of elements in the periodic table:

Table 6
Z Ionisation energy Z Ionisation energy
1 1310 10 2072
2 2360 11 494
3 517 12 734
4 895 13 575
5 797 14 783
6 1087 15 1051
7 1397 16 994
8 1307 17 1250
9 1673 18 1540
  1. Draw a line graph to show the relationship between atomic number (on the x-axis) and ionisation energy (y-axis).
  2. Describe any trends that you observe.
  3. Explain why...
    1. the ionisation energy for Z=2 is higher than for Z=1
    2. the ionisation energy for Z=3 is lower than for Z=2
    3. the ionisation energy increases between Z=5 and Z=7

Elements in the Periodic Table

Refer to the elements listed below:

Lithium (Li); Chlorine (Cl); Magnesium (Mg); Neon (Ne); Oxygen (O); Calcium (Ca); Carbon (C)

Which of the elements listed above:

  1. belongs to Group 1
  2. is a halogen
  3. is a noble gas
  4. is an alkali metal
  5. has an atomic number of 12
  6. has 4 neutrons in the nucleus of its atoms
  7. contains electrons in the 4th energy level
  8. has only one valence electron
  9. has all its energy orbitals full
  10. will have chemical properties that are most similar
  11. will form positive ions

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 in measured in atomic mass units (u). 1 u = 1,67 ×× 10-24-24 g.
  • 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 ZAZAX, 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.
  • An ion is a charged atom. A cation is a positively charged ion and an anion is a negatively charged ion.
  • When forming an ion, an atom will lose or gain the number of electrons that will make its valence energy level full.
  • An element's ionisation energy is the energy that is needed to remove one electron from an atom.
  • Ionisation energy increases across a period in the Periodic Table.
  • Ionisation energy decreases down a group in the Periodic Table.

Summary

  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
  2. For each of the following, say whether the statement is True or False. If it is False, re-write the statement correctly.
    1. 10201020Ne and 10221022Ne 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.
  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
    2. The charge of an atom is...
      1. positive
      2. neutral
      3. negative
    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
    4. Consider the isotope 9223492234U. Which of the following statements is true?
      1. The element is an isotope of 9423494234Pu
      2. The element contains 234 neutrons
      3. The element has the same electron configuration as 9223892238U
      4. The element has an atomic mass number of 92
    5. The electron configuration of an atom of chlorine can be represented using the following notation:
      1. 1s22 2s88 3s77
      2. 1s22 2s22 2p66 3s22 3p55
      3. 1s22 2s22 2p66 3s22 3p66
      4. 1s22 2s22 2p55
  4. The following table shows the first ionisation energies for the elements of period 1 and 2.
    Table 7
    PeriodElementFirst ionisation energy (kJ.mol-1kJ.mol-1)
    1H1312
     He2372
     Li520
     Be899
     B801
     C1086
    2N1402
     O1314
     F1681
     Ne2081
    1. What is the meaning of the term first ionisation energy?
    2. Identify the pattern of first ionisation energies in a period.
    3. Which TWO elements exert the strongest attractive forces on their electrons? Use the data in the table to give a reason for your answer.
    4. Draw Aufbau diagrams for the TWO elements you listed in the previous question and explain why these elements are so stable.
    5. It is safer to use helium gas than hydrogen gas in balloons. Which property of helium makes it a safer option?
    6. 'Group 1 elements readily form positive ions'. Is this statement correct? Explain your answer by referring to the table.

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