Periodic Table Table. Modern Periodic Law The properties of the elements repeat in a regular pattern...
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PeriodicPeriodic TableTable
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Modern Periodic Law
The properties of the elements repeat in a regular pattern when arranged by
their atomic numbers.
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History
• Johann Dobereiner – 1829 (friend of Goethe)
• He was the first to organize elements by their properties
• He grouped them in groups of three called triads
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triads
• He noticed that the atomic mass of the middle member of the group was close to the arithmetic mean of the others.
• Chlorine = 35.5, Bromine = 80, Iodine = 127 (average of Cl and I = 81)
• Properties in common: – All react vigorously with first column metals to form
soluble salts (compounds of a metal and nonmetal)– Hydrogen compounds are strong acids – All form -1 ions
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triads
• Lithium = 7, Sodium = 23, Potassium = 39 (average of Li and K = 23)
• Properties in common: – All salts are soluble– All give brightly colored flames– All react vigorously with water– All form +1 ions
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Other triads
• Calcium = 40, Strontium = 88, Barium = 137 (average of Ca and Ba = 88.5)
- All give +2 ions
• S = 32, Se = 79, Te = 127.6 (average of S and Te = 79.8)
- All give smelly compounds with hydrogen
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Failure of triads
• Not all elements could be fit into triads: iron, manganese, nickel, cobalt, zinc and copper are similar elements but cannot be placed in the triads.
• Newly discovered elements did not fit into triads
• Very dissimilar elements could be fit into triads
• Dobereiner’s triads were discarded
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Newlands’ octaves
• John Newlands 1838 - 1898
• Law of Octaves (1863)• Elements can be
arranged in “octaves” because certain properties repeated every 8th element when the elements are arranged in order of increasing atomic mass.
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Newlands’ Octaves
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Newlands’ octaves
• Newlands’ Octaves also failed– It was not valid for elements that had atomic
masses higher than Ca.– The octaves mixed metals and nonmetals –
for example he put iron (metal) in the same group as oxygen and sulfur (non-metals)
– When more elements were discovered, such as noble gases He, Ne, Ar, they could not be accommodated in his table.
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Newlands’ importance
• Concept of groups of eight carried over to modern table• Reinforced concept of periodicity from Dobereiner’s table
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Mendeleev and Meyer
First useable periodic table (1869)Dmitri Mendeleev 1834 – 1907 Lothar Meyer 1830 – 1895
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Modern Periodic table
• The table was organized by atomic mass (not atomic number) and by properties.
• When organized by atomic mass, both found that the chemical properties repeated on a regular basis – “Periodicity”
• Both scientists noticed holes in the periodic table where elements seemed to be missing.
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Modern Periodic Table
• However, Mendeleev….
….published first (1869, Meyer in 1870)
….corrected the atomic mass of several elements
….classified anomalous elements by properties rather than atomic mass – he said that future measurements would correct anomalous masses
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Modern Periodic table
Ar and KCo and Ni Te and I
Th and Pa….accurately predicted the properties
of missing elements Sc, Ga, and Ge Mendeleev is remembered as the
inventor of the modern periodic table, not Meyer.
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Moseley and Seaborg
• Henry Moseley discovered the proton and atomic number in 1913
• Arranging the periodic table by atomic number eliminated the problem of anomalous atomic weights.
• Glenn Seaborg came up with the idea of the actinide series – last major modification
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Structure of the table
• Rows = periods– All elements in a period have the same
valence shell and the same number of occupied energy levels
• Columns = groups or families– All elements in a group have the same dot
structure– All elements in a group have similar properties
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Coloring time!
• Label the representative elements (s and p blocks)– The number of valence electrons of these
elements increases by one moving left to right
• Label the transition elements (d block)• Label the inner transition elements (f
block)– Transition elements all considered to have
two valence electrons
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More coloring!
• Label the dividing line between metals (on the left) and nonmetals (on the right)
• Label the following groups:• Column 1: Alkali metals (Li to Fr)• Column 2: Alkaline earth metals (Be to Ra)• Representative column 6: Chalcogens (oxygen
family)• Representative column 7: Halogens (fluorine
family)• Representative column 8: Noble gases (include
helium)
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Even more coloring!
• First row of inner transition metals: Lanthanide Series
• Second row of inner transition metals: Actinide Series
• Label the metalloids (B, Si, Ge, As, Sb, Te, Po)
• Label the “other metals” (Al, Ga, In, Sn, Tl, Pb, Bi)
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periodic trends
• Atomic radius decreases across a period
• Result of increasing nuclear charge
• Radius increases down a column
• Valence electrons are in higher energy levels
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Periodic trends
• Ionic radius: ions are atoms that have gained or lost an electron
• Ions have a charge equal to
# protons - # electrons• “Isoelectronic species” are atoms or ions
with the same number of electrons• Na+, F- and Ne are isoelectronic (10 e-)
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Periodic trends• Radius of isoelectronic ions decreases left to
right• Metals lose electrons and make + ions• Nonmetals gain electrons to make - ions
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Ionization energy
• Ionization energy is the energy needed to remove the highest energy electron from an atom (makes a +1 ion)
• Increases across a row due to increased nuclear charge
• Decreases down a column – electrons in higher energy levels are easier to remove, and are shielded by inner shell electrons
• Alkaline earth metals and nitrogen family are slightly higher than expected due to breaking symmetry of half-filled and completely filled shells
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First ionization energy
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Electronegativity
• Electronegativity is an atom’s attraction for electrons in a bond
O
H H
• Metals have low electronegativity, nonmetals high
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Electron affinity
• The energy gained or lost when a gaseous atom of an element gains an electron
• Sometimes defined as the energy required to detach an electron from a -1 charged ion
• Values are generally positive (endothermic process)
• Values generally increase from left to right, with more exceptions than ionization energy
• Values for noble gases are very small or negative
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Electron affinity
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Properties of metals
• Physical properties:• Shiny (Luster)• Flexible (malleability – can be hammered into a
sheet)• Ductility (can be drawn into wire)• Conductors of heat and electricity• Hardness – transition metals are the hardest (Ti,
Cr) though they are less hard than C (diamond) or B. Alkalis are soft; Alkaline earths are hard.
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Physical properties of metals
• Most are solids – only mercury is a liquid• Magnetism
– Diamagnetism: no unpaired electrons, unaffected or repelled by magnet
– Paramagnetism: Unpaired electrons, attracted to magnet – Ferromagnetism: Ability to form a permanent magnet (Fe, Co,
Ni, some inner transitions, some alloys and compounds of these metals)
• Curie temperature: temperature at which a material loses its ferromagnetic properties (1388K for Co, 88K for Dy, 1043K for Fe, 627K for Ni)
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Metals
• Chemical properties:• Tend to lose electrons and form + ions
– The further left on the table, the more readily the metal loses electrons
– Left side of table are better conductors, more malleable, etc.– Charge of ions depends on column; transition metals vary– More reactive metals are at the bottom of the group because of
shielding
• Form salts with non-metals• Many react with acids to give hydrogen gas and a salt
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Alkali metals in water
http://www.youtube.com/watch?v=QSZ-3wScePM
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Transition metals
• All considered to have two valence electrons, though many different valence states (charges on ions) can exist
• Most tend to be hard and dense
• Tc and all metals past Bi are radioactive; many others have radioactive isotopes as well
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Nonmetals
• Physical properties:• Can be solids, liquid (Br only) or gas • Solids are generally hard• Gases are the Noble Gases and the seven
diatomic gases (BrINClHOF: Br2, I2, N2, Cl2, H2, O2, F2)
• Br2 is a volatile liquid, and I2 an easily sublimed solid
• Many are colored (S is yellow, Cl pale green, Br orange, I purple, O pale blue)
• Most are diamagnetic, except oxygen
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Chemical properties of nonmetals
• Nonmetals tend to gain electrons and form negative ions
• Will react with metals to form salts – for example, Fe2O3 (rust)
• When forming compounds with each other, electrons are shared rather than transferred
• Noble gases are monatomic and don’t react with anything except fluorine (only Xe, Kr and Rn)
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Metalloids
• Properties are intermediate between metals and nonmetals
• Poor conductors, semi-shiny solids
• Tend to share electrons rather than transfer
• Used in semiconductors