Previous Video 8. The energy required to remove the next electron is called the second ionization energy, and so on. Moving down a column, the ionization energies decrease.
Recall that the highest principal quantum number of valence electrons increases down the column leading to larger atomic sizes. Thus, the farther the outermost electrons are, the easier they are to remove. For main-group elements, the ionization energy increases across the period. The reason lies in the increasing atomic number, where valence electrons experience a higher effective nuclear charge making the removal of outermost electrons more difficult. This explains why chlorine has a higher ionization energy than sodium, for example.
Generally, ionization energy is a minimum for an alkali metal and rises to a peak with each noble gas. Transition metals display a small increase in the ionization energy across the period, and the f -block elements show an even smaller change.
Boron has a smaller ionization energy than beryllium, even though it is farther to the right on the periodic table. Beryllium has lower energy 2 s electrons, whereas boron has a higher energy 2 p electron making its removal energetically more favorable.
Another exception is oxygen, which has lower first ionization energy than nitrogen. Compared to nitrogen, oxygen has four p -electrons, and removing one electron eliminates the electron-electron repulsion.
Thus, less energy is required for the ionization. These exceptions are observed in succeeding periods too. Electron removal from cations is more difficult than from neutral atoms. Generally, the successive ionization energies increase for elements.
Consider potassium. The second ionization energy is significantly higher, as it involves the removal of a core electron from an ion with a noble gas configuration. Similarly, for calcium, there is a high increase from the second to third ionization energy as a core electron is removed from a cation with a noble gas configuration.
The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy IE 1. The energy required to remove the second most loosely bound electron is called the second ionization energy IE 2. The energy required to remove the third electron is the third ionization energy, and so on. Energy is always required to remove electrons from atoms or ions, so ionization processes are endothermic and IE values are always positive.
For larger atoms, the most loosely bound electron is located farther from the nucleus and so is easier to remove. Thus, as size atomic radius increases, the ionization energy should decrease. Within a period, the IE 1 generally increases with increasing Z. Down a group, the IE 1 value generally decreases with increasing Z. There are some systematic deviations from this trend, however. Note that the ionization energy of boron atomic number 5 is less than that of beryllium atomic number 4 even though the nuclear charge of boron is greater by one proton.
This can be explained because the energy of the subshells increases as l increases, due to penetration and shielding. Within any one shell, the s electrons are lower in energy than the p electrons. This means that an s electron is harder to remove from an atom than a p electron in the same shell.
The electron removed during the ionization of beryllium [He]2 s 2 is an s electron, whereas the electron removed during the ionization of boron [He]2 s 2 2 p 1 is a p electron; this results in lower first ionization energy for boron, even though its nuclear charge is greater by one proton.
Thus, we see a small deviation from the predicted trend occurring each time a new subshell begins. Whether a sheep chooses to stay with the herd or go its own way depends on the balance between attraction to the herd and attraction to the outside influence.
There is an on-going tension between the electrons and protons in an atom. Reactivity of the atom depends in part on how easily the electrons can be removed from the atom. We can measure this quantity and use it to make predictions about the behaviors of atoms. Ionization energy is the energy required to remove an electron from a specific atom.
The ionization energies associated with some elements are described in the Table 1. For any given atom, the outermost valence electrons will have lower ionization energies than the inner-shell kernel electrons. As more electrons are added to a nucleus, the outer electrons become shielded from the nucleus by the inner shell electrons. This is called electron shielding.
If we plot the first ionization energies vs. Moving from left to right across the periodic table, the ionization energy for an atom increases. We can explain this by considering the nuclear charge of the atom. The more protons in the nucleus, the stronger the attraction of the nucleus to electrons. This stronger attraction makes it more difficult to remove electrons.
Within a group, the ionization energy decreases as the size of the atom gets larger. If you put a ruler on the first and second points to establish the trend, you'll find that the third, fourth and fifth points lie above the value you would expect.
That is because the first two electrons are coming from pairs in the 3p levels and are therefore rather easier to remove than if they were unpaired. Again, if you put a ruler on the 3rd, 4th and 5th points to establish their trend, you'll find that the 6th and 7th points lie well above the values you would expect from a continuation of the trend.
That is because the 6th and 7th electrons are coming from the 3s level - slightly closer to the nucleus and slightly less well screened. People sometimes get confused with these graphs because they forget that they are removing electrons from the atom. For example, the first point refers to the first electron being lost - from a 3p orbital. Basically, you start from the outside of the atom and work in towards the middle.
If you start from the 1s orbital and work outwards, you are doomed to failure! To plot any more ionisation energies for chlorine needs a change of vertical scale. The seventeenth ionisation energy of chlorine is nearly , kJ mol -1 , and the vertical scale has to be squashed to accommodate this.
This is now a "log graph" - plotted by finding the logarithm of each ionisation energy press the "log" button on your calculator. This doesn't simply squash the vertical scale. It distorts it as well, to such an extent that the only useful thing the graph now shows is the major jumps where the next electron to be removed comes from an inner level. The distortion is so great in the first 8 ionisation energies, for example, that the patterns shown by the previous graph are completely and misleadingly destroyed.
If this is the first set of questions you have done, please read the introductory page before you start.
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