1. The Periodic Table
Ch 6

2. History of the Periodic Table
Only 13 elements had been discovered by 1700
As time went on and more elements were discovered (5 in the decade of ) a method to organize the elements was needed.
Around the world chemists used different mass values for the elements which made it hard to reproduce experiments.

3. Example: John Newland’s arrangement of the (then known) elements
H 1 Li 2 G 3 Bo 4 C 5 N 6 O 7
F 8 Na 9 Mg 10 Al 11 Si 12 P 13 S 14
Cl 15 K 16 Ca 17 Cr 19 Ti 18 Mn 20 Fe 21
Co & Ni 22 Cu 23 Zn 25 Y 24 In 26 As 27 Se 28
Br 29 Rb 30 Sr 31 Ce & La 33 Zr 32 Di & Mo 34 Ro & Ru 35
Pd 36 Ag 37 Cd 38 U 40 Sn 39 Sb 41 Te 43
I 42 Cs 44 Ba & V 45 Ta 46 W 47 Nb 48 Au 49
Pt & Ir 50 Os 51 Hg 52 Tl 53 Pb 54 Bi 55 Th 56

4. Mendeleev
In 1869 Russian scientist Dimitri Mendeleev, demonstrated a connection between atomic mass and elemental properties.
Using element cards, it took him 18 months to set it up.
Mendeleev developed his table while making a textbook for his students.

5. (continued)
He wrote each of the about 60 elements each on a note card and arranged them in order of increasing atomic mass.
Unlike other scientists, Mendeleev left blanks in his table for elements that weren’t discovered yet.

6. Mendeleev’s periodic table

7. Discovery
After the publication of his Periodic Table missing elements (Ga, Ge, Sc) were found with masses and properties very close to what Mendeleev predicted.
Mendeleev noticed that some elements seemed out of place in his table. He thought the masses (from other scientists) were wrong.

8. Henry Moseley
In 1913, Moseley determined the atomic number for each element.
Elements are now arranged by atomic number, not mass, in the modern periodic table.

9. Modern Periodic Table

10. Periodic Table seven horizontal rows called periods
18 vertical columns called groups or families
groups 1 and 2 (1A and 2A) and groups (3A – 8A) are called representative elements
groups 3-12 are the transition metals

11. Periodic Table
elements in any group have similar physical and chemical properties
properties of elements in periods change from group to group
symbol placed in a square
atomic number above the symbol
atomic mass below the symbol

12. Periodic Table Three main classes of elements: 1. metals 2. nonmetals
3. metalloids

13. Metals Most elements are metals. Properties:
Good conductors of heat and electricity
Solid at room temperature (except Mercury)
Reflect light (shiny)
Lose electrons in reactions

14. Metallic character
Most metallic element always to the left of a period, least metallic to the right, and 1 or 2 metalloids are in the middle
Most metallic element always at the bottom of a column, least metallic on the top, and 1 or 2 metalloids are in the middle of columns 4A, 5A, and 6A

15. Nonmetals Located in the upper right corner of PT
Greater variation among these than metals.
Most are gases at room temperature.
A few are solids (C, S, P) and one is a liquid (Br)
Tend to have properties opposite of metals.
Gain electrons in reactions.

16. Metalloids Generally have properties similar to metals and nonmetals.
An element in this group may behave like a metal under certain conditions, and then behave like a nonmetal under different conditions.
For example, the metalloid Silicon is a poor conductor of electricity, but it becomes a good conductor when it is mixed with another metalloid, Boron.

17. Main groups Group IA  alkali metals Group IIA  alkaline earth metals
Group VIIIA  noble gases
Group VIIA  halogens – “salt formers”
Group VIA  chalcogens
Group VA  Nitrogen group
Group IVA  IVA group
Group IIIA  IIIA group

18. Other groups s & p block filled (A-groups)  representative elements
d block filled (B-groups)  transition metals
f block filled  inner transition metals
4f  lanthanides
5f  actinides
f elements that are naturally occurring  rare earth elements

19. What are periodic trends?
also called “atomic trends” – take place at the atomic level
trends are general patterns or tendencies
they are general not definite – there are exceptions
when looking at trends we look for increases & decreases
across  periodic
down  group

20. Effect on trends Nuclear Charge the “pull” of the nucleus
proportional to the number of protons in an atom
the greater the number of protons, the stronger the nuclear charge (“pull”)
this generally affects periodic trends

21. Effect on trends Shielding
- the electron protection from the nuclear “pull”
- shield = an energy level of electrons
- we are not concerned with single electrons, only energy levels of electrons
- these electrons reduce the nuclear pull
- affects group trends

22. Effect on trends Stability
- where electron arrangement is compared to stable octet (or other special stabilities)
- determines if atom gains or loses electrons
- can be used to explain anomalies in trends

23. Trends in Atomic Size Increases down column Decreases across period
valence shell farther from nucleus because of increased shielding
Decreases across period
left to right because of the nuclear “pull”
adding electrons to same valence shell
valence shell held closer because more protons in nucleus

24. Ionization Energy
Minimum energy needed to remove a valence electron from an atom
1 mole of electrons in the gaseous state (kJ/mol)
The lower the ionization energy, the easier it is to remove the electron
metals have low ionization energies

25. Trends in Ionization Energy
Ionization Energy decreases down the group
valence electron farther from nucleus
Ionization Energy increases across the period
left to right
harder to remove an electron from the atom because of the increased nuclear “pull”
Exceptions: Group 3, Group 6 (chalcogens)

27. Ionization Energy Li + energy  Li+ + e- Li+ + energy  Li+2 + e-
1st ionization = 520 kJ/mol
Li+ + energy  Li+2 + e-
2nd ionization = 7297 kJ/mol
Li+2 + energy  Li+3 + e-
3rd ionization = 11,810 kJ/mol
Notice, each successive ionization energy is greater than the preceding one – there is a greater “pull” between the nucleus and the electron and thus more energy is needed to break the attraction.

28. Examining ionization energies can help you
predict what ions the element will form.
easy to remove an electron from Group IA, but difficult to remove a second electron. So group IA metals form ions with a 1+ charge.

29. Electron Affinity atoms attraction to an electron
it is the energy change that accompanies the addition of an electron to a gaseous atom
“opposite” of ionization energy (Concept NOT actual trend)

30. Exceptions: Nitrogen Group & Noble Gases
Across a Period
electron affinity increases because of increased “pull”
Down a Group
electron affinity decreases because the electrons are shielded from the pull of the nucleus
Exceptions: Nitrogen Group & Noble Gases

31. Electronegativity
the ability of an atom to attract electrons when the atom is in a compound
very similar to electron affinity
Across a Period
increases because of increased pull
Down a Group
decreases because of shielding
F– most electronegative element

32. Decrease
Increase

33. Ionic size cations – lose electrons (positively charged)
anions – gain electrons (negatively charged)
elements gain or lose e- to become stable – being like noble gases (filled outer sublevel)
IA VA - -3
IIA VIA - -2
IIIA VIIA - -1
IVA – share VIIIA – 0, stable

35. Ionic size GOOD RULE OF THUMB
anions are always larger than their neutral atom
cations are always smaller than neutral atom
Across a Period:
cations decrease (I-III) because of greater pull on electrons
anions decrease (V-VII) because of less pull on electrons and repulsion of the electrons
Down a Group:
both cations and anions increase size

36. Reactivity
Reactivity of metals increases to the left on the Period and down in the column
follows ease of losing an electron
Reactivity of nonmetals (excluding the noble gases) increases to the right on the Period and up in the column