Periodic trends: Ionization energy

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  1. Looking at Periodic Trends
Topic Notes
In this lesson, we will learn:
  • The definition of ionization energy and understand its significance to studying the elements.
  • To explain the trend in ionization energy by applying principles of electrostatic forces.
  • The anomalies in the ionization energy data to help develop understanding of electron shells.


  • As seen in Periodic trends: Atomic radius, chemists have found, through experimenting, some principles of electrostatic forces โ€“ forces that exist because charged particles attract or repel each other. The principles are:
    • #1: Oppositely charged particles attract each other, while particles of like charge repel each other.
    • #2: The greater the charge difference of two particles, the greater their force of attraction (for example, the attractive force between a 2+ ion and a 2- ion is stronger than the attractive force between a 1+ ion and a 1- ion).
    • #3: Attractive forces between oppositely charge particles decrease with distance.
    • #4: Repulsive forces between like charged particles decrease with distance.

  • Together the principles form a theory that explains what chemists see in the data of their experiments, such as the atomic radius of chemical elements and their 1st ionization energies.

  • As seen in this chapter so far, arranging the elements by their proton number shows a number of trends in the properties of the elements. This is true going down the table or โ€œgoing down the groupโ€, and going across the table or โ€œacross the periodโ€. The fact that these patterns repeat themselves โ€“ they are periodic โ€“ is why the table of elements is called the periodic table of elements!

  • Ionization energy is defined as the energy required to remove one mole of electrons from one mole of gaseous atoms to form a positive ion.

  • More specifically: The first ionization energy is the energy required to remove one mole of the most weakly-held electrons from one mole of gaseous atoms to form one mole of gaseous ions with a single positive charge.

  • The successive ionization energies follow from the first: it is the energy required to remove one mole of the next most weakly-held electrons from one mole of gaseous ions to form gaseous ions with a one-greater positive charge. For example, the second ionization energy would be the energy required to remove one mole of the most weakly-held electrons from one mole of 1+ charged gaseous ions, forming one mole of 2+ charged gaseous ions.

  • The 1st ionization energies of the elements show a very distinct pattern in the periodic table. For chemists, it is very revealing to study the ionization energies in elements across a period because it shows how difficult it is to remove one extra electron from the same outer electron shell!

  • As briefly talked about in Structure of the Periodic Table, the distinct shape of the periodic table, where the s, p, d, and f blocks exist, is because of ionization energies.

  • The trend in ionization energy across a period (for example, period 2) is explained using electrostatic forces:
    • As you go across the period from left to right, each element contains one extra proton in the nucleus, increasing its charge.
    • Each further element also has one extra electron in its outer shell. This greater charge difference between the positively charged nucleus and negative outer shell electrons results in greater force of attraction (see principle #2) and the electrons being attracted (principle #1) more strongly.
    • This means extra energy is required to be put in to overcome the force of attraction and remove an outer shell electron - in general then, moving to the right of a period, first ionization energy increases.
    • There is an anomaly in this trend for boron: boron's outer shell electron configuration is 2s2 2p1 - it has one electron in the 2p subshell, which is being shielded from the nucleus by the 2s subshell that is already full, causing repulsion (principle #1), while the 2p orbital is further away from the positive nucleus to begin with so is less strongly attracted to it (principle #4). This effect overcomes the greater charge difference from an extra electron and proton. It therefore costs less energy to remove the first electron from boron's outer shell than the general trend would suggest.
    • There is another anomaly in this trend for oxygen: oxygen's outer shell electron configuration is 2s2 2p4, where one of the 2p orbitals is now full with two electrons paired for the first time (until oxygen, the electrons fill up one p orbital by themselves, see Hund's rule). This increases repulsion (see principle #1) and overrides the effect of greater charge difference attracting the electrons to the more highly charged nucleus. This means less energy is required to remove the first electron from oxygen's outer shell than the general trend suggests.

  • The ionization energy trend occurs in the 3rd period too. After each noble gas, there is a massive drop in ionization energy e.g. from Ne to Na. These data helped developed understanding of electron shells and subshells and the number of electrons they can hold:
    • Using electrostatic principles (more negative electrons being attracted to a more positively charged nucleus), we would expect greater attraction of the electrons by the nucleus, and even more energy needed to remove (one mole of) electrons. So what do the repeating โ€“ or periodic - drop in ionization energy mean?
    • Our current theory says that the extra electron in boron must be in a different โ€˜stateโ€™ or sub shell than the last electron in beryllium. Why else would it cost a lot less energy than beryllium to remove an electron?
    • The idea of electrons being in shells and subshells was developed by quantum mechanics, which also established the number of electrons the subshells could hold.

  • The trend in ionization energies practically shows you how easily an atom can form a positive ion โ€“ by losing an electron, a positive ion is formed. Based on this, we can observe that it is easier for metals to lose electrons and form positive ions than non-metals. This is related to the electronegativity of an atom โ€“ the focus of the next lesson!