Trends in Ionization Energy
What is Ionization Energy?
Ionization energy is defined as the amount of energy required to remove the most loosely bound electron
from a gaseous atom or ion. This process is always endothermic, meaning it requires energy input, and is
therefore represented by a positive value.
First Ionization Energy
The first ionization energy is the energy required to remove one electron from each atom in one mole of
gaseous atoms to form one mole of gaseous positive ions with a +1 charge. The general equation is:
X(g) → X+ (g) + e−
Successive Ionization Energies
Successive ionization energies refer to the energy required to remove subsequent electrons from an atom
or ion. Each successive ionization energy is greater than the previous one because as electrons are
removed, the remaining electrons are held more tightly by the nucleus due to increased effective nuclear
charge.
Second Ionization Energy: The energy required to remove one electron from each gaseous +1 ion to
form gaseous +2 ions.
X+ (g) → X2+ (g) + e−
Third Ionization Energy: The energy required to remove one electron from each gaseous +2 ion to
form gaseous +3 ions.
X2+ (g) → X3+ (g) + e−
And so on for higher ionization energies.
Factors Affecting Ionization Energy
Several factors influence the magnitude of ionization energy:
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, Atomic Radius: A larger atomic radius means the outermost electron is further from the nucleus,
experiencing weaker attractive forces. This results in a lower ionization energy as less energy is
required to remove the electron. Conversely, a smaller atomic radius leads to stronger attraction and
higher ionization energy.
Nuclear Charge: A higher nuclear charge (more protons) results in a stronger electrostatic attraction
between the nucleus and the electrons. This increased attraction means more energy is needed to
remove an electron, leading to a higher ionization energy.
Shielding Effect: Electrons in inner energy levels shield the outermost electrons from the full
attractive force of the nucleus. The more inner electron shells there are between the nucleus and the
outermost electron, the greater the shielding effect. Increased shielding reduces the attraction
between the nucleus and the outermost electron, making it easier to remove and thus lowering the
ionization energy.
Electron Configuration and Subshell Stability: Electrons in filled or half-filled subshells are more
stable and require more energy to remove.
Filled subshells (e.g., s2 , p6 ) are very stable.
Half-filled subshells (e.g., s1 , p3 , d5 ) are also relatively stable.
Electrons in paired orbitals within a subshell experience electron-electron repulsion, making
them slightly easier to remove compared to electrons in a half-filled or filled subshell.
Trends in the Periodic Table
Across a Period (Left to Right)
General Trend: Ionization energy generally increases across a period.
Explanation: As you move from left to right across a period, the number of protons in the nucleus
increases, leading to a stronger nuclear charge. Electrons are added to the same principal energy
level, so the shielding effect remains relatively constant. The atomic radius decreases because the
increasing nuclear charge pulls the electrons more strongly. This stronger attraction requires more
energy to remove the outermost electron.
Anomalies:
Group 2 to Group 13 (e.g., Be to B): Ionization energy decreases. This is because beryllium has
a filled 2s2 subshell, which is stable. Boron's outermost electron is in the 2p subshell, which is
at a slightly higher energy level and further from the nucleus, making it easier to remove
despite the increased nuclear charge.
Group 15 to Group 16 (e.g., N to O): Ionization energy decreases. Nitrogen has a stable half-
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filled 2p subshell. Oxygen has a paired electron in its 2p configuration. The repulsion between
the paired electrons in oxygen makes it easier to remove one electron compared to removing
an electron from the stable half-filled p subshell of nitrogen.
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What is Ionization Energy?
Ionization energy is defined as the amount of energy required to remove the most loosely bound electron
from a gaseous atom or ion. This process is always endothermic, meaning it requires energy input, and is
therefore represented by a positive value.
First Ionization Energy
The first ionization energy is the energy required to remove one electron from each atom in one mole of
gaseous atoms to form one mole of gaseous positive ions with a +1 charge. The general equation is:
X(g) → X+ (g) + e−
Successive Ionization Energies
Successive ionization energies refer to the energy required to remove subsequent electrons from an atom
or ion. Each successive ionization energy is greater than the previous one because as electrons are
removed, the remaining electrons are held more tightly by the nucleus due to increased effective nuclear
charge.
Second Ionization Energy: The energy required to remove one electron from each gaseous +1 ion to
form gaseous +2 ions.
X+ (g) → X2+ (g) + e−
Third Ionization Energy: The energy required to remove one electron from each gaseous +2 ion to
form gaseous +3 ions.
X2+ (g) → X3+ (g) + e−
And so on for higher ionization energies.
Factors Affecting Ionization Energy
Several factors influence the magnitude of ionization energy:
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, Atomic Radius: A larger atomic radius means the outermost electron is further from the nucleus,
experiencing weaker attractive forces. This results in a lower ionization energy as less energy is
required to remove the electron. Conversely, a smaller atomic radius leads to stronger attraction and
higher ionization energy.
Nuclear Charge: A higher nuclear charge (more protons) results in a stronger electrostatic attraction
between the nucleus and the electrons. This increased attraction means more energy is needed to
remove an electron, leading to a higher ionization energy.
Shielding Effect: Electrons in inner energy levels shield the outermost electrons from the full
attractive force of the nucleus. The more inner electron shells there are between the nucleus and the
outermost electron, the greater the shielding effect. Increased shielding reduces the attraction
between the nucleus and the outermost electron, making it easier to remove and thus lowering the
ionization energy.
Electron Configuration and Subshell Stability: Electrons in filled or half-filled subshells are more
stable and require more energy to remove.
Filled subshells (e.g., s2 , p6 ) are very stable.
Half-filled subshells (e.g., s1 , p3 , d5 ) are also relatively stable.
Electrons in paired orbitals within a subshell experience electron-electron repulsion, making
them slightly easier to remove compared to electrons in a half-filled or filled subshell.
Trends in the Periodic Table
Across a Period (Left to Right)
General Trend: Ionization energy generally increases across a period.
Explanation: As you move from left to right across a period, the number of protons in the nucleus
increases, leading to a stronger nuclear charge. Electrons are added to the same principal energy
level, so the shielding effect remains relatively constant. The atomic radius decreases because the
increasing nuclear charge pulls the electrons more strongly. This stronger attraction requires more
energy to remove the outermost electron.
Anomalies:
Group 2 to Group 13 (e.g., Be to B): Ionization energy decreases. This is because beryllium has
a filled 2s2 subshell, which is stable. Boron's outermost electron is in the 2p subshell, which is
at a slightly higher energy level and further from the nucleus, making it easier to remove
despite the increased nuclear charge.
Group 15 to Group 16 (e.g., N to O): Ionization energy decreases. Nitrogen has a stable half-
3 4
filled 2p subshell. Oxygen has a paired electron in its 2p configuration. The repulsion between
the paired electrons in oxygen makes it easier to remove one electron compared to removing
an electron from the stable half-filled p subshell of nitrogen.
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