Ionization Energy:
Ionization energy is the energy required to remove one electron from a neutral atoms of an element. It has a trend for periods and groups. This is known as the first ionization energy. Ionization energy is measured in kJ/mol The second ionization energy is the energy required to remove an electron from an Aᐩ¹. The third ionization energy is the energy required to remove an electron from an Aᐩ².
A + energy → (A^+) + electron
The ionization trend for periods is from left to right, the ionization energy increases.The ionization energy increases for one main reason. This reason being the atomic radius and the nuclear charge. This is not two reasons because the two coincide with each other. When the nuclear charge is greater, the atomic radius is smaller. The atomic radius hold more nuclear charge or more protons thus causing a greater attraction or pull between the protons and the electrons. This greater pull or attraction requires more energy to remove an electron from an atom. The trend for ionization energy in groups is the the ionization energy decreases from from top to bottom, meaning that the highest ionization energy is at the top and the lowest ionization energy is at the bottom. The reason for this is shielding and atomic radius. The atomic radius is smaller at the top of the periodic table. There are less electrons thus there is no need for more energy levels. The lack of energy levels causes greater attraction between the nucleus and the electrons causing the need for more ionization energy to remove and electron. The second reason that the ionization energy decreases down a group is the shielding effect. The shielding effect is the blockage of of electrons attraction of those in the outer energy level with the nucleus. There are more electrons the farther down the periodic table thus a larger shielding effect occurs causing for less ionization energy to remove an electron. After the first ionization energy, the trend continues but in a less pronounced way when used in the second or third ionization energy. After the first ionization energy it becomes significantly harder for electrons to be removed because taking away electrons will result in closer attraction between the nucleus and the electrons as well as a smaller shielding effect because there is not enough electrons to block attraction between one another. Once elements reach a noble gas configuration it becomes virtually impossible to remove an electron from the stable state.
Ionization energy is the energy required to remove one electron from a neutral atoms of an element. It has a trend for periods and groups. This is known as the first ionization energy. Ionization energy is measured in kJ/mol The second ionization energy is the energy required to remove an electron from an Aᐩ¹. The third ionization energy is the energy required to remove an electron from an Aᐩ².
A + energy → (A^+) + electron
The ionization trend for periods is from left to right, the ionization energy increases.The ionization energy increases for one main reason. This reason being the atomic radius and the nuclear charge. This is not two reasons because the two coincide with each other. When the nuclear charge is greater, the atomic radius is smaller. The atomic radius hold more nuclear charge or more protons thus causing a greater attraction or pull between the protons and the electrons. This greater pull or attraction requires more energy to remove an electron from an atom. The trend for ionization energy in groups is the the ionization energy decreases from from top to bottom, meaning that the highest ionization energy is at the top and the lowest ionization energy is at the bottom. The reason for this is shielding and atomic radius. The atomic radius is smaller at the top of the periodic table. There are less electrons thus there is no need for more energy levels. The lack of energy levels causes greater attraction between the nucleus and the electrons causing the need for more ionization energy to remove and electron. The second reason that the ionization energy decreases down a group is the shielding effect. The shielding effect is the blockage of of electrons attraction of those in the outer energy level with the nucleus. There are more electrons the farther down the periodic table thus a larger shielding effect occurs causing for less ionization energy to remove an electron. After the first ionization energy, the trend continues but in a less pronounced way when used in the second or third ionization energy. After the first ionization energy it becomes significantly harder for electrons to be removed because taking away electrons will result in closer attraction between the nucleus and the electrons as well as a smaller shielding effect because there is not enough electrons to block attraction between one another. Once elements reach a noble gas configuration it becomes virtually impossible to remove an electron from the stable state.