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Periodic Table
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The Periodic Table Warmup:
Overview of the Periodic Table
Metals Metalloids Nonmetals Noble Gases
Periodic Table
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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 1869 Russian a) 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
b) The amazing part: he predicted 3 elements not yet
discovered (eka-aluminum, eka-boron, eka-silicon)
Periodic Table
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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
Element and year discovered
Properties
Predicted Properties
Observed Properties
Eka-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
Periodic Table
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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) λ 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) 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)
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. Electron configuration ( reactivity and bonding)
2. Atomic radius
3. Ionic radius
4. Ionization energy
5. Electronegativity
Periodic Table
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1. Electron configuration
1A 1
2A 2
B 3-12
3A 13
4A 14
5A 15
6A 16
7A 17
8A 18
H 1s1
Li 2s1
Na 3s1
K 4s1
Rb 5s1
Cs 6s1
Fr 7s1
1+
He 1s2
Be 2s2
Mg 3s2
Ca 4s2
Sr 5s2
Ba 6s2
Ra 7s2
2+
Sc-Zn 3d
Y-Cd 4d
La-Hg 5d
Ac-Une 6d
(followed
by s)
1+ or 2+
(generally)
B 2s22p1
3s23p1
4s24p1
5s25p1
6s26p1
3+
C 2s22p2
3s23p2
4s24p2
5s25p2
6s26p2
2+
4+
4-
N 2s22p3
3s23p3
4s24p3
5s25p3
6s26p3
3-
O 2s22p4
3s23p4
4s24p4
5s25p4
6s26p4
2-
F 2s22p5
3s23p5
4s24p5
5s25p5
6s26p5
1-
Ne 2s22p6
3s23p6
4s24p6
5s25p6
6s26p6
0
B Lanthanide/
Actinide series
Begin in
period 6
Ce-Lu 4f
Th-Lr 5f
(followed by s)
1+ or 2+
(generally)
The position of a valence electron and the ability to remove it from an atom are related to
• the number of protons in the nucleus
• the extent to which the valence electron is shielded from the positively-charged
nucleus by the negatively-charged core electrons
Noble gas configuration = [core] electrons
“outer” electrons = “valence” electrons Elements of groups 1A-8A have valence
electrons in s and p orbitals.
(Note: USA uses A and B.
Rest of world uses 1-18)
Periodic Table
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Isoelectronic Series
= same number of electrons
(iso = equal)
1. Draw the complete electron configuration of each of the following elements.
2. What ions will they form?
3. When ions, how many electrons does each have? How many protons?
4. Predict the relative diameters of the members of this isoelectronic series.
Element Electron config Ion Ion # e- Ion # p+
O
F
Ne
Na
Mg
Prediction (smallest to largest):
Periodic Table
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2. 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
Periodic Table
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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?
______________________________________________________________________
______________________________________________________________________
Periodic Table
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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: _____________________________
___________________________________________________________________
3. Ionic Radius
a) Cations (+) are always SMALLER than the original atom
(↓ # e-’s ↓ repulsion and greater pull on the valence e-’s by the nucleus
remaining e-’s shrink to nucleus)
b) Anions (-) are always LARGER than the original atom
(↑ # e-’s ↑ repulsion shell swells)
Periodic Table
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Ionic Radius
1. In this table of ionic radii, how is the charge of the ions of elements in groups 1A-4A
related to the group number? _____________________________________________
2. a) Divide the radius of Cs with the radius of its ion: ____________________
b) Divide the radius of Li with the radius of its ion: ____________________
c) Divide the radius of Be with the radius of its ion: ____________________
d) Divide the radius of B with the radius of its ion: _____________________
e) Summarize your findings about a)-d) here: _________________________
___________________________________________________________________
3. a) Divide the radius of the F ion with the radius of the neutral F atom:
b) Divide the radius of the O ion with the radius of the neutral O atom:
c) Divide the radius of the N ion with the radius of the neutral N atom:
d) Summarize your findings about a)-c) here: ___________________________________
_____________________________________________________________________
e) Compare and contrast 2 e) and 3 d) ________________________________________
____________________________________________________________________
Periodic Table
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4. 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 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
Periodic Table
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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?
___________________________________________________________________
___________________________________________________________________
Periodic Table
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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: ↑
Periodic Table
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5. Electronegativity = how much one atom pulls on another atom’s electrons in a bond
∴ only refers to atoms in a bond (molecule or compound)