1. A German scientist named Johann Dobereiner put forward his law of triads in 1817.
Each of Dobereiner's triads was a group of three elements. The appearance and reactions of the elements in a triad were similar to each other.
At this time, scientists had
begun to find out the relative
atomic masses of the elements.
Dobereiner discovered that the
relative atomic mass of the middle element in each triad was close to the average of the relative atomic masses of the other two elements.
This suggested that atomic mass might be important to arranging the elements.
This gave other scientists a clue that relative atomic masses were important when arranging the elements

2. With this idea in mind nearly 50 years later, in 1864, an Englishman, John Newlands, arranged the known elements into seven rows.
Newlands' Octave Arrangement from Chemical News [1866]
No.
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
The elements were placed in order of increasing atomic mass and arranged so elements with similar chemical properties are in the same group.

3. Since the properties were repeated every eighth element, Newlands referred to his arrangement as the
Unfortunately for Newland’s, there were a number of problems with his arrangement and it really only worked up through calcium.
Law of Octaves.
About five years later, Dmitri Mendeleev a Russian chemist who was unaware of the work of Newlands,
designed his own table.
After about 18 months of gathering information and arranging element cards, it was finished.
Like Newlands, Mendeleev arranged the 63 known elements in order of increasing atomic mass, having elements with similar chemical properties in the same group.

4. Unlike Newlands his table had groups of varying lengths.
He also left gaps in the table for elements he believed had not yet been discovered.
He even made predictions about the properties of some of these elements.

5. Although his successful predictions allowed many to accept his periodic idea,
there were still anomalies that Mendeleev could not explain.
One of these is the order of iodine and tellurium.
The problem was finally solved in 1913 by a 25-year-old English physicist named Henry Moseley.
Moseley showed that the ordering of X-ray spectral lines was dependent upon the ordering of nuclear charge,
that is, in order of the atomic number.
When the elements were placed in order of increasing atomic number,
the anomalies in Mendeleev’s table were eliminated.
Moseley’s work gave rise to the modern periodic law:

6. The properties of the elements are a periodic function of their increasing atomic numbers.
Tragically for the development of science, Moseley was killed in battle during World War I, only two years later.
Moseley’s revised periodic table looked something like this

7. The Modern Periodic Table
seven horizontal rows called periods
18 vertical columns called groups or families
groups 1 and 2 and groups are called representative elements
groups 3-12 are the transition metals

8. Modern Periodic Table cont.
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

9. Metallic Character Metals Metalloids Nonmetals malleable & ductile
shiny, lustrous
conduct heat and electricity
lose electrons in reactions
Metalloids
Also known as semi-metals
Show some metal and some nonmetal properties
Nonmetals
brittle in solid state
dull
electrical and thermal insulators
gain electrons in reactions

10. Metallic Character
Metals are found on the left of the table, nonmetals on the right, and metalloids in between
Most metallic element always to the left of the 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

11. Other Important Groups to Know
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

12. Other Groups s & p block filling  representative elements
d block filling  transition metals
f block filling  inner transition metals
4f  lanthanides
5f  actinides
f elements that are naturally occurring  rare earth elements

13. What are Periodic Trends
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

14. Effects on the 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 has its greatest effect across a period

15. Effects on Trends Cont. 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

16. Effects on Trends Cont. 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

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

19. Trend in 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
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 less than Group 2, Group 6 (chalcogens) less than Group 5

21. Ionization Energy Cont.
Li + energy  Li+ + 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.
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.

22. Ionization Energies for aluminum (kj/mol):
1st - 578
2nd
3rd
4th
5th

23. Electron Affinity atoms tendency to attract (gain) an electron
it is the energy change that accompanies the
addition of an electron to a gaseous atom
Basically the “opposite” of ionization energy
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: 2A; Nitrogen Group & Noble Gases

24. 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
Illustration

25. Ionic Size Cont. Across a Period Down a Group GOOD RULE OF THUMB
cations decrease (I-III) because of greater pull on electrons (protons pull on fewer electrons = more +)
anions decrease (V-VII) because of decrease in electron repulsion (protons pull on fewer electrons i. e. less negative)
Down a Group
both cations and anions increase size
GOOD RULE OF THUMB
anions are always larger than their neutral atom
cations are always smaller than neutral atom

26. Electronegativity
the ability of an atom to attract electrons when the atom is in a compound
electron “tug-of-war”
similar to electron affinity, but not the same
Across a Period
increases because of increased pull
Down a Group
decreases because of shielding
Fluorine – most electronegative element

27. 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
Reactivity Video

28. Practice Which element has a greater ionization energy – Mg or Ba
Which element has a greater atomic radius – N or F
Which element has a greater electron affinity – S or Pb