1. Electron Configurations and Periodicity
Chapter 8
Electron Configurations and Periodicity

2. Electron Spin
Electron spin was first observed in by Otto Stern and Walter Gerlach.
A beam of silver atoms was directed in the magnetic field of a specially designed magnet.

3. Electron Spin
The same thing can be done with hydrogen atoms, and when it is, the beam is split in two—half going in one direction, half in the other direction.

4. Electron Spin
The reason this happens is because the electrons in each atom behaves as a tiny magnet with only two possible orientations.
The resulting directions of spin magnetism are +½ and -½.
Atoms have a particular electron configuration—which is the distribution within the available subshells.

5. Electron spin ↑ms +½; ↓ms -½
s,p, & d are examples of subshells. They have coefficients indicating which subshell it is in along with a superscript telling how many electrons are in each one.
We can draw orbital diagrams to represent this. Below is the diagram from carbon:
↑ms +½; ↓ms -½

6. The pauli exclusion principle
The Pauli Exclusion Principle summarizes experimental observations stating that no 2 electrons in an atom can have the same 4 quantum numbers, n, l, ml, ms.
In other words, you can’t write +½, +½, because electrons can’t have the same spin.
Because of this, ms can only have values (+½ & -½)

7. The pauli exclusion principle
Also, each subshell holds a maximum of 2x the number of electrons as there are orbitals in the subshell:

8. The pauli exclusion principle
For example, a 2p subshell, which has 3 orbitals, (ml = -1, 0, and +1) can hold a maximum of 6 electrons.

9. Electron configurations
Atoms have an infinite number of electron configurations. The lowest energy level is called the ground state, any other state is referred to as an excited state.
The normal electron configuration of Na is 1s22s22p63s1.

10. Electron Shell Filling
When filling electron shells, we follow the aufbau principle. This gives us the electron configuration of the ground state of atoms.

11. Observed Patterns
As we pass from period to period, some of the patterns we can observe are that with the Noble gases, the p subshell is completely filled, and with the alkaline metals, the s subshell is just filled.

12. Observed Patterns
Each of the alkaline earth metals (IIA) are said to have a noble-gas core, meaning that there is an inner shell configuration equivalent to the nearest gas + 2 outer electrons.
For example, the electron configuration for magnesium is:
1s22s22p63s2
Or,
[Ne]3s2

13. Valence Electrons and pseudo-noble gas configurations
Elements such as gallium that have an (n-1)d10 electron configuration are called pseudo-noble-gas cores because these electrons are rarely involved in chemical reactions.
Any electrons outside of the noble-gas or pseudo- noble-gas core are referred to as valence electrons and are the electrons primarily involved in chemical reactions.

14. Elements and their properties
Many of the elements in the same column have similar electron configurations, and hence similar chemical properties.
Alkali metals
Main group elements have s and p subshells that are being filled. (Valence configuration ns2np6)

15. Transition elements
Transition elements (d-block elements) the d subshells are being filled.

16. Inner transition elements
The inner transition elements (lanthanides and actinides, f-block elements) have the f- subshell being filled.
There are some exceptions to the building up rules, but we need not be too concerned with them.
For example: Cr is predictd to be [Ar]3d44s2, but is experimentally shown to be [Ar]3d54s1.
I prefer that you write the electron configurations according to the Aufbau principle, (the building up order.)

17. Orbital diagrams
When we write orbital diagrams, recall no 2 electrons can have the same spin. Also, we must keep in mind Hund’s rule.
Hund’s rule says that electrons be placed in their subshells with an arrangement depicting the lowest possible energy configuration and the same spin first.
For example: Carbon, z = 6, 1s22s22p2

18. Orbital diagrams

19. mendeleev
Mendeleev organized what was known about the elements as of 1869 into a table. He noted their similar properties and left open spots between some of the elements that were undiscovered but he believed (correctly) they would eventually be.

20. mendeleev
He made predictions about these undiscovered elements masses and he was remarkably close.

21. Mendeleev and His Table’s predictions
For instance, Mendeleev predicted Eka- aluminum, an element under Al, but not yet discovered to weigh/have a mass of 68 amu from the oxide called Ea2O3, to have a density of 5.9 g/cm3, and have a low melting point and a high boiling point.
When Paul-Emile Lecoq de Boisboudrar discovered gallium in 1874, it had a mass of amu, had an oxide Ga2O3, and a density of 5.91 g/cm3, a melting point of 30.1°C, and a boiling point of 1983°C!

22. Periodic properties
As the atomic number of an atom increases, atoms display a periodic variation. As a result of this, the physical and chemical behavior of atoms vary.
Periodic law asserts that when elements are arranged by atomic number, their physical and chemical properties vary periodically.

23. Properties of the Atom
We’ll take a look at three properties of the atom:
1. Atomic radius
2. Ionization energy
3. Electron affinity
Items 2 and 3 are important in bonding.

24. 1. Atomic radius
Within each period, the atomic radius decreases with increasing atomic number. The largest atom is an alkali, the smallest is a noble gas.

25. 1. Atomic radius
Within each group, the atomic radius increases in size with an increase in atomic number.

26. 1. Atomic radius
These observations are due to the fact that as the principle quantum number (n) increases, the size of the orbital increases.

27. 1. Atomic radius
It is also due to the fact that as the effective nuclear charge increases, the electrons get pulled inward and this reduces the size of the orbitals.
How does this occur?
The effective nuclear charge is the positive charge that an electron “feels” from the nucleus. It is equal to the nuclear charge but reduced by any shielding from any intervening electrons.

28. Effective Nuclear Charge
The effective nuclear charge on Li (1s22s1) is 3e. The effective nuclear charge on the 2s electron is reduced to the 2 intervening electrons located between the nucleus and the 2s electron.

29. Atomic properties
Consider a given period of elements. The princple quantum number (n) of all outer orbitals remains constant.
The effective nuclear charge increases because the number of core electrons remains constant (while the number of protons increases).
Thus, the size of the outermost orbital, along with the radius of the atom decreases with increasing atomic number (z) in any period.

30. 2. Ionization energy
The first ionization energy of an atom is the minimum energy needed to remove the outermost electrons from a neutral atom in the gaseous state.
The unit is usually energy/mol (eg., Li = 520kJ/mol)
The values tend to increase with increasing atomic number (as one goes from left to right on the periodic table). The atoms/elements with the highest ionization energies are the noble gases—the atoms that are the most stable.

31. 2. Ionization energy
As you go across the table, the atomic radius decreases, the electrons are closer to the nucleus due to the increased nuclear charge, and hence the electrons are harder to get away.

32. 2. Ionization energy
As each electron is removed, more energy is required to remove an additional electron.
Consider Li again:
1st ionization energy = 520kJ/mol
2nd ionization energy = 7,298kJ/mol
3rd ionization energy = 11,815kJ/mol
Valence electrons are easier to remove than are core electrons.

33. 2. Ionization energy

34. 3. Electron affinity
Electron affinity refers to the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion.
For instance, going from Cl to Cl-
If the ion formed is stable, its energy change for its formation will be negative.
The larger the negative value, the more stable the ion.
The halogens form stable negative ions.

35. 3. Electron affinity
The addition of the electron and the stability of the ion is dependent upon which orbital the electron goes into.

36. 3. Electron affinity
The stability of the ion formed when adding an electron to an atom depends on which orbital the electron is going into.
Adding an e- to Group IA forms a fairly stable ion.

37. 3. Electron affinity
The stability of the ion formed when adding an electron to an atom depends on which orbital the electron is going into.
Adding an e- to a Group IIA does not happen because we would have to add an e- to the next higher energy level (a p-orbital).
For the same reason, Group VIIIA atoms do not form stable ions.

38. 3. Electron affinity
Generally, from IIIA to VIIA, more stable bonds form due to the more negative electron affinity values.

39. periodicity
The variations in metallic and nonmetallic character are attributed to variations in ionization energies.
Elements with low ionization energies tend to be metallic; those with high ionization energies non-metallic.
The basic-acidic behavior of oxides of these elements tends to be a good inidcator of the metallic-nonmetallic character of the elements.

40. Properites of Atoms
Oxides are classified as basic or acidic, depending on their reactions with acids and bases.
A basic oxide reacts with acids. Most metal oxides are basic.
An acidic oxide reacts with bases. Most nonmetal oxides are acidic.
An amphoteric oxide is one which has both acidic and basic properties.

41. Group IA oxides All alkali metals form basic oxides.
There is a slight increase in metallic character as we go down a column.

42. Group IIA oxides Alkaline metals also form basic oxides.
There is a slight increase in metallic character as we go down a column.

43. Group IIIA
There is a significant increase in metallic character as we go down a column. The oxides start off as acidic, but become amphoteric—indicative of an increase in metallic character.

44. Group IVA
These elements start out as a nonmetal, become a metalloid, and finish as metals.
The oxides start out as acidic, and finish as amphoteric.

45. Group Va
This group of elements also shows a transition from nonmetal to metalloid to metal. Additionally, this group’s oxides go from acidic to amphoteric to basic.

46. Group VIA
Again, these elements start out as a nonmetal, become a metalloid, and finish as metals. Their oxides start out as acidic and become amphoteric.

47. Group VIIA
These elements are the halogens, are gases, and form acidic oxides.

48. Group VIIIA
These elements are generally unreactive due to their full valence shell.