Ionization Energy Trends in the Periodic Table

The ionization energy of an atom is the amount of power required to rerelocate an electron from the gaseous develop of that atom or ion.

You are watching: Why is the second ionization energy greater than the first ionization energy?

first ionization energy - The energy compelled to rerelocate the greatest energy electron from a neutral gaseous atom.

For Example:

Na(g) → Na+(g)+ e-I1 = 496 kJ/mol

Notice that the ionization power is positive. This is bereason it requires power to remove an electron.

second ionization power - The power compelled to remove a second electron from a singly charged gaseous cation.

For Example:

Na+(g)→ Na2+(g)+ e-I2 = 4560 kJ/mol

The second ionization energy is practically ten times that of the initially bereason the number of electrons leading to repulsions is lessened.

3rd ionization power - The power required to remove a third electron from a doubly charged gaseous cation.

For Example:

Na2+(g)→ Na3+(g)+ e-I3 = 6913 kJ/mol

The third ionization power is even better than the second.


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Successive ionization energies increase in magnitude bereason the variety of electrons, which reason repulsion, steadily decrease. This is not a smooth curve There is a huge jump in ionization power after the atom has actually lost its valence electrons. An atom that has actually the same electronic configuration as a noble gas is really going to host on to its electrons. So, the amount of energy essential to rerelocate electrons beyond the valence electrons is considerably higher than the energy of chemical reactions and bonding. Therefore, only the valence electrons (i.e., electrons external of the noble gas core) are affiliated in chemical reactions.

The ionization energies of a specific atom depend on the average electron distance from the nucleus and also the efficient nuclear charge

These factors have the right to be shown by the following trends:


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first ionization energy decreases down a team.

This is because the highest energy electrons are, on average, farther from the nucleus. As the major quantum number rises, the size of the orbital increases and also the electron is simpler to remove.

Examples:

I1(Na) > I1(Cs)

I1(Cl) > I1(I)

first ionization power boosts across a period.

This is because electrons in the very same primary quantum shell carry out not totally shield the increasing nuclear charge of the proloads. Hence, electrons are organized more tightly and also need more power to be ionized.

Examples:

I1(Cl) > I1(Na)

I1(S) > I1(Mg)

The graph of ionization energy versus atomic number is not a perfect line bereason tbelow are exceptions to the rules that are conveniently explained.

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Filled and also half-filled subshells show a small rise in stcapability in the very same method that filled shells display increased stcapacity. So, once trying to remove an electron from one of these filled or half-filled subshells, a slightly better ionization energy is found.

Example 1:

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I1(Be) > I1(B)

It"s harder to ionize an electron from beryllium than boron because beryllium has a filled "s" subshell.

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Example 2:

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I1(N) > I1(O)

Nitrogen has actually a half-filled "2p" subshell so it is harder to ionize an electron from nitrogen than oxygen.