1. The most useful tool in the Lab
Periodic Table
The most useful tool in the Lab

2. Early Organization J.W. Dobereiner (1829) organized elements in triads
Triad – three elements with similar properties (ex: Cl, Br, I)
J.R. Newlands (1864) organized elements in octaves
Octave – repeating group of 8 elements

3. Development of the PeriodiceTable
Dmitri Mendeleev taught chemistry in terms of properties.
Mid 1800’s - molar masses of elements were known.
Wrote down the elements in order of increasing mass.
Found a pattern of repeating properties.

4. Mendeleev’s Table
Grouped elements in columns by similar properties in order of increasing atomic mass.
Found some inconsistencies - felt that the properties were more important than the mass, so switched order.
Also found gaps.
Must be undiscovered elements.
Predicted their properties before they were found.

5. The Modern Periodic Table
Henry Moseley – British physicist
Arranged elements according to increasing atomic number
The arrangement today
Symbol, atomic number & mass

6. The New Way Elements are still grouped by properties.
Similar properties are in the same column.
Order is by increasing atomic number.
Added a column of elements (noble gases)
Weren’t found because they are unreactive.

7. Organization the same # of energy levels and
Horizontal rows = periods
There are 7 periods
Each period represents an energy level
Every element in the same period has
the same # of energy levels and
the same core electron configuration

8. Organization Vertical column = group or family
Similar physical & chemical prop.
Same # of valence electrons
Same common oxidation state
Identified by number & letter

9. Horizontal rows are called periods
There are 7 periods

10. Group 1A are the alkali metals
Group 2A are the alkaline earth metals

11. Group 7A is called the Halogens
Group 8A are the noble gases

12. The group B are called the transition elements
These are called the inner transition elements, and they belong here

13. The elements in the A groups are called the representative elements
outer s or p filling 1A 8A 2A 3A 4A 5A 6A 7A

14. Lanthanides – the 4f orbital fills for these elements

15. Actinide series – the 5f orbitals are being filled for these elements.

16. Types of elements
Metals
Non-metals
Metalloids or semi-metals

17. Metals Good conductor of heat and electricity Malleable Ductile
High tensile strength
High luster
Solid at room temperature
React by losing electrons

18. Nonmetals Poor conductors of heat and electricity
React by gaining electrons
Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)

19. Semi-Metals Heavy, stair-step line Metalloids border the line
Properties intermediate between metals and nonmetals
Learn the general behavior and trends of the elements, instead of memorizing each element property
B, Si, Ge, As, Sb, Te

20. Families Group IA – alkali metals most reactive metals
Silvery in appearance
Soft
Combine easily with non-metals
Melting point is higher than the boiling point of water
Have 1 valence electron

21. Families Group 2 – Alkaline Earth Metal Family
Harder, stronger, denser, higher melting point, and less reactive than alkali
Usually not found as free elements, but as compounds
Have 2 valence electrons

22. Families Group 7 – Halogens Most reactive family Non-metals
Have seven valence electrons
Group 8 – Noble Gas
Inert, unreactive
Have full set of valence electrons

23. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87
1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f 145d106p67s1

24. S- block s1 s2 Alkali metals all end in s1
Alkaline earth metals all end in s2
really should include He, but it fits better later.
He has the properties of the noble gases.

25. 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

26. The P-block p1 p2 p3 p4 p6 p5

27. Each row (or period) is the energy level for s and p orbitals.1 2 3 4 5 6 7
Each row (or period) is the energy level for s and p orbitals.

28. Areas of the periodic table
Group A elements = s & p blocks
representative elements
Wide range of phys & chem prop.

29. Transition Metals -d block

30. d orbitals fill up after previous energy level, so first d is 3d even though it’s in row 4.1 2 3 4 5 6 7 3d

31. F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
inner transition elements f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13

32. 1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f

33. } Atomic Size Radius First problem: Where do you start measuring from?
The electron cloud doesn’t have a definite edge.
Atomic Radius = half the distance between two nuclei of a diatomic molecule.

34. Trends in Atomic Size Influenced by three factors: 1. Energy Level
Higher energy level is further away.
2. Charge on nucleus
More charge pulls electrons in closer.
3. Shielding effect
(blocking effect)

35. WHAT HAPPENS TO ATOMIC RADII?
Does a negative ion (anion) get larger or smaller?
Does a positive ion (cation) get larger or smaller?

36. Trends in Ionic Size Cations form by losing electrons.
Cations are smaller than the atom they come from.
Metals form cations.
Cations of representative elements have noble gas configuration.

37. Ionic size Anions form by gaining electrons.
Anions are bigger than the atom they come from.
Nonmetals form anions.
Anions of representative elements have noble gas configuration.

38. WHAT IS IONIZATION ENERGY?
The energy required to remove an electron
Which element has the highest ionization energy? Why?

39. What determines Ionization Energy?
The greater the nuclear charge, the greater IE.
Greater distance from nucleus decreases IE
All the atoms in the same period have the same energy level.
But, increasing nuclear charge
So IE generally increases from left to right.

40. Ionization Energy
The energy required to remove the first electron is called the first ionization energy
The second ionization energy is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st or 2nd IE.

41. Driving Force
Full Energy Levels require lots of energy to remove their electrons.
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration.

42. WHAT IS ELECTRONEGATIVITY?
The ability of an atom to pull off an electron.
Which element has the highest electronegativity? Why?

43. Periodic Trend Metals are at the left of the table.
They let their electrons go easily
Low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away from others
High electronegativity.

44. Trends in Electron Affinity
The energy change associated with adding an electron to a gaseous atom.
Easiest to add to group 7A.
Gets them to full energy level.
Increase from left to right: atoms become smaller, with greater nuclear charge.
Decrease as we go down a group.