The History of the Modern Periodic Table

Periodic Table Geography

The horizontal rows of the periodic table are called PERIODS.

The elements in any group of the periodic table have similar physical and chemical properties! The vertical columns of the periodic table are called GROUPS, or FAMILIES.

Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

Alkali Metals

Alkaline Earth Metals

Transition Metals

InnerTransition Metals These elements are also called the rare-earth elements. InnerTransition Metals

Halogens

Noble Gases

The s and p block elements are called REPRESENTATIVE ELEMENTS.

During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table.

Johann Dobereiner Model of triads 1780 - 1849 In 1829, he classified some elements into groups of three, which he called triads. The elements in a triad had similar chemical properties and orderly physical properties. (ex. Cl, Br, I and Ca, Sr, Ba) Model of triads 1780 - 1849

John Newlands Law of Octaves 1838 - 1898 In 1863, he suggested that elements be arranged in “octaves” because he noticed (after arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th element. Law of Octaves 1838 - 1898

John Newlands 1838 - 1898 Law of Octaves Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society. 1838 - 1898 Law of Octaves

John Newlands 1838 - 1898 Law of Octaves His law of octaves failed beyond the element calcium. WHY? Would his law of octaves work today with the first 20 elements? 1838 - 1898 Law of Octaves

Dmitri Mendeleev In 1869 he published a table of the elements organized by increasing atomic mass. 1834 - 1907

Lothar Meyer At the same time, he published his own table of the elements organized by increasing atomic mass. 1830 - 1895

Elements known at this time

Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass. Both left vacant spaces where unknown elements should fit. So why is Mendeleev called the “father of the modern periodic table” and not Meyer, or both?

Mendeleev... stated that if the atomic weight of an element caused it to be placed in the wrong group, then the weight must be wrong. (He corrected the atomic masses of Be, In, and U) was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown.

After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions for Sc, Ga, and Ge were amazingly close to the actual values, his table was generally accepted.

However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights determined. By looking at our modern periodic table, can you identify what problems might have caused chemists a headache? Ar and K Co and Ni Te and I Th and Pa

Henry Moseley In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number. *“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.” 1887 - 1915

Henry Moseley His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.

Glenn T. Seaborg After co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became known as the Actinide series. 1912 - 1999

Glenn T. Seaborg He is the only person to have an element named after him while still alive. "This is the greatest honor ever bestowed upon me - even better, I think, than winning the Nobel Prize." 1912 - 1999

Modern Periodic Table

The “Real” Periodic Table

Modern Periodic Table

Groups

Unstable!!! Alkali Metals (Group 1) Brainiac Video most reactive metals never found alone in nature lose 1 e- to achieve stability Brainiac Video

Alkaline Earth Metals (Group 2) harder, denser, stronger, higher melting points than Group 1 less reactive than Group 1 will lose 2 e- to achieve stability

Transition Metals (Groups 3-12) metals with 1 or 2 e- in their outer energy level less reactive than Group 1 & 2 some exist in nature as free elements (Au, Pd, Pt, Ag) lose e- to achieve stability

Inner Transition Metals Lanthanides—elements 57-70 all very similar to each other this is why they are not placed in a group all are metals, some are very strong magnets Actinides—elements 89-102 all are radioactive most are synthetic all are metals

CARBON can lose, gain, or share 4 e- Prefers to share. Why? Element of life

Unstable!!! Halogens (Group 17) most reactive nonmetals never found alone in nature gain or share e- to achieve stability Chlorine Iodine Bromine

Noble Gases (Group 18) Discovered between 1858 and 1900 inert Newest Group

HYDROGEN (H) doesn’t belong to any group unique properties nonmetal placed where it is based on e- config. can lose, gain, or share its 1 e-

Metals, Nonmetals, & Metalloids

Metals hard, shiny (lustrous), good conductor, loses e- to form compounds, usually 2 or less e- in the outer energy level malleable – hammered into sheets ductile – drawn into wires

Nonmetals gases or brittle solids insulators, gain or share e- to form compounds usually have 5 or more e- in outer energy level

Metalloids semiconductors properties of both metals and nonmetals

Electron Configurations An element’s chemical properties are determined by its e- config. Differences in properties of elements are due to different e- configurations, not just placement on the chart.

Ions Ion – an atom or group of atoms having a charge can be (+) or (-) charged due to gain or loss of electrons Cation – positive ion, forms when metals LOSE e- Anion – negative ion, forms when nonmetals GAIN e-

Ions

Ions

Oxidation Number The Oxidation Number (charge) tells the number of electrons an atom will lose or gain to achieve stability Group 1 = +1 Aluminum = +3 Group 2 = +2 Zinc = +2 Group 18 = 0 Cadmium = +2 Group 17 = -1 Silver = +1 Group 16 = -2 Carbon = -4 or +4 Group 15 = -3 (nonmetals) Hydrogen = +1 or -1 Transition metals can lose e- from “s” or “d” sublevel or both, multiple oxidation #s (usually +1 or +2).

Periodic Trends

Atomic Radius Atomic Radius – how big an atom is (p. 163) decreases across a period b/c Zeff increases increases down a group b/c more shells are added

Zeff Zeff = effective nuclear charge This is how much charge the negative valence electrons “feel” from the positive Zeff can be approximated by subtracting the number of inner (non-valence) electrons from the atomic number.

*Fluorine's outer electrons experience a stronger attraction to the nucleus and are pulled in closer.

Periodic Trends Ionic Radius – how big an ion is (charged particle); Metal ions – have lost electrons, charge is (+), cation Ionic radius is smaller than atomic radius isoelectronic with preceding noble gas size of metal cations decreases across a period increases down a group Nonmetal ions – have gained electrons, charge is (-), anion Ionic radius is larger than atomic radius isoelectronic with following noble gas size of nonmetal anions

Ionic Radius

Electronegativity Electronegativity – the ability of an atom to attract e- when in a compound (p.168) Increases across a period. Why? Increased Zeff attracts electrons closer to nucleus. decreases down a group

Ionization Energy Ionization Energy – energy required to remove an e- from an atom (units = kJ/mol) {p167.} First Ionization Energy – energy required to remove most loosely held e- increases across a period decreases down a group Metals have low ionization energies. Why? They easily lose electrons in order to become stable. Nonmetals have high ionization energies. Why? Losing electrons makes them more unstable.

Shielding Trends affecting ionization energy Shielding – inner e- block Zeff attraction of (+) nucleus for outer e- constant across a period increases down a group e- feel less charge How does this affect ionization energy? It causes IE to increase

Summary of Periodic Trends

The periodic table is the most important tool in the chemist’s toolbox!