The first ionization energy is a fundamental property of an element that measures the energy required to remove its most loosely held electron. This measurement provides a direct insight into the strength of the attractive force between the nucleus and the valence electrons. Understanding this energy is essential because it helps predict an atom’s tendency to form positive ions, which is a significant factor in determining its chemical reactivity and bonding behavior.
The Process and Energy Measurement
The definition of the first ionization energy (IE1) specifies the minimum energy needed to convert one mole of neutral atoms into one mole of gaseous ions, each with a single positive charge. The atom must be isolated and in its gaseous state to ensure the energy measured is solely for overcoming the nuclear attraction on the electron. This process is represented by the standard chemical equation: X(g) + Energy yields X\(^{+}\)(g) + e\(^{-}\).
Because energy is required to remove an electron, ionization is an endothermic process, meaning the IE1 value is positive. The standard unit used by chemists is kilojoules per mole (kJ/mol), which quantifies the energy for a large collection of atoms. Values generally range from the low end, such as 381 kJ/mol for Cesium, to the high end, like 2370 kJ/mol for Helium.
Fundamental Factors Influencing the Value
The magnitude of the first ionization energy is determined by the balance of forces within the atom, primarily the electrostatic attraction between the positive nucleus and the negative outer electron. Three main factors govern this attraction. The first is the nuclear charge, which is the number of protons in the nucleus. As the number of protons increases, the positive charge attracting the electrons increases, pulling them in more tightly and requiring more energy to remove the outermost one.
The second major influence is electron shielding, also known as the screening effect. Inner-shell electrons lie between the nucleus and the outermost valence electron, effectively reducing the net positive charge that the valence electron “feels.” The more inner electron shells an atom has, the greater the shielding, which weakens the hold on the valence electron and lowers the ionization energy.
The third factor is the distance from the nucleus, which is directly related to the atomic radius. Since the electrostatic force weakens dramatically with distance, an electron located in a higher principal energy level is farther from the nucleus and is held less strongly. Therefore, larger atoms, which have greater atomic radii, generally have lower ionization energies because their outermost electrons are easier to remove.
Predicting Changes Using the Periodic Table
The factors influencing ionization energy create predictable patterns across the periodic table. Moving from left to right across a period, the first ionization energy generally increases. This occurs because the nuclear charge steadily increases while the outermost electrons are added to the same main energy shell, meaning the shielding effect remains relatively constant. The increasing attraction from the nucleus holds the electrons more tightly, making them harder to remove.
Conversely, moving down a group, the first ionization energy generally decreases. In this case, each successive element adds a new electron shell, significantly increasing the atomic radius and the distance of the valence electron from the nucleus. The effect of the increased distance and the greater number of inner-shell electrons providing shielding outweighs the slight increase in nuclear charge.
There are minor irregularities to these trends, such as the slight decrease in ionization energy observed when moving from Group 2 to Group 13, or from Group 15 to Group 16. These deviations are explained by the stability associated with removing an electron from a fully-filled subshell or the electron-electron repulsion that occurs when electrons are first paired in an orbital. The overall trend is a strong indicator of an element’s metallic or non-metallic character.
Why Successive Ionizations Are Different
The concept of first ionization energy is distinct because removing a second, third, or subsequent electron requires a different amount of energy, collectively known as successive ionization energies (IE2, IE3, etc.). The second ionization energy (IE2) is greater than the first (IE1), and subsequent ionizations continue to increase. This consistent increase occurs for two main reasons related to the atom’s charge and structure.
First, once the first electron is removed, the remaining electron is being pulled away from a positively charged ion, X\(^{+}\)(g), rather than a neutral atom, X(g). The positive charge remaining on the ion exerts a stronger attractive force on the remaining electrons, making it harder to remove the next one. Second, a particularly large jump in successive ionization energy occurs when the electron is removed from a stable, filled inner electron shell, which is much closer to the nucleus. This sharp increase explains why elements like sodium form an Na\(^{+}\) ion but not an Na\(^{2+}\) ion.