Ionization energy is the energy required to remove the most loosely held electron from a neutral atom. The energy required to remove a second electron is called the second ionization energy.
Both trends are increasing for the most part. There are some exceptions like boron and oxygen in period 2 and aluminum and sulfur in period 3. All of the elements before these ones are stable and it is easier to remove the electron from them than from an electron of boron, oxygen, aluminum, and sulfur. This makes boron, oxygen, aluminum, and sulfur have a lower ionization energy.
First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels increases downward within each group.
Ionization energy decreases moving down a group. This happens because when you are moving down a family, the atomic radius is increasing. For the most part, the larger the atomic radius, the lower the ionization energy.
Electronegativity is the attraction of an atom for electrons in the covalent bond.
On graph paper
The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei don’t exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. There is an exception. Electronegativity increases to the right. This does not include noble gasses. This is because the definition of electronegativity...
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