# 4.1 Ionisation energy and trends  (Page 2/3)

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## 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)

 1. A positive ion that has 3 less electrons than its neutral atom A. ${\mathrm{Mg}}^{2+}$ 2. An ion that has 1 more electron than its neutral atom B. ${\mathrm{Cl}}^{-}$ 3. The anion that is formed when bromine gains an electron C. ${\mathrm{CO}}_{3}^{2-}$ 4. The cation that is formed from a magnesium atom D. ${\mathrm{Al}}^{3+}$ 5. An example of a compound ion E. ${\mathrm{Br}}^{2-}$ 6. A positive ion with the electron configuration of argon F. ${\mathrm{K}}^{+}$ 7. A negative ion with the electron configuration of neon G. ${\mathrm{Mg}}^{+}$ H. ${\mathrm{O}}^{2-}$ I. ${\mathrm{Br}}^{-}$

## 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.

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.

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

 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 $Z=7$

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.
• 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.
• Transition metals: The differences between groups in the transition metals are not usually dramatic.

what is the stm
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