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Ionisation energy and trends

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

Ionisation Energy and the Periodic Table

Ions

In the previous section, we focused our attention on the electron configuration of neutral atoms. In a neutral atom, the number of protons is the same as the number of electrons. But what happens if an atom gains or loses electrons? Does it mean that the atom will still be part of the same element?

A change in the number of electrons of an atom does not change the type of atom that it is. However, the charge of the atom will change. If electrons are added, then the atom will become more negative. If electrons are taken away, then the atom will become more positive. The atom that is formed in either of these cases is called an ion. Put simply, an ion is a charged atom.

Definition 1: Ion

An ion is a charged atom. A positively charged ion is called a cation e.g. Na+Na+, and a negatively charged ion is called an anion e.g. F-F-. The charge on an ion depends on the number of electrons that have been lost or gained.

But how do we know how many electrons an atom will gain or lose? Remember what we said about stability? We said that all atoms are trying to get a full outer shell. For the elements on the left hand side of the periodic table the easiest way to do this is to lose electrons and for the elements on the right of the periodic table the easiest way to do this is to gain electrons. So the elements on the left of the periodic table will form cations and the elements on the right hand side of the periodic table will form anions. By doing this the elements can be in the most stable electronic configuration and so be as stable as the noble gases.

Look at the following examples. Notice the number of valence electrons in the neutral atom, the number of electrons that are lost or gained and the final charge of the ion that is formed.

Lithium

A lithium atom loses one electron to form a positive ion:

Figure 1: The arrangement of electrons in a lithium ion.
Figure 1 (CG10C3_012.png)
In this example, the lithium atom loses an electron to form the cation Li+Li+.

Fluorine

A fluorine atom gains one electron to form a negative ion:

Figure 2: The arrangement of electrons in a fluorine ion.
Figure 2 (CG10C3_013.png)

You should have noticed in both these examples that each element lost or gained electrons to make a full outer shell.

Investigation : The formation of ions

  1. Use the diagram for lithium as a guide and draw similar diagrams to show how each of the following ions is formed:
    1. Mg2+Mg2+
    2. Na+Na+
    3. Cl-Cl-
    4. O2+O2+
  2. Do you notice anything interesting about the charge on each of these ions? Hint: Look at the number of valence electrons in the neutral atom and the charge on the final ion.

Observations:


Once you have completed the activity, you should notice that:

  • In each case the number of electrons that is either gained or lost, is the same as the number of electrons that are needed for the atoms to achieve a full outer energy level.
  • If you look at an energy level diagram for sodium (NaNa), you will see that in a neutral atom, there is only one valence electron. In order to achieve a full outer energy level, and therefore a more stable state for the atom, this electron will be lost.
  • In the case of oxygen (OO), there are six valence electrons. To achieve a full energy level, it makes more sense for this atom to gain two electrons. A negative ion is formed.

Exercise: The formation of ions

Match the information in column A with the information in column B by writing only the letter (A to I) next to the question number (1 to 7)

Table 1
1. A positive ion that has 3 less electrons than its neutral atom A. Mg2+Mg2+
2. An ion that has 1 more electron than its neutral atom B. Cl-Cl-
3. The anion that is formed when bromine gains an electron C. CO32-CO32-
4. The cation that is formed from a magnesium atom D. Al3+Al3+
5. An example of a compound ion E. Br2-Br2-
6. A positive ion with the electron configuration of argon F. K+K+
7. A negative ion with the electron configuration of neon G. Mg+Mg+
  H. O2-O2-
  I. Br-Br-
Click here for the solution

Ionisation Energy

Ionisation energy is the energy that is needed to remove one electron from an atom in the gas phase. The ionisation energy will be different for different atoms.

Tip:

When we talk of ionisation energies and calculate these energies the atoms or molecules involved are in the gas phase.

The second ionisation energy is the energy that is needed to remove a second electron from an atom, and so on. As an energy level becomes more full, it becomes more and more difficult to remove an electron and the ionisation energy increases. On the Periodic Table of the Elements, a group is a vertical column of the elements, and a period is a horizontal row. In the periodic table, ionisation energy increases across a period, but decreases as you move down a group. The lower the ionisation energy, the more reactive the element will be because there is a greater chance of electrons being involved in chemical reactions. We will look at this in more detail in the next section.

Trends in ionisation energy

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

Table 2
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=2Z=2 is higher than for Z=1Z=1
    2. the ionisation energy for Z=3Z=3 is lower than for Z=2Z=2
    3. the ionisation energy increases between Z=5Z=5 and Z=7Z=7
Click here for the solution

Figure 3
Khan academy video on periodic table - 2

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. Note that although hydrogen appears in group 1, it is not an alkali metal.
    Figure 4: Electron diagrams for some of the Group 1 elements, with sodium and potasium incomplete; to be completed as an excersise.
    Figure 4 (CG10C3_015.png)
  • 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 13 elements have three valence electrons.
  • Group 16: These elements are sometimes known as the chalcogens. These elements are fairly reactive and tend to gain electrons to fill their outer shell.
  • Group 17: 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. These elements are also very reactive.
  • Group 18: These elements are the noble gases. All of the energy shells of the halogens are full and so these elements are very unreactive.
    Figure 5: Electron diagrams for two of the noble gases, helium (HeHe) and neon (NeNe).
    Figure 5 (CG10C3_016.png)
  • Transition metals: The differences between groups in the transition metals are not usually dramatic.

Note: Interestng fact:

Group 15 on the periodic table is sometimes called the pnictogens.

Investigation : The properties of elements

Refer to Figure 4.

  1. Use a periodic table to help you to complete the last two diagrams for sodium (NaNa) and potassium (KK).
  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 about the 'ionisation energy').

You should also be able to indicate where metals, non-metals and metalloids are found on the periodic table. If you do not recall where these lie, then refer to classification of matter.

By now you should have an appreciation of what the periodic table can tell us. The periodic table does not just list the elements, but tells chemists what the properties of elements are, how the elements will combine and many other useful facts. The periodic table is truly an amazing resource. Into one simple table, chemists have packed so many facts and data that can easily be seen with a glance. The periodic table is a crucial part of chemistry and you should never go to science class without it.

The following presentation provides a summary of the periodic table

Figure 6

Summary

  • Elements are arranged in periods and groups on the periodic table. The elements are arranged according to increasing atomic number.
  • A group is a column on the periodic table containing elements with similar properties. A period is a row on the periodic table.
  • The groups on the periodic table are labeled from 1 to 8. The first group is known as the alkali metals, the second group is known as the alkali earth metals, the seventh group is known as the halogens and the eighth group is known as the noble gases. Each group has the same properties.
  • Several trends such as ionisation energy and atomic diameter can be seen across the periods of the periodic table
  • 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.

End of chapter exercises

  1. For the following questions state whether they are true or false. If they are false, correct the statement.
    1. The group 1 elements are sometimes known as the alkali earth metals.
    2. The group 2 elements tend to lose 2 electrons to form cations.
    3. The group 8 elements are known as the noble gases.
    4. Group 7 elements are very unreactive.
    5. The transition elements are found between groups 3 and 4.
    Click here for the solution
  2. Give one word or term for each of the following:
    1. A positive ion
    2. The energy that is needed to remove one electron from an atom
    3. A horizontal row on the periodic table
    4. A very reactive group of elements that is missing just one electron from their outer shells.
    Click here for the solution
  3. For each of the following elements give the ion that will be formed:
    1. sodium
    2. bromine
    3. magnesium
    4. oxygen
    Click here for the solution
  4. The following table shows the first ionisation energies for the elements of period 1 and 2.
    Table 3
    PeriodElementFirst ionisation energy (kJ.mol-1kJ.mol-1)
    1HH1312
     HeHe2372
     LiLi520
     BeBe899
     BB801
     CC1086
    2NN1402
     OO1314
     FF1681
     NeNe2081
    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.
    Click here for the solution

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