*Humboldt State University* ® *Department of Chemistry*

Richard A. Paselk
Chem 107	Fundamentals of Chemistry	Fall 2009
Lecture Notes: 14 October	© R. Paselk 2005

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Electrons in Atoms, cont.

For chemistry on Earth we can assume that all atoms are in their ground states. That is, each of the electrons in an atom will be at the lowest energy it can attain.

What is the basis of the periodicity of properties?

As we noted above, electrons are arranged in an atom into specific energy levels. These energy levels are called shells.

A shell indicates the average distance of its electrons from the nucleus, since higher energy electrons are more loosely held (much like planets in the Solar system, where higher energy = faster speed moves planets away from the Sun).

At energy levels greater than n = 1 get subshells, with the theoretical number of subshells = n. So for n = 1, one subshell, for n = 2 two subshells, etc.

Within the shells and subshells electrons occupy geometrical regions of space called orbitals.

The lowest energy orbital within each shell is spherically symmetrical and is called an s-orbital.
At a slightly higher energy, but still in the same shell (new subshell), are the p-orbitals. The p-orbitals show planar symmetry with two lobes. There are three p-orbitals arrayed along mutually perpendicular axis (x, y, and z), giving p_x, p_y, and p_z orbitals with each p-subshell.

Note that for atoms in the grond state all filled shells and subshells have spherical symmetry.

So let's look at the elements in the Periodic Table in light of this model.

The first shell holds only 2 electrons in what is called the 1s orbital. Thus helium has its first shell filled. There is no more room for electrons, so it can't react by picking up another electron. On the other hand, as a crude thought model, we can consider that each electron is held by both charges in the He nucleus, so they are much more tightly held than the electron in H, so He won't give up an electron either - its inert.
The second shell is larger (its out further from the nucleus) so holds 2 electrons in a 2s orbital, but there is now room for an additional three 2p orbitals. Thus 8 electrons can be accommodated in the second shell. Note the consequences:
- for lithium (Li) the inner two electrons of the 1s shell cancel the attraction of two of the three protons, so the outer 2s electron "sees" only a single charge. But its out further than the electrons in the 1s shell were, so its not held as strong, so Li loses its outer electron more readily than H and is more reactive.
- for Fluorine on the other side of the chart we can think of the outer shell electrons being attracted to the nucleus by 9 - 2 = 7 charges, so the last open space in an orbital will be super attractive to an outside electron, so F will be be very reactive, but in an opposite way to Li - it wants to steal electrons instead of giving them up.
- for neon all of the orbitals will be filled, and the electrons will be strongly attracted to the nucleus, so Ne will again be inert, like He above it.
The third shell is larger yet (further from the nucleus), but still crowded, so initially it can only accommodate another eight electrons.
- Of course the electrons in the 3s orbitals are even farther out from the nucleus, so we would expect Na to be even more reactive than Li, and so on for K, Rb, etc. each giving up its outermost electron more readily than the element above it in the Periodic table.
- On the other hand Cl will also attract electrons less than F, so it will be less reactive etc. for the halogens.
- So we will expect the most reactive elements to be on the opposite corners of the table - lower left and upper right.

The second Period is special, having greater tendency to form covalent bonds (share electrons). In particular the second period elements tend to from the strongest covalent bonds - C is the only element to form strong, stable, multiple covalent bonds to itself, making C based polymers possible.

A Quantum Picture of Atomic Orbitals

A brief introduction to orbitals is illustrated with QuickTime movies based on quantum calculations in the Atomic orbitals supplement.

Electronic Configurations

How do we figure out where electrons go in atoms? Two common methods:

Follow along the Periodic Table, noting the "s-block", "p-block", and "d-block" elements in the table.
Use the Aufbau (filling) diagram:

7s 7p
6s 6p 6d
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s

and draw diagonal arrows at an angle such that they parallel np - (n+1)s. The order at which the arrows cross the orbitals gives the filling order.

There are a number of different notation conventions for electronic configurations.

Spectroscopic notation: in this convention we indicate shells (main energy levels) by numbers, orbitals within these shells by letters (s, p, d, or f), and the number of electrons in each orbital type by superscript. For example: (Slide)
- H:
- He:
- B:
- P:
- V:

Note that when we get to the d electrons they are added to the next inner orbital - they are added inside the atom and are not outermost! Note also that you may write them in the order they show up on the Periodic Table, or, if you prefer, you may group them by shell

Spectroscopic notation using the Noble gas core convention. Notice that at the end of each period the outermost shell (s & p) is filled, and when you go to the next element (e.g. Na) its as if you are adding onto the electrons configuration of the Noble gas. So we can save a lot of writing if we substitute its symbol for these inner electrons (note we are not really assuming an inner noble gas, we are just creating a type of short-hand). For example: (Slide)
- P:
- V:
- U:

The savings here is really obvious! Note that if you use the Periodic chart that comes on exams you will fill the f's before the d's of a given period (this is not the case for the wall chart).

Orbital Filling Diagrams: In orbital filling diagrams we provide a little more information - noting that the electrons come in two different spins and that they fill into orbitals with their spins paired in opposite directions. I have no preference for the direction of unpaired arrows other than that they should be the same. For example: (Slide)
- H:
- He:
- Li:
- N:
- O:

Orbital Filling Diagrams using the Noble gas core convention. You can also use the Noble gas core convention with orbital filling diagrams, just like the spectroscopic notation, but using arrows etc.

So how do we decide which electrons are lost in ionization?

Ions: When an atom loses electrons we would expect it to lose its outermost electrons first. But which are outermost? Remember the "last added" electrons in the transition elements at in the d orbitals of the next outermost shell. Thus the d orbital electrons should not be the outermost electrons in an atom. Thus we will lose the s & p electrons first then the d electrons if any are present. If additional electrons are lost then we can go into the d shell. Examples: (Slide)
- Na⁺ =
- Cu²⁺ =
- Fe²⁺ =
- Br^- =

We also find a number of "special cases where filling diagrams and tools such as the Aufbau diagram fail in their predictions. What's going on? We will leave that as a mystery for other course such as Chem 109.

Periodic Table and Periodic Properties

Read Chapter 7 on the Periodic Properties of Elements

Let's look again at trends for: (use Clickers to discuss vertical, horizontal, and diagonal trends)

Atomic size = Atomic radii (hi - lo goes from Cs - F).
- The pattern shown in the plot below:

plot of atomic radii versus atomic number

Can be represented by the trend shown on the Periodic Table:

First ionization energy (lo - hi goes from Cs - He).
- The pattern shown in the plot below:

plot of ionization energy versus atomic number

Can be represented by the trend shown on the Periodic Table:

Electronegativity (Cs, lowest to F, highest).
- Electronegativity is a measure of how electrons are shared between two associated atoms. The pattern shown in the plot below:

plot of electronegativity versus atomic number

Can be represented by the trend shown on the Periodic Table:

Two additional properties:

Density (bump in middle of transition elements).
- The elements with the highest densities under standard laboratory conditions are shown on the Periodic Table below:

This pattern can be explained by a combination of the atomic sizes (shown on the plot below),

the number of nucleons (increases left - right), and the physical states of the elements. Thus densities of elements on the extreme right are low since they are gases, etc.

Melting points
- Highest melting points for elements include C and a bump near the middle of the transition elements as shown on the Periodic Table below: