[Explained] Why Does Ionization Energy Decrease Down A Group?

You might have wondered, why does ionization energy decrease as you move down a group in the periodic table? 

It is mostly due to the increase in the size of the atoms.

But first let us brush up our knowledge about ionization energy and the factors that affect it.

Ionization Energy of an atom is the amount of energy required to remove the most loosely bound electron from an from an isolated gaseous atom to form a gaseous atom.

It is affected by mainly five factors:

1. Size of the atom
2. Nuclear Charge
3. Screening/Shielding effect
4. Penetration effect of electrons
5. Electron configuration

Let us understand which best explains why ionization energy of elements decrease going from top to bottom in a group.

 As we all know that the value of the ionization energy decreases as regularly with increase in atomic number in going down the group.

The regular decrease in going down the group is mainly due to the following two factors.

1. The size of the atoms increase with increase in atomic number in going down the group. The number of shells also increase as a matter of fact.

Let us take the example of the second group of elements in the periodic table. They are popularly known as the alkaline earth metals. This is because they form basic solutions with water. Below is the list of all the elements in the group 2 of the periodic table.

1. Beryllium
2. Magnesium
3. Calcium
4. Strontium
5. Barium
6. Radium

The number of valence electrons in them does not change. They all have 2 valence electrons in their highest energy orbitals (ns2).

When we move from the left to the right on the periodic table the size of atoms gradually keeps decreasing. Therefore they are smaller than the alkali metals that is the group 1 of the periodic table.

 When we move down the group the number of energy levels increases which causes the valence electrons to be present farther than the valence electron of the element above it on the same group of the periodic table.

 

Element	Size in pm (picometers)
Beryllium	105
Magnesium	150
Calcium	180
Strontium	200
Barium	215
Radium	215

We can see this is true by studying the empirical values of the size of the atoms of the respective second group elements. The empirical atomic size of beryllium is 105 pm, Magnesium is 150 pm, Calcium is 180 pm, Strontium is 200 pm, Barium and Radium are both 215 pm. 

As their size increases the distance between the nucleus and the outermost electrons also increase.

What this increase in the distance between nucleus and valence electron does is that the energy required to remove the outermost electron is decreased.

To put it in simple terms, as the distance between the nucleus and the valence electron increases the energy required which is defined as the ionization energy keeps decreasing. 

Next we shall see and learn about the changes in the effective nuclear charge.

Nuclear charge also is reduced as we move from top to bottom on a periodic table.

To make things a bit more clearer we shall compare the values of effective nuclear charge of elements of the second group. 

Let us just brush up the definition of effective nuclear charge and see what it means.

According to Wikipedia's Effective Nuclear Charge page, 

The effective nuclear charge of an atom is the number of protons that an electron in the element effectively sees due to screening by inner shell electrons. It is a measure of the electrostatic interaction between the negatively charged electrons and the positively charged protons in the atom.

Element	Nuclear Charge
Beryllium	1.912
Magnesium	3.308
Calcium	4.398
Strontium	6.071
Barium	NA
Radium	NA

Now you might say that the statement above said that the effective nuclear charge decreases down the group. 

While we are at it, let me remove your confusion about nuclear charge and effective nuclear charge. 

Nuclear charge is the electric charge of a nucleus of an atom, equal to the number of protons in the nucleus times the elementary charge. 

Whereas, the effective nuclear charge is nothing but the attractive positive charge of the nuclear protons acting on valence electrons which is always less than the total number of protons present in a nucleus due to the shielding effect.

Now here another question arises. What is Shielding effect?

In a multi-electron atom, the electrons present between the nucleus and the valence shell electron ( the inner electrons) shield the valence electron present from the nucleus. 

This decreases the effect of the nucleus on the valence electron, due to the presence of inner electrons, the nucleus exerts less force of attraction on the valence electron.

The values of nuclear charge increases but the shielding effect more than compensates for the increase, therefore the effective nuclear is lower for elements that are lower in the same group of the periodic table.

We can compare this statement to the data from Wikipedia's molar ionization energies of the elements page. Here we have taken the first group of the periodic table the alkali metals.

Ionization Energy Values of Alkali Metals Group 1
Element	Ionization Energy Values in kJ/mol
Hydrogen	1312
Lithium	520
Sodium	495.8
Potassium	418
Rubidium	403
Cesium	375.7
Francium	380

We can see that there is a gradual decrease in the ionization energy of the alkali metals when moving down the group from Lithium to Cesium.

This is due to the increase in the size of the atoms. As we go down the group there is an addition of atomic orbitals. More atomic orbitals means that there will be more shielding effect which will eventually make it easier to remove the outermost electron.

Electron Configuration of Alkali Metals
Element	Electron Configuration
Hydrogen	1s1
Lithium	[He] 2s1
Sodium	[Ne] 3s1
Potassium	[Ar] 4s1
Rubidium	[Kr] 5s1
Cesium	[Xe] 6s1
Francium	[Rn] 7s1

Trend of Ionization Energy In the Periodic Table

2. Size of the atoms increase due to increase in the number of shells, the shielding effect increases and the valence electron get more and more loosely bound to the nucleus as atomic number increases in going down a group. This best explains why ionization energy tends to decrease from the top to the bottom of a group.

Ionization Energy In-Depth

The decrease is very sharp for first group of elements as we can see in the graph below.

Trend of Ionization Energy in alkali metals

Factors affecting Ionization Energy