Download History of the Periodic Table

Periodic Table S1
The Periodic Table
History of the Periodic Table
I.
Early attempts
Made the task a little easier:
Jöns Jakob Berzelius
1828
Swedish
- developed a table of atomic weights
- introduced letters to symbolize elements
a)
Johann Döbereiner
1829
German
- described triads of elements
(e.g. Cl, Br, I; Ca, Ba, Sr; S, Se, Te)
– first indication that elements were related to one another
– atomic mass is related to chemical properties
Karlsruhe Congress (big Chemistry Conference) 1860 Germany
b)
John Newlands
1865
English
- arranged elements in order of relative atomic masses;
- described the Rule of Octaves – every 8th element has similar properties
c)
Julius Lothar Meyer
1870 German
graph of atomic volume (atomic weight/density) against
atomic weight  periodic trends in elements’
properties; established concept of valency
II.
Dmitri Mendeleev
a)
1869
Russian
How:
While writing a book on inorganic chemistry

to get organized, wrote elements on notecards with
some properties and atomic weight/mass: ULTIMATE SOLITAIRE
 arranged elements in order of atomic masses
 noticed a repetition of properties every 8 or 18
elements

elements with similar properties in horizontal rows
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Periodic Table S1
b)
The amazing part: he predicted 3 elements not yet
discovered (eka-aluminum, eka-boron, eka-silicon)
c)
Problems : Ar/K, Te/I, Co/Ni
1st element in each pair has greater atomic mass
 places reactive K in unreactive noble gases
d)
Importance –
1)
realized elements yet to be discovered;
2)
characteristics of element could be predicted from its atomic weight (and
position on the tables)
Properties of Some Elements Predicted by Mendeleev
Predicted
Elements
Eka-aluminum
Eka-boron
Eka-silicon
Element and year
discovered
Gallium, 1875
Scandium, 1877
Germanium, 1886
Properties
Density of metal
Melting point
Predicted
Properties
6.0 g/mL
Low
Observed
Properties
5.96 g/mL
30oC
Oxide formula
Ea2O3
Ga2O3
Density of metal
Oxide formula
3.5 g/mL
Eb2O3
3.86 g/mL
Sc2O3
Solubility of oxide
Dissolves in acid
Dissolves in acid
Melting point
Density of metal
Color of metal
Oxide formula
High
5.5 g/mL
Dark gray
EsO2
900oC
5.47 g/mL
Grayish white
GeO2
Density of oxide
Chloride formula
4.7 g/mL
EsCl4
4.70 g/mL
GeCl4
Discovery of the Noble Gases 1890s
•
Lord Rayleigh (physicist) and Sir William Ramsay (chemist)
•
1894 - Argon “the lazy one”, discovered when Ramsay was trying to isolate nitrogen
•
1895 - Helium – found on earth in uranium minerals (found in the sun in 1868)
•
1898 - Neon “the new one”, Krypton “the hidden one”, Xenon “the alien one”
•
1910 – Radon
Properties:
Largely unreactive, 8 electrons in valence shell, low boiling and melting points
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Periodic Table S1
Nucleus discovered – 1910 - Rutherford predicted that the charge of an
atom is proportional to its mass
III.
Henry Moseley
1913
English
(worked with Rutherford)
a)
 of emitted X-rays corresponded to # protons
 atomic number
“Do other properties match atomic numbers?” Yes!
 arranged the periodic table by atomic #’s, not mass
b)
IV.
Law of Atomic Numbers (Law of Chemical Periodicity)
-
the properties of elements are periodic functions of their atomic numbers
-
corrected incorrect placement of cobalt and nickel, and iodine and tellurium
Glenn Seaborg
1940s
American
1912-1999
a)
“transuranium” elements – formation of elements beyond uranium (93-103)
 reorganization of periodic table to include both series of radioactive elements
(lanthanides and actinides)
b)
note the names of elements 95-103, reflect Seaborg’s academic life – scientists
and institutions (UC-Berkeley)
Trends of the Periodic Table
“periodic” = repeating pattern
Overall theme = electrons’ positions relative to each other and the nucleus determine the
following properties:
1.
Atomic radius
2.
Ionization energy
3.
Electronegativity
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Periodic Table S1
1.
Atomic Radius
½ distance between nuclei
a)
Trend down a GROUP: 
i.
larger atoms – valence e-’s are farther away from nucleus
ii.
shielding effect – the number of e-’s between the nucleus and valence e’s helps keep the valence e-’s farther away from the nucleus, thus  the
pull of the nucleus on the valence e-’s.
b)
Trend across a PERIOD: same principal energy level)
i.
for every added e-, one more p+
 pull on outer e-’s by nucleus
ii.
not as noticeable in periods with heavier elements
(inner e-‘s shield the valence e-’s  greater distance
between nucleus and valence e-’s)
iii.
shielding effect is constant across a period, as e-’s are added only to the
valence, or outermost energy level
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Periodic Table S1
Atomic Radii
1. Which groups and periods of elements are shown in the table of atomic radii?
______________________________________________________________________
2. In what unit is atomic radius measured? __________Express this unit in m __________
3. What are the values of the smallest and largest atomic radii shown? What elements have
these atomic radii? _______________________________________________________
4. What happens to atomic radii within a period as the atomic number increases?
______________________________________________________________________
5. What accounts for the trend in atomic radii within a period?
______________________________________________________________________
______________________________________________________________________
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Periodic Table S1
6. What happens to atomic radii within a group? ____________________________
7. What accounts for the trend in atomic radii within a group?
_______________________________________________________________________
_______________________________________________________________________
8. a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant
figures. Cs:Li _______________
b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant
figures. Cs:Rn ______________
c) Summarize your findings about a) and b) here: _____________________________
___________________________________________________________________
2.
Ionization Energy
Definition: the energy required to remove an electron from an atom in the gas phase
(in J or kJ)
a)
Successive ionization energies for each atom (since > 1
electron can be removed)
Removing each subsequent electron requires more energy
Diagram - removing successive electrons from Be:
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Periodic Table S1
Ionization Energies of Na, Mg, and Al (in kJ/mol)
Successive ionization energies (kJ/mol)
Element
First
Second
Third
Fourth
Na
496
4,562
6,912
9,543
Mg
738
1,451
7,733
10,540
Al
578
1,817
2,745
11,577
1. What happens to the values of the successive ionization energies of an element?
___________________________________________________________________
2. How is a jump in ionization energy related to the valence electrons of the element?
___________________________________________________________________
___________________________________________________________________
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Periodic Table S1
1. What is meant by first ionization energy? _______________________________________
_______________________________________________________________________
2. Which element has the smallest first ionization energy? The largest? What are their
values? ________________________________________________________________
3. What generally happens to the first ionization energy of the elements within a period as
the atomic number of the elements increases? ________________________________
4. What accounts for the general trend in the first ionization energy of the elements within a
period? ________________________________________________________________
______________________________________________________________________
5. Based on the graph, rank the group 2A elements in periods 1-5 in decreasing order of first
ionization energy. _________________________________________________________
8. What generally happens to the first ionization energy of the elements within a group as the
atomic number of the elements increases? _____________________________________
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Periodic Table S1
9. What accounts for the general trend in the first ionization energy of the elements within a
group? _________________________________________________________________
_______________________________________________________________________
b)
Summary of trends in first ionization energies:
trend down a GROUP:  
3.
trend across a PERIOD: 
Electronegativity
= how much one atom pulls on another atom’s electrons in a bond

only refers to atoms in a bond (molecule or compound)
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