THE PERIODIC TABLE HISTORY & STRUCTURE

Historical Background Classification of the elements began in 1817 with J.W. Dobereimer. Grouped the elements in groups of 3 by atomic number. Problem: did not leave room or system for the addition of new elements.

And then….. 1866: John Newlands (musician) grouped the elements by atomic mass in groups of 8 Problem: Still no room for new elements, and atomic mass was a problem when you consider isotopes.

But then 1869: Dmitri Mendeleev (Russia) and Julius Meyer (Germany) Grouped elements in sequence from left to right and top to bottom by increasing atomic mass and reaction with chlorine and oxygen Gave us the first significant organization of elements in that new elements could be added by their reactivity

But then: In 1905, Henry Mosely discovered that the elements could be grouped according to their energy relationships in order of increasing atomic number THIS LED TO……….

THE PERIODIC LAW: The physical and chemical properties of elements are the periodic functions of their atomic number

Elements are arranged in GROUPS & PERIODS Vertical columns of elements= GROUPS similar chemical properties same number of electrons in their outermost shell (valence shell)

Horizontal Rows of elements = PERIODS No similar chemical properties Members have the same number of electron shells

Elements may be categorized as: METALS Chemical Properties Form + charged ions by losing electrons Have 3 or less electrons in their valence shells Physical Properties Good conductors of heat and electricity High density Luster Malleable and ductile

NONMETALS Chemical Properties Physical Properties Form – ions by gaining electrons (except 8A) Have 5-8 electrons in the valence shell Physical Properties Poor conductors of heat and electricity Solids tend to be brittle and non-ductile Low density

METALLOIDS Have characteristics of metals and nonmetals Traditionally used as semiconductors

Some groups have special names and characteristics Group 1A: The Alkali Metals 1 electron in valence shell Form +1 ions Highly reactive with oxygen and water Mixed in compounds to form alkaline mixtures

Group 2A: The Alkali Earth Metals 2 electrons in valence shell Form +2 ions Reactive Form alkaline solutions in water Fire resistant (do not melt in fire)

Group 6A: The Chalcogens 6 electrons in valence shell All found in copper ores (named from Greek word chalcos, meaning ore) Show increasing metallic properties as atomic number increases. Form -2 ions

Group 7A: The Halogens Very active because of need to fill valence shell Form - 1 ions 7 electrons in valence shell Tend to form salts with elements from 1A and 2A

Group 8A: The Noble Gases 8 electrons in valence shell “filled” Inert gases (stable and unreactive)

Arrangement of electrons determines chemical properties Model of arrangement is called Noble Gas Shell Model Shell: region of space about the nucleus where electrons reside Each shell represents collection of probability clouds Shells are 3-D spheres Each shell holds max number of electrons filling from the inside to the outside 1st shell: 2 electrons 2 nd shell: 8 electrons 3 rd shell: 18 electrons Electrons fill shells in one direction until ½ filled, then pair in the opposite direction

Effective Nuclear Charge: Z* Effective nuclear (Z*) charge is the essentially the positive charge from the nucleus that an outer electron “sees.” The greater the Z*, the more attracted an electron is to the nucleus. Outer shells have diminished nuclear charges Effective nuclear charge diminishes with distance from the nucleus Larger the atom, lower the effective nuclear charge The less filled the valence shell, the lower the effective nuclear charge Effective nuclear charge increases from the lower left hand corner to the upper right hand corner of the periodic table

Higher the Z* = Higher Ionization Energy Relates to an atom’s ability to lose electrons Represents the amount of energy needed to remove an electron from the valence shell Higher the Z* = Higher Ionization Energy Higher Ionization Energy = greater tendency to gain electrons Lower Ionization Energy = greater tendency to lose electrons

Electron Affinity Relates to the ease by which atoms gain electrons Atoms on the upper right side of the periodic table have a greater electron affinity Exception: Noble Gases have no electron affinity