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 à periodic trends in elements’
properties; established concept of valency
II. Dmitri Mendeleev 1869 Russian
a) How:
While writing a book on inorganic chemistry
à to get organized, wrote elements on notecards with
some properties and atomic weight/mass: ULTIMATE SOLITAIRE
à arranged elements in order of atomic masses
à noticed a repetition of properties every 8 or 18
elements
à elements with similar properties in horizontal rows
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
à 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 / Element and year discovered / Properties / Predicted Properties / Observed PropertiesEka-aluminum / Gallium, 1875 / Density of metal / 6.0 g/mL / 5.96 g/mL
Melting point / Low / 30oC
Oxide formula / Ea2O3 / Ga2O3
Eka-boron / Scandium, 1877 / Density of metal / 3.5 g/mL / 3.86 g/mL
Oxide formula / Eb2O3 / Sc2O3
Solubility of oxide / Dissolves in acid / Dissolves in acid
Eka-silicon / Germanium, 1886 / Melting point / High / 900oC
Density of metal / 5.5 g/mL / 5.47 g/mL
Color of metal / Dark gray / Grayish white
Oxide formula / EsO2 / GeO2
Density of oxide / 4.7 g/mL / 4.70 g/mL
Chloride formula / EsCl4 / 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
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) l of emitted X-rays corresponded to # protons
à atomic number
“Do other properties match atomic numbers?” Yes!
à arranged the periodic table by atomic #’s, not mass
b) 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
IV. Glenn Seaborg 1940s American
1912-1999
a) “transuranium” elements – formation of elements beyond uranium (93-103)
à 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
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+
à pull on outer e-’s by nucleus
ii. not as noticeable in periods with heavier elements
(inner e-‘s shield the valence e-’s à 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
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?
______
______
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:
Ionization Energies of Na, Mg, and Al (in kJ/mol)
Successive ionization energies (kJ/mol)
Element / First / Second / Third / FourthNa / 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?
______
______
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? ______
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: ¯ trend across a PERIOD:
3. 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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