The reactivity of alkali metals increases as we go down the group. Lithium is the most stable while Caesium is one of the most reactive alkali metals.
The reactivity of alkali metals and other electropositive elements can be explained in terms of their ionisation energy.
Ionisation energy
- Ionisation energy is defined as the amount of energy required to remove the most loosely bound valence electron from the isolated gaseous atom in its lowest energy state or ground state, in order to convert it into a gaseous cation.
- The process of converting an atom into an ion is called ionisation.
- Ionisation energy is also known as ionisation potential or ionisation enthalpy.
- If the valence electron is to be removed, then, work must be done to overcome this electrostatic force. That’s why energy is required to remove the valence electrons. That’s why ionisation processes are always endothermic. That means the ionisation energy will always have a positive value.
Successive ionisation energies
- All electrons can be removed from an atom until only the nucleus of an atom is left.
- If an electron is removed from a neutral atom, it results in the formation of a monovalent cation. The energy required for this is called the first ionisation energy.
- If the second electron is removed from the monovalent cation forming a divalent cation, the energy required is called the second ionisation energy.
- The third, fourth and ultimately every following electron can be removed in a similar way, one at a time. This sequence of ionisation energies is called successive ionisation energies.
- For every element, the successive ionisation energy keeps increasing and follows this order.
  $IE_1\ \ \ \ <\ \ \ \ IE_2\ \ \ \ <\ \ \ \ IE_3$ and so on…
- The reason why successive ionisation energies increase is because the charge on the ion gets larger as each electron is removed, which increases the effective nuclear charge. There is a greater attractive force between the positively charged protons in the nucleus and the remaining negatively charged electrons. More energy is required to overcome these increased attractive forces.
  - The effective nuclear charge is the net charge that a valence electron experiences in an atom with multiple electrons.
- The successive ionisation energies of carbon can be seen in the table here:
  
  There is a big difference between the fourth and the fifth ionisation energies of carbon because, for the fifth ionisation energy of carbon, the next electron has to be removed from the next inner shell which is closer to the nucleus. After the fourth ionisation energy, all four valence electrons of carbon in its second shell are removed. So, the fifth electron has to be removed from the first shell. The electrons in the first shell are quite strongly attracted by the nucleus and hence very high energy is required to remove these electrons.
Ionisation energy mainly depends upon the following factors:
- Atomic radius
  The atomic radius is the distance between the nucleus of an atom and its valence electron. The electrostatic force of attraction between the charged particles decreases with the increase in the distance between them. So it becomes easier to remove the electron from the atom. Hence, there is an inverse relationship between the atomic radius and the ionisation energy. This is the same in the case of ionic radius as well.
- Nuclear charge
  The bigger the nuclear charge, the greater the attractive force between the nucleus andthe electrons. So, more energy is required to overcome this force to remove an electron. Hence, ionisation energy increases with the increase in the nuclear charge.
- Shielding effect by the electrons
  In a multi-electron atom, the electrons of the valence shell are attracted to the nucleus. Since all electrons are negatively charged, these valence electrons are repelled by the electrons present in the inner shells. On account of this, the actual force of attraction between the nucleus and the valence electrons is somewhat decreased by the repulsive forces acting in opposite directions. This decrease in the force of attraction is called the shielding effect or screening effect.
  
  The greater the shielding effect, the easier it becomes to remove the valence electron and, consequently, the lower the ionisation energy.
The periodic trend of ionisation energies
- While going down the group, even though there’s a noticeable increase in nuclear charge, there’s also an increase in the distance from the nucleus and an increase in the shielding effect. The combined effect of the increase in the atomic radius and the shielding effect outweighs the increase in the nuclear charge. Hence, the ionisation energy decreases down the group.
- Moving from left to right in a period, the atomic number, as well as the magnitude of nuclear charge, increases. The increased effective nuclear charge attracts the electron cloud of the atom nearer to the nucleus, which causes a decrease in the atomic radius. The ionisation energy increases across a period because there is an increase in the nuclear charge and a decrease in atomic radius.
- Halogens have the second-highest ionisation energy in a given period after noble gases.
- Abnormalities in the periodic trend
  - Going across a period, there are two ways in which the ionisation energy may be affected by the electron configuration.
    - In period $2$, there is an anomaly when moving from the $2s$ subshell to the $2p$ subshell. Beryllium has its outermost electrons in the $2s$ subshell, while the boron atom has its outermost electron in the $2p$ subshell, which is slightly further away from the nucleus as compared to the $2s$ subshell. Since $p$ electrons have slightly more energy than $s$ electrons, it takes a little less energy to remove than the electron from boron.
      
      One other factor is that the electrons in the $2s$ subshell shield the electrons in the $2p$ subshell. These factors allow the first ionisation energy for boron to be less than that for beryllium, even though the boron atom has a comparatively higher nuclear charge.
      In the same way, magnesium has a slightly higher first ionisation energy than aluminium.
    - Nitrogen has one electron occupying each of the three $p$ orbitals in the second sub-level. On the other hand, oxygen has an additional electron in one of the three $2p$ orbitals. The presence of paired electrons in an orbital leads to a greater electron-to-electron repulsion being experienced by these $2p$ electrons. The extra repulsion between the pair of electrons in that orbital results in a lower amount of energy being needed to remove one of these electrons. Hence, oxygen has slightly lower ionisation energy due to this spin-pair repulsion.
    - The first ionisation energy of some elements is shown in the following table.
Position of an element from successive ionisation energies

There is a large difference between the fifth and the sixth ionisation energies. This confirms that the element has five valence electrons, which are comparatively easier to remove than the sixth electron, which is to be removed from the inner shell and closer to the nucleus. Hence, the element is a group $5$ element.

1.05 - Ionisation Energy

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