1. Unit 2 – Electrons and Periodic Behavior
Cartoon courtesy of NearingZero.net

2. Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!

3. The Wave-like Electron
The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.
Louis deBroglie

4. Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum

5. Electron transitions involve jumps of definite amounts of energy.
This produces bands
of light with definite
wavelengths.

6. Quantum Numbers
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
Principal quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number

7. Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
Wolfgang
Pauli

8. Principal Quantum Number
Generally symbolized by n, it denotes the shell (energy level) in which the electron is located.
Number of electrons that can fit in a shell:
2n2

9. Angular Momentum Quantum Number
The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.

10. Magnetic Quantum Number
The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.

11. Assigning the Numbers
The three quantum numbers (n, l, and m) are integers.
The principal quantum number (n) cannot be zero.
n must be 1, 2, 3, etc.
The angular momentum quantum number (l) can be any integer between 0 and n - 1.
For n = 3, l can be either 0, 1, or 2.
The magnetic quantum number (m) can be any integer between -l and +l.
For l = 2, m can be either -2, -1, 0, +1, or +2.

12. Principle, angular momentum, and magnetic quantum numbers: n, l, and ml

13. Spin Quantum Number
Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:

14. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.

15. Sizes of s orbitals Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.

16. The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape

17. P orbital shape There are three dumbbell-shaped p orbitals in
each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.

18. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells”
d orbital shapes
…and a “dumbell
with a donut”!

19. Shape of f orbitals

20. Orbital filling table

21. Electron configuration of the elements of the first three series

22. Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half
fill its 3d sublevel
Copper steals a 4s electron to FILL
its 3d sublevel

23. Mendeleev’s Periodic Table
Dmitri Mendeleev

24. Modern Russian Table

25. Stowe Periodic Table

26. A Spiral Periodic Table

27. “Mayan” Periodic Table

28. Period
The Periodic Table
Group or Family
Group or family
Period

29. The Properties of a Group: the Alkali Metals
Easily lose valence electron
(Reducing agents)
React violently with water
Large hydration energy
React with halogens to form
salts

30. Properties of Metals
Metals are good conductors of heat and electricity
Metals are malleable
Metals are ductile
Metals have high tensile strength
Metals have luster

31. Examples of Metals
Potassium, K reacts with water and must be stored in kerosene
Copper, Cu, is a relatively soft metal, and a very good electrical conductor.
Zinc, Zn, is more stable than potassium
Mercury, Hg, is the only metal that exists as a liquid at room temperature

32. Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.
Nonmetals are poor conductors of heat and
electricity
Nonmetals tend to be brittle
Many nonmetals are gases at room temperature

33. Examples of Nonmetals
Microspheres of phosphorus, P, a reactive nonmetal
Sulfur, S, was once known as “brimstone”
Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

34. Properties of Metalloids
Metalloids straddle the border between metals and nonmetals on the periodic table.
They have properties of both metals and nonmetals.
Metalloids are more brittle than metals, less brittle than most nonmetallic solids
Metalloids are semiconductors of electricity
Some metalloids possess metallic luster

35. Silicon, Si – A Metalloid
Silicon has metallic luster
Silicon is brittle like a nonmetal
Silicon is a semiconductor of electricity
Other metalloids include:
Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te

36. Determination of Atomic Radius:
Half of the distance between nucli in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels

37. Table of Atomic Radii

38. Increases for successive electrons taken from the same atom
Ionization Energy - the energy required to remove an electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
    Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove
Tends to decrease down a group
Outer electrons are farther from the
nucleus

39. Table of 1st Ionization Energies

40. Ionization of Magnesium
Mg kJ  Mg+ + e-
Mg kJ  Mg e-
Mg kJ  Mg e-

41. Another Way to Look at Ionization Energy

42. Electronegativity A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase across
a period
Electronegativities tend to decrease down a
group or remain the same

43. Periodic Table of Electronegativities

44. Summation of Periodic Trends

45. Ionic Radii Cations Anions Positively charged ions formed when
an atom of a metal loses one or
more electrons
Cations
Smaller than the corresponding
atom
Negatively charged ions formed
when nonmetallic atoms gain one
or more electrons
Anions
Larger than the corresponding
atom

46. Table of Ion Sizes