Course Lecture Topic Information

Chem 105 Web Site
Dr. Kimberly Lawler-Sagarin
Elmhurst College

There will be no discussion for the week of February 23 through February 29
Here are some virtual lecture notes to help you with Homework 1

• Chapter 3: Most important - Section 3.2 (pg 59-67) Also - Section 3.1 has an introduction to the periodic table which you may want to skim through.

### Electrons and their Configuration in Atoms

As we learned last week, electrons occupy different energy levels or "shells" in an atom. This week, we will learn a little about how the electrons are arranged in the atom in more detail.

Atomic energy levels are discrete (or not continuous) - that is, only certain values for the energy are allowed. Electrons in the lower energy levels are generally closer to the nucleus. Each energy "shell" is given a number (1, 2, 3, etc.) designated by the letter "n". Each electron energy shell can only hold a certain number of electrons.

The higher the energy level, the larger the number of electrons it can hold.

 Electron Shell (n) Maximum # Electrons (2n2) 1 2 2 8 3 18 4 32

### Electron Shells and Subshells

The number n, associated with each electron energy shell is known as the principle quantum number or energy quantum number.

Each electron shell can be further broken down into one or more smaller units called subshells.

Each of these subshells also has an associated quantum number, l.

• For a given shell, l can take on any integer value from 0 to n-1.

• Each value of l is associated with a subshell.

• We also use letters (s, p, d, f...) to identify these different subshells. s is used for l=0, p for l=1, etc.

• Each subshell can be represented by a symbol which combines the numerical value of n (1,2,3, etc.), with the symbolic value for l (s,p,d,f), as in:
3p    for n=3   and  l= 1 = symbol p

One can determine the number of subshells for a given shell from the value of n for that shell.

 Example: The first, or lowest, energy shell is the n=1 shell. Given that l can take on only integer values from 0 to n-1: l=0 is the only allowed value for l for this shell (n-1=0 when n=1). The symbol for l=0 is s Thus, the n=1 shell is composed of only one subshell, with the symbol: 1s Another Example: The next highest energy shell is the n=2 shell. Given that l can take on only integer values from 0 to n-1: There are two allowed values for lL l=0 and l=1 (n-1=1 when n=2). Thus, the n=2 shell is composed of two subshells, with the symbols: 2s and 2p.

Here's how things break down for the first four energy shells. Note that everytime you increase "n", an additional subshell is added.

 Shell (n) allowed l(0 to n-1) subshells 1 0 1s 2 01 2s2p 3 012 3s3p3d 4 0123 4s4p4d4f

### Energy Ordering of the Subshells

Within a given shell, the ordering of the energy of the subshells is given by:

s < p < d < f

Each subshell consists of one or more orbitals. These orbitals can be thought of as representing, in an average way, the region of space that the electrons inhabit. We'll be talking about these more later in the course. The type of subshell determines the number of orbitals in the subshell and the total number of electrons the subshell can hold, as follows:

 Subshell Type # orbitals max # electrons s 1 2 p 3 6 d 5 10 f 7 14

You may notice the the number of electrons is twice the number of orbitals. This is because each orbital can hold two electrons.

Putting this information together with the number of subshells allowed in each shell, we come up with a more detailed picture of how many electrons are allowed in each energy shell.

• Shell 1 (n=1): one subshell, the 1s. An s subshell holds 2 electrons, so a maximum of 2 electrons fit into the n=1 level.

• Shell 2 (n=2): two subshells, the 2s and 2p. An s subshell holds 2 electrons, a p holds 6, so a maximum of 8 electrons fit into the n=2 level.

• Shell 3 (n=3): three subshells, the 3s, 3p and 3d. An s subshell holds 2 electrons, a p holds 6, and a d holds 10, so a maximum of 18 electrons fit into the n=3 level.

Here's the above in tabular form with the addition of the n=4 level.

 Shell (n) subshells e-'s in subshell total e-'s in shell 1 1s 2 2 2 2s2p 26 8 3 3s3p3d 2610 18 4 4s4p4d4f 261014 32 . . . . .

### Writing Electron Configurations

The electron configuration for an atom is a way of writing down how electrons of an atom occupy the different energy shells and subshells.

Electron configurations for atoms are written by first grouping electrons into subshells. Next, write the designation of each subshell (nl) followed by a superscript indicating the number of electrons in that subshell.

n l (number of electrons in subshell)

(n = 1,2,3...., l = s,p,d,f....)

Terms for each subshell are placed after one another until all subshells containing electrons are included. An electron configuration looks something like this:

1s22s22p63s23p6......

But, how do we get there? Use the following rule:

• Electrons fill the lowest energy subshell available. Once that is full, electrons are placed into the next lowest subshell available.

Let's start with hydrogen (H). Hydrogen, with atomic number 1, has 1 electron. The lowest energy subshell is the 1s level, so the electron is placed there, and the configuration looks like:

1s1

Helium, atomic number 2, has 2 electrons. The second electron fills the 1s subshell (remember, s subshells can only hold 2 electrons).

1s2

Lithium, atomic number 3, has 3 electrons. The first two filled the 1s orbital, so the third one has to go into the next subshell, the 2s:

1s22s1

Next in line is Beryllium (Be), with 4 electrons.

1s22s2

When we get to boron (B), with 5 electrons, we have to begin a new subshell because the 2s is filled. The 2p in next in line:

1s22s22p1

As we proceed up in atomic number, we fill he 2p subshell and eventually have to start filling the 3s level. The electrons for the first 18 elements are placed into subshells in the following order: 1s, 2s, 2p, 3s, 3d. Their configurations are below:

 Electron configurations for the first 18 elements Element Atomic Number Electron Configuration H 1 1s1 He 2 1s2 Li 3 1s22s1 Be 4 1s22s2 B 5 1s22s22p1 C 6 1s22s22p2 N 7 1s22s22p3 O 8 1s22s22p4 F 9 1s22s22p5 Ne 10 1s22s22p6 Na 11 1s22s22p63s1 . . . Al 13 1s22s22p63s23p1 . . . Cl 17 1s22s22p63s23p5 Ar 18 1s22s22p63s23p6

### Noble Gas Notation

A simplified way of writing electron configurations makes use of the noble gas notation.

Consider the electron configuration of Li:

1s22s1

Noting that 1s2 is the electron configuration of He, this can also be written as:

[He]2s1

Similary, for sodium, we can replace 1s22s22p6 with [Ne]:

[Ne]3s1

 Element Atomic Number Electron Configuration Be 4 [He]2s2 F 9 [He]2s22p5 Ne 10 [He]2s22p6 = [Ne] Cl 17 [Ne]3s23p5 Ar 18 [Ne]3s23p6 = [Ar]

### Beyond the first 18 elements

Beyond the first 18 elements, things get a little more complicated. Basically, the energy shells overlap. For example, the 4s subshell actually has a lower energy than the 3d. Thus, the n=3 and n=4 levels overlap a bit. Electrons will fill the lowest energy subshell, so they will fill the 4s BEFORE the 3d. Below is a table of the energy ordering for the various subshells. As you go up vertically in the table, the subshells increase in energy. Note the overlapping shells.

### Energy Level Ordering for Subshells

 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s

### Electron Configurations After Element 18

The 4s subshell lies lower in energy than the 3d subshell, thus it is filled before the 3d.

The electron configuation of potassium (K), element 19 is:

1s22s22p63s23p64s1

Using the simplified ''noble gas notation'' we can write this in the simplified way:

[Ar]4s1

where [Ar] represents the electron configuration of an argon atom.

The electron configuation of calcium (Ca), element 20 is:

1s22s22p63s23p64s2

or

[Ar]4s2

Starting with scandium (Sc, atomic number 21), we begin to fill the 3d:

1s22s22p63s23p64s23d1

Once the 3d is filled, we can return to the n=4 shell by beginning to fill the 4p subshell with the element gallium (Ga, atomic number 31):

1s22s22p63s23p64s23d104p1

A table of the elements K through Kr is given below. Note the there are a few exceptions due to the fact that 1/2 filled subshells are especially stable, hoever this is not something you will be held responsible for.

 Element Atomic Number Electron Configuration K 19 [Ar]4s1 Ca 20 [Ar]4s2 Sc 21 [Ar]4s23d1 Ti 22 [Ar]4s23d2 V 23 [Ar]4s23d3 Cr 24 [Ar]4s13d5 - note: 1/2 filled subshells Mn 25 [Ar]4s23d5 Fe 26 [Ar]4s23d6 Co 27 [Ar]4s23d7 Ni 28 [Ar]4s23d8 Cu 29 [Ar]4s13d10 - note: 1/2 filled s subshell Zn 30 [Ar]4s23d10 Ga 31 [Ar]4s23d104p1 Ge 32 [Ar]4s23d104p2 As 33 [Ar]4s23d104p3 Se 34 [Ar]4s23d104p4 Br 35 [Ar]4s23d104p5 Kr 36 [Ar]4s23d104p6 = [Kr]

* 1/2 filled subshells are especially stable