CHEM 1411 - General Chemistry I
Course Lecture Notes
Summer 2007
Instructor: Dr. Shawn Amorde
Web page:
Chemical Periodicity
Text Chapters (6.1-6.6)
Recommended Problems: (2, 8, 10, 12, 18, 22, 24, 28, 33, 36, 42, 50)
The Periodic Table
A. More about the Table
B. Atomic Radii
C. Ionization Energy
D. Electron Affinity
E. Ionic Radii
F. Electronegativity
The Periodic Table
We just discussed electronic configurations; now let’s apply this to the Periodic Table.
Let’s start with Group 8A, the noble gases. The outer most energy level, shell, has eight electrons filling both the s and p orbitals. These gases were known for many years not to react with anything; however recent studies show the heavier gases do sometimes react with F and O.
The Group 1A and 2 A elements all have their last electrons assigned to s orbitals, and show very similar reactivity to each other.
Group 3A-8A elements all have their last electrons assigned in p orbitals, and also have similar reactivity the each element in the individual columns.
The Group B elements all have electrons in d orbitals and are known as the transition elements, or as the d-transition elements.
The Lanthanides and Actinides are the f-transition elements or also known as the inner transition elements. They start with Cerium (Ce) which has an atomic number of 58.
Atomic Radii
We have just discussed how to predict where the electrons are distributed are the nucleus, let us consider the entire atom’s size.
We can’t see an atom to measure the diameter and our calculated energy orbitals do not have a definite “edge”, thus we do not actually know the size of an individual atom.
However, we can measure the distances of atoms in solids and the lengths of chemical bonds, so the atomic radii can be calculated in many elements.
Atomic radii size decreases from left to right, and bottom to top in the Periodic Table.
Doesn’t this seem backwards? We are adding more electrons, but our atom size is getting smaller. Why?
First, let us consider charge attraction and repulsion. The inner most electrons are attracted to the nucleus, but what about the electrons in the outer shells? They “feel” the attraction to the nucleus less because of the counterbalanced repulsion to the inner shell electrons. This is also why the size increases as you move down each column.
As we move from left to right across the table the number of protons increases and “over balances” the charge of the two electrons in the 1s shell.
Ionization Energy
This is the energy required to remove the most loosely bound electron from a gas atom to form an ion.
These energies increase as we move left to right across the Periodic Table and decrease as we move down each column.
Electron Affinity
Electron affinity is the energy absorbed when an electron is added to an atom to form an anion.
Elements with negative electron affinities gain electrons easily to form negative ions (anions).
The electron affinities are little harder to predict in the middle metals, but the halogens are always the most negative. The noble gases and the Group 2A elements are all zero.
Ionic Radii
Cations are always smaller than the neutral atoms, and conversely anions are always bigger than the neutral atoms.
Isoelectronic species have the same number of protons. For example, Li+1 and Be+2 are isoelectronic.
Electronegativity
The electronegativity is the measure of an element to attract electrons to itself when chemical combined (bonded) to another atom.
Elements with high electronegativities often gain electrons to form anions and elements with low electronegativities often lose electrons to form cations.
For the elements, electronegativity increases from left to right across the table and decreases from top to bottom within each group.