1. Electron Configuration and Periodic Properties
The Periodic Law
Electron Configuration and Periodic Properties

2. Atomic Radii
It is difficult to measure the radius of an atom because the edge of the electron cloud is fuzzy and varies under different conditions.
One way to measure radius is to measure the distance between the nuclei of two identical atoms that are chemically bonded together and dividing this distance by 2.

3. Atomic Radii The definition of atomic radius is:
One-half the distance between the nuclei of identical atoms that are bonded together.

4. As you go across a period, the atomic radius tends to get smaller.
This occurs even though more electrons are being added.
As electrons are added to the s and p sublevels in the same main energy level, they are gradually pulled closer to the nucleus which is getting more positively charged.

5. As you move down a group, the atomic radius tends to increase.
With each move down a group, another main energy level is added
This increases the size of the atom.

7. Ionic Radii
An ion is an atom that has gained or lost one or more electrons.
Atoms tend to form ions in such a way that their outermost s and p orbitals are full with an octet of electrons.
In other words ions are formed so that the atom attains a noble gas configuration.
This makes the ion more stable than the atom

8. A positive ion is called a cation.
Cations have lost one or more electrons
Cations always have a smaller radius than the neutral atom from which they were formed.

9. A negative ion is called an anion.
Anions have gained one or more electron
Adding an electron increases the radius so anions have a larger radius than the atom from which they were formed.

11. The metals tend to form cations
The nonmetals tend to form anions

12. Valence Electrons
Valence electrons are the electrons in the outer energy level that are available to be gained, lost, or shared in the formation of chemical compounds.
For main group elements, the valence electrons are the ones in the outer s and p orbitals.

13. Group 1 elements (the alkali metals) have one valence electron.
Group 2 elements (the alkaline earth metals) have two valence electrons.

14. Group 13 – 3 valence electrons
Group 17 (halogens) – 7 valence electrons
Group 18 (noble gases) – 8 valence electrons

15. Lithium has 1 valence electron.
Looking at the number of valence electrons is a good way to figure out what type of ion an atom is likely to form
Lithium has 1 valence electron.
That electron is in the 2s orbital
If it loses that electron, then its highest energy level is the first level, and that level is full
Therefore lithium tends to form a 1+ cation

16. What about elements in group 2?
Why don’t elements in Group 18 tend to form ions??
What about group 14?

17. Ionization Energy
An ion is an atom or group of bonded atoms that has a positive or negative charge.
Ionization is any process that results in the formation of an ion.
Ionization Energy is the energy required to REMOVE one electron from a neutral atom
This is called the FIRST ionization energy

18. Some atoms give up an electron more easily than others.
Group 1 elements tend to lose 1 electron easily, giving them a noble-gas electron configuration.
Noble gases (group 18) are difficult to remove an electron from.

19. The second ionization energy is the amount of energy required to remove a second electron from a positive ion.
The second ionization energy is always higher than the first ionization energy

20. Ionization energies generally decrease down the groups.
In general, ionization energies of the main-group elements increase across each period.
It gets harder and harder to remove electrons as you move across a period towards the noble gases
Ionization energies generally decrease down the groups.
When the outermost electron is farther from the nucleus, the electron is easier to remove
This is because the negatively charged electron is further from the positively charged nucleus

23. Electron Affinity
Electron affinity is the energy change that occurs when a neutral atom gains an electron
If that energy is given off, then the value for electron affinity is negative.
If energy must be absorbed for an electron to be gained, then the value for electron affinity is positive.
In each period, the halogens gain electrons most readily.
When a halogen gains one electron, its electron configuration becomes like a noble gas

24. Electron affinity tends to increase across each period.
Group electron affinity trends are not as regular as the trend for ionization energy
In general, electron affinity tends to decrease down a group.

25. It is always more difficult to add a second electron to a negatively charged ion.
Therefore, it always requires an input of energy for the second electron to be added and the values for second electron affinities are always positive.
What is the likelihood that a Chlorine atom will add 2 electrons?

26. Electronegativity
Valence electrons hold atoms together in chemical compounds.
In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another.
This leads to an uneven distribution of charge which has a large effect on the chemical properties of a compound.

27. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in a compound.

28. Fluorine is the most electronegative element.
It attracts valence electrons more strongly than any other atom in a compound
Fluorine’s electronegativity is 4
All other atoms have an electronegativity that is measured in comparison with fluorine

29. Across a period, the electronegativity tends to increase
The alkali and alkaline earth metals are the least electronegative
Since these atoms tend to lose electrons rather than gain them, they don’t tend to attract electrons strongly
Nitrogen, oxygen, and the halogens are the most electronegative
These atoms tend to form anions; they attract electrons strongly in the quest for a noble gas configuration.

30. In a group, electronegativities tend to decrease or stay about the same.
The noble gases are not assigned an electronegativity because they don’t tend to form compounds.