Electron Orbital Diagrams. Placing Electrons in Orbitals. Every orbital may hold up to 2 electrons These 2 electrons have opposite spins: an up spin and a down spin You can determine how many electrons are possible for an energy level by 2n 2 If n=1, only 2 electrons are possible
Noble Gas Configurations
Orbital energy diagrams tell us what electrons/orbitals an element fills in its ground state (most stable state).
But they are tedious to draw and they must be in a figure.
So how can we get the same info but be able to write it in a sentence?
Electron configurations let us do this.
All we do is translate the vertical orbital diagram into a horizontal line. (well, almost all)
What about the electron configuration for Sr?
When you write the correct electron configuration for the ground state of Sr, notice that you group the n values together.
This is different from how you wrote the orbital diagram which is strictly based on filling order.
Why do we write the electron configuration this way?
Convention and because that way the valence electrons are at the end of the configuration: and valence electrons are lost first!
But even electron configurations get long and tedious to write.
Write the electron configuration for gold!
How can we make this shorter?
Well, the Noble Gases have filled s and p subshells, so they are good reference elements.
So we write Noble Gas configurations, which are based on the LAST Noble Gas BEFORE the element we are working with.
The Periodic Table is laid out in order of increasing Atomic Number.
It is also organized by the order in which the electrons fill! (Except for the exceptions.)
Thus, we have the s-block, the p-block, the d-block, and the f-block.
Electron Configurations and
Members of a group have similar chemical properties. Why?
They have the same number of valence electrons, or their outermost electron configurations are the same.
And valence electrons are responsible for chemical properties.
Valence electrons are outer shell electrons and so are furthest from the nucleus (generally).
Thus they are not as tightly held to the nucleus.
So they are removed more easily than core electrons (inner shell electrons).
Also, the core electrons with their negative charges block or insulate or "shield" the valence electrons from feeling the full nuclear charge.
So the shielded valence electrons don't feel the full nuclear charge, Z, they instead feel what we call Zeff, or the effective nuclear charge.
Think of picking up iron filings or paper clips with a magnet, eventually the magnet can’t attract any more filings as it is “shielded”.
Members of the same group have the same number of valence electrons and have similar properties as those valence electrons behave similarly.
For example, all of the Alkali Metals have 1 valence electron, with an outer electron configuration of ns1.
This is why Alkali metals are so reactive, their 1 valence electron is easily removed in chemical rxns.
And what is their e- configuration then?
Alkaline Earth metals all have an ns2 outer electron configuration, or they have 2 valence electrons.
What is the outer electron configuration for halogens?
But chemical reactivity is not the only periodic property or trend which we can see in the Periodic Table.
Another important Periodic Trend or Property is the atomic radii, or the radius of an atom, or 1/2 the diameter of the atom.
We may measure the radius of an atom.
It is also common to measure the distance between 2 nuclei in a diatomic element.
Then the atomic radius would be half of this distance.
There are 2 trends in atomic radii that you can explain and see from the Periodic Table:
As you go across a Period from left to right, the atomic radius of the elements DECREASES.
This is because you are adding protons so Zeff is increasing. As Zeff increases, the outer electrons are pulled closer to the nucleus, so the radius shrinks.
This is particularly true as you add s and p electrons in the same energy level. They don't shield each other very well, so Zeff increases as you add protons.
What about the different orbital types, do they have different attractions, or different Zeff?
For example, for n = 4, there are 4s, 4p, 4d, and 4f orbitals. They don’t have the same energy, do they have different Zeff?
In general, as the subshell energy increases on the same principal energy level, the Zeff decreases.
So for the same n value, the s has the highest Zeff.
As you go down a Group, the atomic radius increases.
Now you are adding an energy shell which by definition is further from the nucleus!