8.2
Ionization Energy
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…
The chemical behavior of atoms and ions is greatly affected by how easy or difficult it is to remove their electrons, especially the outermost electrons which participate in chemical bond formations.
The energy required to remove an electron from a gaseous atom in its ground state is called the first ionization energy and is given in kJ/mol. 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.
But there are some exceptions to consider.
Boron has a smaller ionization energy than beryllium, even though it is farther to the right on the periodic table. Beryllium has lower energy 2s electrons, whereas boron has a higher energy 2p 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.
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Q1: What is first ionization energy and why does it matter in chemistry?
First ionization energy is the energy required to remove the outermost electron from a gaseous atom in its ground state, measured in kJ/mol. It directly affects chemical behavior because these valence electrons participate in chemical bond formation. Understanding ionization energy helps predict how readily atoms lose electrons and form cations during chemical reactions.
Q2: Why does ionization energy decrease moving down a group on the periodic table?
Ionization energy decreases down a group because atomic size increases, placing valence electrons farther from the nucleus. The increased distance weakens the electrostatic attraction between the nucleus and outermost electrons, making them easier to remove. This trend is fundamental to understanding periodic properties and element reactivity.
Q3: How does ionization energy change across a period and what causes this trend?
Ionization energy generally increases across a period because atomic number increases, raising the effective nuclear charge experienced by valence electrons. Stronger nuclear attraction makes electron removal more difficult. This trend explains why chlorine has higher ionization energy than sodium, demonstrating the relationship between nuclear charge and electron removal difficulty.
Q4: Why does boron have lower ionization energy than beryllium despite having greater nuclear charge?
Boron has lower ionization energy than beryllium because it loses a 2p electron, which is higher in energy than beryllium's 2s electron. The 2p electron is easier to remove due to increased orbital energy from penetration and shielding effects. This exception demonstrates that subshell type, not just nuclear charge, determines ionization energy.
Q5: What explains the exception where oxygen has lower first ionization energy than nitrogen?
Oxygen has lower first ionization energy than nitrogen because removing one electron eliminates electron-electron repulsion from the paired electrons in the 2p orbital, creating a more stable half-filled configuration. This energetic advantage outweighs the effect of increased nuclear charge, making electron removal easier for oxygen despite its position farther right on the periodic table.
Q6: Why do successive ionization energies increase dramatically for an element?
Successive ionization energies increase because each removal requires more energy as electrons are stripped away. The largest jump occurs when removing core electrons from a cation with a noble gas configuration, which experience much stronger nuclear attraction. For example, potassium's second ionization energy is significantly higher because it removes a core electron from a stable noble gas configuration.
Q7: How does removing an electron from a cation compare to removing one from a neutral atom?
Removing an electron from a cation is more difficult than from a neutral atom because the cation has greater electrostatic attraction to its remaining electrons. Additionally, removing electrons from cations with higher positive charges requires even more energy. This principle explains why successive ionization energies always increase for any element.