Ionization energy is the amount of energy required to remove an electron from a gaseous atom or ion.
Ionization energies are endothermic because it requires energy to remove an electron.
The first ionization energy is the amount of energy required to remove the first electron from a neutral gaseous atom.
We can represent this process using the equation shown,
where the electron is a product because it's being removed.
The second ionization energy is the amount of energy required to remove the second electron from an already positive gaseous ion,
as shown in the equation.
From this graph we can see that first ionization energies for alkali metals are the lowest,
meaning it's easiest to remove an electron from an alkali metal.
This makes sense, since alkaline metals tend to give away one electron to have the same electron configuration has a noble gas.
We can also see that the first ionization energies for noble gases are the highest;
because noble gases are already stable, it requires significant amounts of energy to remove an electron.
We can see from this graph that first ionization energies generally increase going to the right across a row of the periodic table
and generally decrease going down a column of the periodic table.
First ionization energies decrease going down a column of the periodic table because the principle energy level increases,
which means the outermost electrons are being held at a higher energy farther from the nucleus as we go down a column.
So removing an electron requires less energy.
First ionization energies tend to increase going to the right across a row of the periodic table because affective nuclear charge increases.
This means electrons feels a higher in higher positive charge
attracting them to the nucleus as we move to the right across a row.
So it requires more and more energy to separate the negative and positive charges.
Successive ionization energies increase.
It always requires more energy to remove the second electron than the first.
This trend exists because as we remove electrons the resulting ion becomes more and more positive,
so it requires more and more energy to remove a negative electron from an increasingly positive ion.
There's a large jump in ionization energies when we transition from removing a valence electron to removing a core electron.
To demonstrate this, we'll examine successive ionization energies for sodium and magnesium.
A neutral sodium atom has one valence electron in the 3s subshell
and a neutral magnesium atom has two valence electrons in the 3s subshell.
So when we remove the first electron from sodium and magnesium, we're removing a valence electron from each atom.
So we follow our periodic trend that ionization energy increases up into the right in the periodic table.
Thus, it requires more energy to remove the first electronic from magnesium than from sodium.
After we've removed the first electron, sodium plus is leftover and it only has core electrons
but magnesium plus still has one valence electron.
It requires significantly more energy to remove the core electron than a valence electron,
so sodium's second ionization energy is larger than magnesium's second ionization energy.
After we've removed the second electron both sodium two plus and magnesium two plus only have core electrons left.
When comparing removing a core electron between these two ions,
the third ionization energy of magnesium is a larger than the third ionization energy of sodium.
So if we're comparing the relative ionization energies of two elements
and we're removing the same type of electron--either valence and valence electron or core and core electron--
we can follow the general trend that ionization energy increases up into the right in the periodic table.
But if we're removing a core electron from one element and a valence electron from the other,
it should require more energy to remove the core electron.
We can see this trend as we look at the successive ionization energies of our third row in the periodic table.
With sodium we begin removing core electrons with the second ionization energy.
With magnesium we begin removing core electrons with the third ionization energy.
For aluminum it's with the fourth ionization energy.
And so on and so forth.
When we transition from removing a valence electron to removing the core electron,
we can see that there is a very large a jump in the ionization energy.