1. The Periodic Table
chapter 6

2. Developing the Periodic Table
In the early 1800s, scientists began to find ways to classify the elements.
German chemist Dobereiner grouped elements based on similar properties.
English chemist Newlands, arranged elements based on atomic mass.

3. Developing the Periodic Table
Russian chemist, Dmitri Mendeleev organized elements into a table based on atomic mass and similar properties.
Mendeleev stated that the properties of elements are a periodic function of their atomic masses.

4. Mendeleev’s Periodic Table

5. Mendeleev’s Prediction
Mendeleev’s table had several missing elements. When these elements were discovered, they were almost exactly as Mendeleev predicted.
The following is an example of the element we know as Germanium.

6. Germanium is located below silicon
Germanium is located below silicon. Mendeleev predicted its properties based on this location in his table.
Ekasilicon (Es)
Germanium (Ge)
1. Atomic mass: 72
1. Atomic mass: 72.61
2. High melting pt.
2. Melting pt: 945° C
3. Density: 5.5g/cm3
3. Density: 5.323g/cm3
4. Dark gray metal
4. Gray metal
5. Will obtain from K2EsF6
5. Obtain from K2GeF6
6. Will form EsO2
6. Forms oxide (GeO2)

7. Modern Periodic Law Henry Moseley revised Mendeleev’s periodic law.
He used atomic number to organize elements.
Atomic number is the basis for our current periodic law.

8. Periodic Table

9. Periodic Table Review:
Rows on the periodic table are called PERIODS
Columns on the periodic table are called GROUPS or FAMILIES

10. Periodic Table Review There are 7 periods and 18 groups.
Electron arrangements are repeated in periods.
Elements with similar e- configurations are placed in the same group.
Elements in groups are also listed in order of their increasing principal quantum numbers.

11. Electron Configuration
Sublevel / e- capacity s p d f

12. S - block Contains elements in Group 1, Group 2, and He from Group 18.
Electrons are added to the s – orbitals.
EX: H = 1s1
He = 1s2
Li = 1s22s1
Be = 1s22s2

14. P - block
Contains elements in Group 13, Group 14, Group 15, Group 16, Group 17, and the remaining elements from Group 18 (except He)
Electrons are added to the p – orbitals.
Ex: B = 1s22s22p1
C = 1s22s22p2
N = 1s22s22p3

16. D – block Contains elements from the center of the periodic table.
These elements are called transition metals.
Electrons are added to the d – orbitals of the transitions metals as well as La and Ac of the inner transition elements (rare earth).

18. F - block
Contains elements from the inner transition metals (rare earth elements)
Electrons are added to the f – orbitals.
Ex: Ce  Lu
Th  Lr

19. Octet Rule Atoms with full outer levels are stable (less reactive)
For elements (except He) this stable configuration would have eight e-.
(two in the outer s sublevels and six in the outer p sublevels)
These outer eight e- (valence electrons) are called an octet.

20. Octet Rule
Eight electrons in an outer level render an atom unreactive.
This is referred to as the Octet Rule.
When atoms react with one another, they do so to obtain a stable config.
Some atoms gain or lose e- (ions) and some share e- (molecules).

21. Organizing Information on the Periodic Table
Use a pen to label the following:
Group 1 Alkali metals
Group 2 Alkaline earth metals
Group 16 Chalcogens
Group 17 Halogens
Group 18 Noble gases
Sc – Uub Transition metals
La – Lu Lanthanoids
Ac – Lr Actinoids

22. Organizing Information on the Periodic Table
Draw a stair step dark line starting between B and Al.
Label the right side: metals
Label the left side: nonmetals
Write METALLOID along stair step line.
Label the valence e- (outer electrons).
Use colored pencils to shade each group or category a different color.

23. Basic Properties of Metals, Nonmetals, and Metalloids
1. Dense and shiny (luster).
2. Conduct heat and electricity well.
3. Have high melting points.
4.Malleable and ductile.

24. Nonmetals:
1. Generally gases or brittle solids.
2. If solid, dull surface.
3. Good insulators.

25. Metalloids:
1. Properties of both metals and nonmetals.
EX: Silicon, for example, possesses a metallic luster, yet it is an inefficient conductor and is brittle.

26. Properties of Alkali Metals
Group 1 metals
Soft silver metals.
Less dense than other metals and lower melting points.
Very reactive due to large size and one loosely held valence electron.
Too reactive to be found free in nature.

27. Properties of Alkaline Earth Metals
Group 2 Metals
Shiny silvery-white metals
Have 2 valence electrons
Not as reactive as alkali metals but very reactive
All found in the Earth’s crust in mineral form
Too reactive to be found in free element form

28. Properties of Halogens
Group 17 nonmetals
All diatomic gases at room temperature EX: F2, Br2
Too reactive to be found as free elements in nature
Most important group to be used in industry

29. Properties of Chalcogens
Group 16 nonmetals
Diverse group that includes nonmetals, metalloids, and metals

30. Properties of Noble Gases
Group 18 nonmetals
Complete octet of valence electrons
Largely unreactive
Monotomic gases

31. Periodic Trends

32. Using the Periodic Table to Predict Properties of Elements
The basis of the periodic table is the atomic structures of the elements.
Position on the table and properties of these elements arise from the e- configurations of the atoms.
Properties such as density, atomic radius, oxidation numbers, ionization energy, and e- affinity can be predicted.

33. Atomic Radius
As principal quantum number increases, the size of the electron cloud increases. Size of atoms increase moving down Per. Table.
Atoms in the same period have the same quantum number; however, positive charge on the nucleus increases by one proton for each element in a period. This pulls the e- cloud in tighter, decreasing atomic radius.

34. Predicting Atomic Radius
General rule: atomic size increases as you move diagonally from top right corner to bottom left corner.

35. When graphed, atomic radii demonstrates a periodic trend

37. Radii of ions: Ions are atoms that have gained or lost e- from the outer orbitals.
Cations: (+)
Become smaller
1. Positive charged nucleus attracting fewer e-.
2. Reduced the number of energy levels.
EX:
Anions: (-)
Become larger
1. Positive charged nucleus attracting more e- expands e- cloud.

39. Trends in Oxidation Numbers
Our knowledge of e- configurations and the stability of noble gases allows us to predict oxidation numbers for elements.
Oxidation numbers represent the charge an ion obtains after losing or gaining valence electrons.

40. 1+
2+ or 4+ 2+
Tend to have more than one oxidation number 3+ 3- 2- 1- 3+
3+ or 4+

41. Two hydrogen atoms are walking down the road
Two hydrogen atoms are walking down the road. One said, “I think I lost an electron!”.
“Really”, the other replied, “ Are you sure?”.
“Yes, I’m positive”.

42. Ionization Energy The energy required to remove an e- from an atom.
The larger the atom, the less energy is required because the e- are farther from the positive center.
Remove the most loosely held e- is first ionization energy.
Measured in kilojoules per mole
kJ/mol

43. Ionization energy increases diagonally from bottom left corner to top right corner.

45. Classification based on First Ionization Energy
METAL
1. Low 1st ionization energy.
2. Located on left side of Periodic Table.
3. Form positive ions.
NONMETAL
1. High 1st ionization energy.
2. Located on the right side of Periodic Table.
3. Form negative ions.

46. Multiple Ionization Energies
Additional e- can be lost from an atom and the ionization energies can be measured.
IONIZATION ENERGIES (kilojoules per mole)
Element 1st nd rd th th H He Li Be B

47. Electronegativity
Electronegativity is the ability of an atom to capture an electron.
It increases from bottom left to top right corners.

48. Review

49. Review Based on our trends: The most reactive metal element would be
Francium
The most reactive nonmetal element would be
Fluorine

50. Electron Affinity
e- affinity is a measure of an atom’s attraction for an e-.
Metals have low e- affinities.
Nonmetals have high e- affinities.
Chemical reactions occur between atoms with high e- affinity and those with low e- affinity.
EX: Al Br  Al2Br3
(low) (high) (more stable)

51. In Summary
Periodic table is a chart of elements in which the elements are arranged based on their e- configurations which dictates their properties.
Moving down a group in the periodic table, atomic radii becomes larger because more energy levels are needed for more e-.

52. In Summary
As the size becomes larger, the e- are located farther away from the positive center.
This decreases the affinity of that atom to hold on to these outer e-, thus decreasing e- affinity.
Ionization energy is low because it is easy for the atom to lose these outer e-.

53. In Summary
Moving across a period in the periodic table, atomic radii becomes smaller because the energy levels of periods are the same but the positive centers of atoms increase. This pulls the e- cloud closer to the nucleus, making the atom smaller.
Ionization energy and e- affinity increases for these smaller atoms.

54. THE END