Electrons in the outermost s and p orbitals. These are the ... · Electrons in the outermost s and p orbitals. These are the electrons most often involved in bonding. The organization

Electrons in the outermost s and p orbitals. These are the electrons most often involved in bonding.

The organization of electrons in an atom or ion, from the lowest energy orbital (“s”) to the highest energy orbital (“f”).

s

p

d

f

FULL ELECTRON CONFIGURATION: • ALL electrons are shown, according to the energy

level and orbital type. • Start at n = 1 • Fill each energy level before moving on to the next.

EXAMPLES:

1s1 1s2

1s2 2s1 1s2 2s2 1s2 2s2 2p1

1s2 2s2 2p5

1s2 2s2 2p6 Noble gases have a FULL VALENCE SHELL,

that is, a STABLE OCTET 1s2 2s2 2p6 3s2 3p6 4s2

1s2 2s2 2p6 3s2 3p6 4s2 3d1

Note the change in the numbering as you move from 4s to 3d. Note that all of the exponents add up to the atomic number of the element.

1s2 2s2 2p6 3s2 3p6 4s2 3d8

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

5s2 4d10 5p6 6s2 4f2

Note the change in the numbering as you move from 6s element # 56, Ba to element #57, La which starts to fill its 4f orbitals.

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

5s2 4d10 5p6 6s2 4f14 5d6

Note the change in the numbering as you move from 6s element # 56, Ba to elements #57 through 70, the Lanthanides, which fill all 14 of the 4f orbitals, and then back up to the 5d orbitals that begins with Lu (element # 71) and ends with Hg. (Ac, # 89 to No, # 102 fill their 4f orbitals) (Lr, # 103 to Cn , #112 fill their 6 d orbitals)

CORE ELECTRON CONFIGURATION: • Show the electrons in the energy levels past the PREVIOUS noble gas.

EXAMPLES:

2s1

2s2 2p5

4s2 3d8

4s2 3d10 4p2

5s2 4d5

5s2 4d10

6s2 4f14 5d4

6s2 4f14 5d10

ORBITAL REPRESENTATION DIAGRAMS: Show the pairs of electrons in each orbital.

EXAMPLES:

1s2 2s2 2p5

↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s2 2s2 2p6 3s2 3p6 4s2

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

[Ar] 4s2 3d8

↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ [Ar] 4s2 3d10 4p2

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ _

[Xe] 6s2 x4f14 5d4

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ ↑ _

A “drawing” that represents the nucleus and its outermost electrons. Core Electron Config Lewis Dot Diagram

1s1 1s2 ↑ ↑↓

[He] 2s1

↑ [He] 2s2

↑↓ [He] 2s2 2p1

↑↓ ↑ _ _

[He] 2s2 2p2

↑↓ ↑ ↑ _

[He] 2s2 2p3

↑↓ ↑ ↑ ↑

[He] 2s2 2p4

↑↓ ↑↓ ↑ ↑

[He] 2s2 2p5

↑↓ ↑↓ ↑↓ ↑

[He] 2s2 2p6

↑↓ ↑↓ ↑↓ ↑↓

IONS: Atoms usually lose s and p electrons first since these are the outermost electrons.

EXAMPLES: ION CORE CONFIGURATION LEWIS DOT

[He] 2s2 2p6

Added one electron

[Ar] 4s2

Lost 2 electrons

[Ar] 4s2 3d10 4p1

Lost 3 valence p and then s electrons

[Ar] 4s2 3d10 4p2

[He] 2s2 2p6

[He] 2s2 2p6

Sample Lewis dot diagrams for ions:

ELECTRONIC CONFIGURATIONS

EXCITED STATE: Atoms absorb energy and exist at higher energy levels. As the atom returns to a lower energy level, it gives off energy. This relates to emission spectra / spectral lines. Example: Carbon atom.

[He] 2s2 2p2

↑↓ ↑ ↑ _

[He] 2s1 2p3

↑ ↑ ↑ ↑ Electron excited from s orbital to a higher energy level. This makes a more stable arrangement, as having all orbitals half filled is more stable than having some orbitals with pairs, some with singles, and some empty. This means that C can easily combine with 4 atoms such as H to make CH4.

More on this in Organic Chemistry.

Consider the element GALLIUM, atomic number 31. This neutral (uncharged) element has 31 electrons.

We NOW know that each of the electrons exist in complex three dimensional orbitals: Since Gallium is in the 4th row of the periodic table: n = 4 0 < ℓ < n – 1, so ℓ= 0, 1, 2, 3, meaning that orbitals represented by EACH of the l values, s, p, d, f, are possible . (But they may not all be filled by electrons!!!) The mℓ and ms just tell us the orientations in space of the orbitals and the spin of the paired electrons so it is NOT necessary to look at these each time you write an electron configuration.

NOTE: The  complete  understanding  of  this  concept  will  occur  over  several  classes,  and  requires  you  to  practice,  daily,  until  you  reach  proficiency.      Reading  these  pages  is  not  enough  to  gain  a  full  understanding.      This  only  scratches  the  surface.  

In fact, the practice that we will do in class will put this into a more practical approach, rather than the mathematical approach as indicated by the quantum numbers, above.

Full electronic configuration Showing ALL of the electrons in every energy level from n = 1, for Ga: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1 Core electronic configuration Showing electrons of an element, starting at the PREVIOUS noble gas [Ar] 4s2 3d10 4p1

Orbital representation diagram Showing paired electrons with opposite spins, in each orbital:

Some textbooks, and even teacher worksheets, show full electron configurations as ENERGY DIAGRAMS, so that the energy differences of the electrons in each level is clearly demonstrated:

Practice done in class, with the teacher, on the white board, is essential for the understanding of electron configurations. We will start with basic FULL configurations, understanding the progression of the electron configurations from Hydrogen to Helium to Lithium, to Beryllium, etc. We will learn how to best distribute electrons in p orbitals so that we know when to have paired and unpaired electrons. As we enter the 3rd energy level, and d orbital electrons are possible, it is necessary to understand how to write the MOST STABLE electron configurations. And entering the 4th energy level, which introduces f orbital electron distributions, we must understand how an element would fill it’s orbitals in the proper energy level order. Practice in class will be accompanied by three dimensional images and videos on screen, as well has examples of emission spectra, so that you can understand this work that you are doing, IN CONTEXT. The sample questions will evolve organically, as we develop understanding of each aspect, so providing a full notes package on the practice questions would not be beneficial to your understanding, nor is it possible to predict each question in the order it will be presented. Once we are comfortable, we will also introduce ions, so that we understand WHERE an element would lose or gain electrons from, in order to achieve the most stable configuration. This will lead us into the ionic and covalent (molecular) bonding that we will be studying next. You will learn better by being able to work through this unit in consultation with your lab partners.