Electronegativity

We met electronegativity briefly last year. It is a relative measure of how tightly an atom attracts electrons in a chemical bond.

We use the Pauling Scale to represent the strength of this attraction. Fluorine is the most electronegative element with a value of 4.0.

The difference in electronegativity can be used to infer the type of bonding between atoms. The higher the difference, the more ionic a bond will be in nature (transfer of electrons). The closer the electronegativity of the two atoms (to each other) the more covalent a bond will be in nature (sharing of electrons).

We need to know:

1. A definition for electronegativity
2. The periodic trend for electronegativity
3. The reasons for these trends (relating to atomic structure)

Ionisation Energy

1st Ionisation Energy values have helped us infer the structure of atoms. We need to be able to use our understanding of atomic structure to justify the periodic trends (and exceptions) in 1st Ionisation Energy.

Ionisation Energy is the amount of energy required to remove one mole of electrons from one mole of atoms (of one element) in its gaseous state.

Ionic Radius

When cations form, they have a smaller radius than their parent atom.

When anions form, they have a larger radius than their parent atom.

Why?

Periodic Trends - Atomic Radius

We have to use our understanding of atomic structure to explain some trends we notice on the Periodic Table of Elements. Today, we looked at Atomic Radius:

Energy Levels

How are the electrons arranged in the energy levels?

Three Principles/Rules guide us with understanding this:

There are two exceptions: Cr and Cu:
Cr: 4s¹ 3d⁵
Cu: 4s¹ 3d¹⁰

One electron from the 4s sub-level "drops" into the 3d sub-level to make the electron arrangement more stable. A half-filled sub-level is stable due to "maximum spin". A complete 3d sub-level is also stable as it "fills" that energy level.

The last skill we need regarding electron configurations is to apply this to ion formation. What are the electron configurations of ions, and why are some coloured?

Quantum Mechanical Model of the Atom

For the last four years, we learned a very simplified model of the atom. It worked very well to explain the chemical properties of the first 20 elements, alkali metals, earth metals, carbon, the halogens and the Noble gases. However, we are about to venture in to the world of the Transition Metals and other atoms that do not obey the Octet Rule. For this, we need a better model...

What has stayed the same?

1. The Nucleus. As chemists, we really only focus on the protons, as they attract electrons.
2. Electron Shells. While it would be more correct to think of these as "Energy Levels", we can still imagine electrons existing in "shells" at different distances from the nucleus. We still (generally) fill these from the inside-out. The first energy level (shell) still takes 2 electrons. The second still takes 8 electrons. After we put 8 in the third energy level, we start using the fourth energy level. However, things get interesting at #21 (Scandium). 

What has changed? What is new?

1. Sub-levels (sub-shells). The electrons are arranged into sub-levels or sub-shells. In the first Energy Level, there is only one sub-level (called "1s"). In the second Energy Level, there are two (called "2s" and "2p"). The third and fourth Energy levels each have three ("3s", "3p", "3d", "4s", "4p" and "4d").
2. Electron configuration. We use the sub-levels to write electron configurations now. For example, magnesium used to be 2,8,2. Now it is 1s² 2s² 2p⁶ 3s²

We are going to use our understanding of atomic structure to:

• give electron configurations of ions
• explain properties of transition metals (and their ions)
• explain periodic trends (such as electronegativity and atomic radius)