1. THE PERIODIC TABLE & ELECTRON CONFIGURATION
Chapters 5 & 6

3. The Element Song

4. Dimitri Mendeleev Invented periodic table
Organized elements by properties
Arranged elements by atomic mass
Predicted existence of several unknown elements
Element 101
1860’s proposed new arrangements of elements.
1869 Published original periodic table
Mendeleev missed receiving the Nobel prize in chemistry by just one vote in 1906, and died before the next year’s election.
Element 101 (discovered in 1955) was named Mendelevium in his honor.

5. Mendeleev’s Early Periodic Table DON’T COPY
TABELLE II
GRUPPE I GRUPPE II GRUPPE III GRUPPE IV GRUPPE V GRUPPE VI GRUPPE VII GRUPPE VIII
___ ___ ___ ___
RH RH RH RH
R2O RO R2O RO R2O RO R2O RO4
REIHEN 1 2 3 4 5 6 7 8 9 10 11 12
H = 1
Li = Be = B = C = N = O = F = 19
Na = Mg = Al = Si = P = S = Cl = 35.5
K = Ca = ? = Ti = V = Cr = Mn = Fe = 56, Co = 59,
Ni = 59, Cu = 63
(Cu = 63) Zn = ? = ? = As = Se = Br = 80
Rb = Sr = ? Yt = Zr = Nb = Mo = __ = Ru = 104, Rh = 104,
Pd = 106, Ag = 108
In the 1860’s, Mendeleev and the German chemist Lothar Meyer, each working alone, made an eight-column table of the elements.
However, Mendeleev had to leave some blank spots in order to group all the elements with similar properties in the same column. To explain these blank spots, Mendeleev suggested there must be other elements that had not yet been discovered. On the basis of his arrangement, Mendeleev predicted the properties and atomic masses of several elements that were unknown at the time.
Mendeleev left blanks in his table for undiscovered elements.
Mendeleev predicted properties and masses of unknown elements correctly
(Ag = 108) Cd = In = Sn = Sb = Te = J = 127
Cs = Ba = ? Di = ?Ce = __ __ __ __ __ __ __
( __ ) __ __ __ __ __ __
__ __ ? Er = ? La = Ta = W = __ Os = 195, Ir = 197,
Pt = 198, Au = 199
(Au = 199) Hg = Tl= Pb = Bi = __ __
__ __ __ Th = __ U = __ __ __ __ __
From Annalen der Chemie und Pharmacie, VIII, Supplementary Volume for 1872, p. 151.

6. Examples of Mendeleev’s DON’T COPY
Mendeleev placed certain elements out of order-he assumed that errors had been made in the atomic masses, but it is clear that some elements remain out of order
Moseley changed that with x-ray spectra
Other additions to the table included the Noble gases discoved by William Ramsey

7. Modern Periodic Table Henry G.J. Moseley
Determined the atomic numbers of elements from their X-ray spectra (1914)
Arranged elements by increasing atomic number
H.G.J. Moseley ( ) while doing post-doctoral work (with Ernest Rutherford) bombarded X-rays at atoms in increasing number and noted that the nuclear charge increased by 1 for each element. This gave him the idea to organize the elements by increasing atomic number
Periodic law – elements organized by increasing atomic number on periodic table (1913)
In 1913, Moseley analyzed the frequencies of X -rays emitted by the elements and discovered that the underlying foundation of the order of the elements was atomic number, not atomic mass
Moseley hypothesized that the placement of each element in his series corresponded to its atomic number Z, which is the number of positive charges (protons) in its nucleus
Moseley proved the periodic law in the region from Z = 13 to 79, and that there could be NO other elements in this region

8. Modern Periodic Table Elements are arranged in 7 horizontal rows,
in order of increasing atomic number from LR and from top bottom
Rows are called periods
Elements with similar chemical properties form vertical columns, called groups,
Groups 1, 2, and 13 through 18 are the main group elements
transition elements: groups 3 through 12 are in the middle of the periodic table
Inner transition elements: The two rows of 14 elements at the bottom of the periodic are the lanthanides and actinides
Periods: are numbered from 1 to 7
Groups: which are numbered from 1 to 18

9. Group 2 = Alkaline Earth Metals
Groups to Know
Group 1 = Alkali Metals
Group 2 = Alkaline Earth Metals
Group 17 = Halogens
Group 18 = Noble Gases

10. World of Chemistry

11. IONS
Positive and negative ions form when electrons are transferred between atoms
Cation: an ion with a + charge
Example: Na+ Ca2+
Anion: an ion with a – charge
Example: O F-

12. ELECTRONEGATIVITY
Electronegativity describes how electrons are shared in a compound
The high number means the element has a greater pull on electrons
Fluorine is the most electronegative element

13. SUMMARY OF PERIODIC TRENDS DON’T COPY
Figure 6.22

14. Light, Energy, and Electrons
e-s are arranged in energy levels (e.l.’s), at different distances from nucleus
Close to nucleus = low energy
Far from nucleus = high energy

15. Rules for “placing” e-s in energy levels
e-s in highest occupied level are “valence e-s”
Only so many e-’s can fit in a particular e.l.
e-s fill lower e.l.’s before being located in higher e.l.’s*
Ground state is the lowest energy arrangement of e-s.
* There are exceptions we will learn later!)

16. Light, Energy, and Electrons
e-s can jump to higher energy levels if they absorb energy.
They can’t keep the energy so they lose it and “fall back” to lower levels.
When they do this, they release the energy they absorbed in the form of light.

17. Light, Energy, and Electrons
(See p 129 of text ChemI/IH)
Electron energy levels are like rungs of a ladder.
Ladder
To climb to a higher level, you can’t put your foot at any level,
you must place it on a rung
Electron energy levels
e-s must also move to higher or lower e.l.’s in specific intervals

18. Bohr Model of the Atom (don’t copy this slide)
Interactive Bohr Model

19. Light, Energy, and Electrons
Quantum-the amount of energy required to move an electron from one E.L. to another.

20. Atomic Emission Spectrum (A.E.S)
Each element emits a color when its excited e-s “fall back.”
Pass this light thru a prism, it separates into specific lines of color.
You can identify an element by its emission spectrum! (no 2 elements have the same AES)

21. Emission Spectra of H, He, Ne (don’t copy this slide)

22. Use of e- waves (don’t copy this slide)
Electron microscope magnifies tiny objects b/c e- wavelength much smaller than visible light
snowflake

23. Heisenburg Uncertainty Principle DON’T COPY
Def: if you want to locate something, you can shine light on it
When you do this to an electron, the photons send the e- off in an unpredictable direction
(def):Therefore, you can never know BOTH the position and velocity of an e- at the same time

24. Electron Sublevels
Each electron has an “address,” where it can be considered to be located in the atom.
Main energy level (principal quantum #)
= “hotel”
Sublevel = “floor”
Orbital = “room”
Regions of space outside the nucleus
All orbitals in a sublevel have the same energy
2 electrons max can fit in an orbital

25. Sublevels in Atoms See Fig 7.5 on p 235 Main energy level
Types of sublevels
# of orbitals
# of electrons 1 s 2 p 3
(4 total) d 5
(9 total)
4-7 f 7
(16 total)
See Fig 7.5 on p 235

26. Orbitals s orbitals are spherical p orbitals are dumbbell shaped
There is only 1 orbital
p orbitals are dumbbell shaped
There are 3 orbitals, all with = energy
Each is oriented on either x, y, or z axis
They overlap
d orbitals have varying shapes
There are 5 orbitals, all with = energy
f orbitals have varying shapes
There are 7 orbitals, all with = energy

27. The Periodic Table & Atomic Structure
Shape of p. table is based on the order in which sublevels are filled
REGIONS OF THE P. TABLE (see p 244 of book)
s REGION (“block”) - Groups 1 & 2
p REGION (block) - Groups 13-18
d REGION (block)- Groups 3-12 (Transition Elements)
f REGION (block)- (Inner Transition Elements)

28. Regions or “Blocks” of the P. Table (don’t need to copy)

29. Writing e- Configurations for Elements Using the P. Table
1. Always start with Period 1-go from L to R.
2. Go to Period 2-from L to R
3. Go to Period 3- from L to R
4. Continue w/Periods #4-7, L to R, until you arrive at the element you are writing e- configuration for.
Exception: elements in d block are 1 main E.L lower than the period where they are located
Exception: elements in f block are 2 main E.L.s lower than the period where they are located

30. Correct Order of Sublevels (lowest to highest energy)
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

31. e- configurations
1. Use the P. Table to write the sublevels in increasing order. 2. Add a superscript next to each sublevel that shows how many e-s are in the sublevel 3. Ex: Hydrogen: 1s1 Helium: 1s2 Lithium: 1s22s1 Oxygen: 1s22s22p4

32. Why are d & f block elements’ sublevels out of order?
When you get to the higher main E.L.’s, the sublevels begin to overlap.

33. Chromium is actually… (copy this!)
1s22s22p63s23p63d54s1
3d54s1Instead of 4s2 3d4
There is less repulsion (lower energy) in the 2nd arrangement 4s 3d

34. Noble Gas Notation Short-cut way of showing e- configuration
A Noble Gas is a Group 18 element.
Identify the noble gas in the period above your element of interest. Write this symbol in brackets.
Write the e- configuration for any additional e-s that your element of interest has, but the noble gas doesn’t have.
Ex: Nitrogen: 1s22s22p5 becomes [He] 2s22p5

35. Arrow Orbital Diagram- Used to show e- configuration.
SYMBOLS:
A box represents an orbital
Label each box with the sublevel :1s 2s 2p p 2p
An arrow represents an electron
2 arrows (e-s) in the same orbital face opposite directions.
Example: oxygen, see above
↑ ↓
↑ ↓
↑ ↓ ↑ ↑

36. Arrow Orbital Diagram- Used to show e- configuration, cont.
INSTRUCTIONS:
Fill electrons from lowest to highest sublevel.
Never place 2 e-s in the same orbital of a sublevel until you have placed one in each of the orbitals (Hund’s Rule)