1. Atomic Structure and Periodicity
Chapter 7
Atomic Structure and Periodicity

2. Electromagnetic Radiation
Energy travels in waves that travel at the speed of light
 (lambda) = distance between two peaks/troughs in a wave *unit is METERS (or a form of meters)
 (nu) = number of waves (cycles) per second (unit is hertz)

4. Equation  = c c is the speed of light (2.9979 X 108 m/s)
Example: Photosynthesis uses light with a frequency of 4.54 X 1014 s-1 (Hz). What wavelength does this correspond to?
Answer: 6.60 X 10-7 m = 660nm

5. Max Planck
Postulated that energy can be gained or lost only by whole number multiples of h
h = Planck’s constant (6.626 X J/s)
 = frequency of the radiation being absorbed/given off (Hz or s-1)
∆E = nh
∆E = change in energy of a system
n = whole number of photons (usually 1)
“Packets” of energy with specific units of size can be transferred (called a quantum).

6. Example
Sodium atoms have a characteristic yellow color when excited in a flame. The color comes from the emission of light of nm.
What is the frequency?
Answer: X 1014s-1 (Hz)
What is the change in energy associated with this photon? Per mole of photons?
Answer: X J/photon = X 105 J/mole

7. Quantum Mechanical Model
Heisenberg, de Broglie, Schrödinger
Shows probability distribution: the darker the color (“cloud”), the more likely an electron can be found in a given area surrounding an atom
“A three-dimensional electron density map”

8. Quantum Numbers Principal Quantum Number: n
Angular Momentum Quantum Number: l
Magnetic Quantum Number: ml
VIDEO?

9. Principal Quantum Number (n)
Whole number values (1, 2, 3, 4…)
SIZE/ENERGY
Higher “n” = larger orbital, electron spends more time farther from nucleus
Higher energy b/c e- is less tightly bound to nucleus

10. Angular Momentum Quantum Number (l) <- cursive l
Values 0  n-1
0=s (spherical), 1=p (dumbell), 2=d (clover…exception), 3=f (complex…see book), 4=g (ummm?)
Deals with shape of the atomic orbital

11. Magnetic Quantum Number (ml)
Values –l  +l  those are cursive L’s
Relates to the orientation of the orbital in space

12. Summary…
n = 1, 2, 3, ….
l = 0, 1, …(n-1)
ml = -l, 0, +l (those are L’s)

13. Electron Spin Electrons can have one of two spin states (orientations)
Has a quantum number associated with it (ms) which can only be +1/2 or -1/2 (it spins one of two ways)
Leads to Pauli exclusion principle: NO TWO ELECTRONS CAN HAVE THE SAME SET OF FOUR QUANTUM NUMBERS (n, l, ml, and ms)

14. Polyelectronic Atoms Atoms with more than one electron
Electrons are attracted to positive nuclei, but repulsed by other electrons surrounding the atom
Called screening/shielding
Electrons prefer to fill orbitals in order (s, p, d, then f)

15. Arrangement of the PT
Dobereiner: groups of three with similar properties called triads
Meyer and Mendeleev: modern PT
Still can be used to predict properties of elements
Was listed in order of atomic MASS
NOW: listed in order of atomic NUMBER!

16. Aufbau Principle
As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogenlike orbitals.
Different atoms have different number of protons, and thus different numbers of electrons.

17. Hund’s Rule
Unpaired electrons are represented as having parallel spins (with spin “up”)
Electrons like to be alone if they can

18. Examples Cool Animation
Use boxes (orbital diagram) to represent the electrons in the following atoms: Be N F

19. Electron Configuration

20. Examples Boron Sulfur Calcium Iron 1s22s22p63s23p64s23d6
Answer:1s22s22p1
Sulfur
Answer: 1s22s22p63s23p4
Calcium
1s22s22p63s23p64s2
Iron
1s22s22p63s23p64s23d6

21. Shorthand Version Se Sr Al I
“Kernel Notation” uses the noble gas in brackets prior to the atom needed… Se
[Ar]4s23d104p4 Sr
[Kr]5s2 Al
[Ne]3s23p1 I
[Kr]5s24d105p5

22. Valence vs. Core Electrons
Using electron configuration, you can see that every group on the PT has the same valence electron configuration
Similar chemical behavior

23. Exceptions to Memorize
Chromium: [Ar]4s13d5 NOT [Ar]4s23d4
Both s and d are half filled
Copper: [Ar]4s13d10
There are more, but these (besides Cr and Cu) do not need to be memorized.

24. Periodic Trends Ionization Energy (first and second) Electron Affinity
Atomic Radius
Alkali Metals

25. Ionization Energy Energy required to remove an electron
Measured in joules
Atoms that generally form positive ions (metals) have higher ionization energies
First ionization energy = first electron
Second ionization energy = second electron
TREND: increases L to R and bottom to top (largest is upper right)

26. Electron Affinity
Energy change with adding an electron to a gaseous atom
TREND: generally increases L to R and bottom to top (largest is upper right)

27. Atomic Radius
TREND: increases R to L and top to bottom (largest lower left)

28. In General… Each group…
Has the same number of valence electrons which determine an atom’s chemistry
Groups from figure 7.36 on page 324 should be memorized
Most reactive metals = lower left (smallest ionization energies)
Most reactive nonmetals = upper right (highest electron affinities)

29. Figure 7.36

30. More trends… Density increases going down a group
In redox reactions between SOLID metals and nonmetals…
Nonmetal = oxidizing agent
Metal = reducing agent
Going down the group = better reducing agent (oxidation = lose electrons)