PeriodicPeriodic TableTable

Modern Periodic Law

The properties of the elements repeat in a regular pattern when arranged by

their atomic numbers.

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

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

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

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

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

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.

Newlands’ Octaves

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.

Newlands’ importance

• Concept of groups of eight carried over to modern table• Reinforced concept of periodicity from Dobereiner’s table

Mendeleev and Meyer

First useable periodic table (1869)Dmitri Mendeleev 1834 – 1907 Lothar Meyer 1830 – 1895

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.

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

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.

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

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

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

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)

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)

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

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-)

Periodic trends• Radius of isoelectronic ions decreases left to

right• Metals lose electrons and make + ions• Nonmetals gain electrons to make - ions

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

First ionization energy

Electronegativity

• Electronegativity is an atom’s attraction for electrons in a bond

O

H H

• Metals have low electronegativity, nonmetals high

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

Electron affinity

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.

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)

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

Alkali metals in water

http://www.youtube.com/watch?v=QSZ-3wScePM

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

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

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)

Metalloids

• Properties are intermediate between metals and nonmetals

• Poor conductors, semi-shiny solids

• Tend to share electrons rather than transfer

• Used in semiconductors