1. Periodic Trends

2. There is a correlation between the arrangement of the elements and their electronic configuration.
We will look at the relationship between the electron configurations and the periodic trends of the elements.

3. Atomic Radii
Ideally, the size of an atom is defined by the edge of its orbital.
However, this boundary varies under different conditions. No set boundary.
One way to express the atomic radius is to measure the distance between the nuclei of two identical atoms that are bonded together, then divide this distance by two.

4. Atomic radius may be defined as one-half the distance between the nuclei of identical atoms that are bonded together.

6. Period Trends There is a gradual decrease across a row.
Some exceptions with the transition elements but the general trend holds.
Exceptions are due to electron-electron repulsions which cause the size to increase.

8. There is a gradual decrease across a row.
The trend to smaller atoms across a period is caused by increasing positive charge of the nucleus. Adding of protons or increasing atomic number.
The electrons are pulled closer to the nucleus.
The increase pull results in a smaller atomic radius.

10. Group Trends
There is an increase down a group.
This trend is because as electrons are added for each row of the periodic table, they are further from the nucleus.
They are in higher energy levels.
Some exceptions due occur.

11. Problem
Of the elements magnesium-Mg, chlorine-Cl, sodium-Na, and phosphorus-P, which has the largest atomic radius and why?
Of the elements calcium-Ca, beryllium-Be, barium-Ba and strontium-Sr, which as the largest atomic radius and why?

12. Ionization Energy
Ionization energy (IE) – the energy required to remove one electron from a neutral atom of an element.
A + energy A+ + e-
An ion is an atom or group of bonded atoms that has a positive or negative charge.

14. Period Trends
In general, ionization energies of the elements increase across a period.
Group 1 elements – have the lowest ionization energies. Therefore they lose electrons most easily. Very reactive.
Group 18 elements - have the highest ionization energies . They do not lose electrons easily. Very low reactivity.

15. The increase is due to increasing nuclear charge (more protons going across a period).
A higher positive charge more strongly attracts electrons in the same energy level.
Therefore, it is tougher to remove an electron from an atom.
Increasing nuclear charge is responsible for both an increasing ionization energy and decreasing atomic radius across a period.

16. Group Trends Ionization energies generally decrease down a group.
Electrons removed from atoms of the elements down a group are farther from the nucleus.
Also, the electrons from the lower energy levels shield the outer electrons.
Therefore, they are removed more easily.

17. Photoelectron Spectroscopy
Higher
energy
Lower
energy

18. Photoelectron Spectroscopy

20. Second and Third Ionization Energies

21. Electron Affinity Neutral atoms can also acquire electrons.
Electron Affinity – the energy change that occurs when an electron is acquired by a neutral atom.
A + e A- + energy
The quantity of energy released is represented by a negative number.

22. Period Trends
Among the elements of each period, the halogens (Group 17) gain electrons most readily.
As the number of positive protons increase across a period, it is easier to add electrons.
As electrons add to the same p sublevel of atoms, electrons affinities increase.

23. Period Trends
An exception to this trend occurs between Group 14 and 15.
Compare the electron affinities of carbon ([He] 2s2 2p2) and nitrogen
[He] 2s2 2p3).
Adding an electron to nitrogen is more difficult because it forces the electron to pair with another electron.

24. Group Trends Electron affinities within groups generally decrease.
The electron being added is further from the positive nucleus so the attraction to the nucleus is less.
Fluorine has the highest electron affinity.

25. Ionic Radii