ATOMS #3: ELECTRON CONFIGURATION
This section is an easily testable material, so it is important to grasp its basic concepts.
^Recall that we can determine the energy contained within each energy shell around an atom using this formula.
En = Energy contained within the "n" shell (n = 1,2,3,...)
RH = Rydberg constant (H for hydrgoen)
n = energy shell or level
^Shown in the middle (blue circle) marks the nucleus containing neutrons an protons (net positive charge). Around the nucleus are energy shells (n = 1, 2, 3,...). As you go up the energy shell, the energy increases.
QUANTUM NUMBERS
^What to catch from the above slide:
The energy formula (En = -RH/n^2) is important to know and be able to utilize
You also need to know the relationship between the various quantum numbers
"n" always starts from 1 and goes up as whole numbers (whole numbers mean no decimals)
"l" always starts from 0 and goes up as whole numbers till "n-1". Meaning if a given "n" value is 2 for instance, the possible "l" values are 0 and 1. Likewise, if a given "n" value is 5 for instance, the possible "l" values are 0, 1, 2, 3, 4.
"ml" is a range of values from -l to +l including "0". If the "l" value is 1 for instance, then the possible "ml" values are -1, 0, +1. If the "l" value is 3 for instance, the possible "ml" values are -3, -2, -1, 0, +1, +2, +3.
^The 2 images above show the basic orbital shapes (denoted by "l") and how many electrons it can accommodate:
"l" value is what determines the shape of the orbital: the most commonly asked questions on the exam would be shapes of "s", "p" and "d" orbitals, which are sphere, dumbbell, and double dumbbell, respectively as shown above. MEMORIZE THEM
The bottom image shows how many electrons each orbital shape can accommodate. For instance, the "s" orbital has a total of 1 orbital and can accommodate 2 electrons. "p" orbital has a total 3 orbitals and can accommodate electrons. Likewise, "d" = 5 orbitals = 10 electrons, and "f" = 7 orbitals = 14 electrons. MEMORIZE THESE AS WELL
^What to catch from the above slide:
There is the 4th quantum number which is defined as "electron spin" (denoted "ms"). This is by far the easiest to grasp out of all the other quantum numbers.
ms can only have 2 possible values: +1/2 and -1/2. These simply represent the two different directions in which electrons can spin (up and down).
Now that we have covered all 4 quantum numbers, we can bring up "PAULI EXCLUSION PRINCIPLE"
Pauli exclusion principle (TESTABLE**): No 2 electrons can have same set of 4 quantum numbers within an atom. I always compare this with a barcode. We cannot have an identical barcode for different products in the store. It is similar with electrons as well.
^No two organisms can have the exact same DNA. The barcode analogy can be applied to both the distinctness of our DNA as well as electrons within an atom.
SUMMARY OF QUANTUM NUMBERS
ELECTRON CONFIGURATION
^What to catch from the above slides:
So how does electron configuration work? Well, it will be done purely by analyzing the periodic table (look at the bottom image). The periodic table is categorized into different blocks, representing which specific orbital shape ("l") it represents. And you simply count from left to right
Take HYDROGEN for example. Its electron configuration (EF) is simply 1s^1. Moving onto the next element to the right, Helium, its EF is 1s^2. And this continues on and on as we move down and right within the periodic table.
^What to catch from the above slide:
What is important to note is the EFs for each element written on the bottom right (includes all elements up to end of row 3). EF is cumulative, meaning you don't just assign "2s^1" for Lithium, but you also write down the previous row "1s^2". This helps keep count of the total electrons within a given atom (TIP: SUM OF EXPONENTS SHOULD GIVE YOU TOTAL # OF ELECTRONS WITHIN THE ELEMENT) --> 1s^2 2s^1 has a sum of 3 for its exponents, which precisely represents the number of electrons within Lithium)
^We can also use "noble gas configuration". As you can imagine, as we go down and right of the periodic table, the EF will get longer and longer. We can't be spending hours writing down the entire notation, which is where we write the NOBLE GAS FROM THE PRVIOUS ROW [in square brackets] and WRITE THE REMAINING EF OF THE ROW THE ELEMENT IS IN.
Above image shows sodium for instance. Instead of writing out the whole notation, we write [Ne], which is the noble gas from previous row, and write the remaining configuration of the row sodium is in, which is 3s^1.
^Here are other examples of the "noble gas configuration" also known as the "abbreviated electron configuration"
FILLING IN "d" ORBITALS
^What to catch from the above slides:
Now, you should notice that the "d" orbital starts from row 4.
I initially questioned why the "d" orbital was notated before the "s" orbital as shown in the first two slides, and the simple reason could be to match the "n" value to that of the previous orbital "3p^6", which has an "n" value of 3. But do note that "4s" has slightly lower energy than "3d" and thus should be filled first based on the Aufbau's principle --> what I personally recommend is to write the "4s" before the "3d", which is different from what is shown on the top two slides
Aufbau's principle (TESTABLE**): fill the subshells of the lowest energy level first before filling higher energy levels.
^What to catch from these images:
Another important concept is the element's or atom's preference to be HALF-FILLED OR FULLY FILLED to reach stability!
Take for example Cr: its abbreviated EF would be [Ar]4s^2 3d^4. But because the "d" orbital wants to be half-filled to achieve stability, one electron from "4s" jumps to the "3d" to give it "3d^5" and "4s^1". When they ask you to write the EF for Cr, the correct notation then becomes [Ar]4s^1 3d^5. MEMORIZE THIS
Likewise, in copper (Cu), one electron jumps from the "4s" to "3d" to FULLY FILL with electrons to achieve stability. Therefore, the notation becomes [Ar]4s^1 3d^10 as seen in the image.
PRACTICE WRITING ELECTRONIC CONFIGURATIONS
ANSWER KEY
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