The History of the Modern Periodic Table

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.

John Newlands

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

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

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

Mendeleev’s discovery

The fateful day for Mendeleev was 17 February 1869 (Julian calendar). He cancelled a planned visit to a factory and stayed at home working on the problem of how to arrange the chemical elements in a systematic way. To aid him in this endeavor he wrote each element and its chief properties on a separate card and began to lay these out in various patterns.

Mendeleev’s discovery

Eventually he achieved a layout that suited him and copied it down on paper. Later that same day he decided a better arrangement was possible and made a copy of that, which had similar elements grouped in vertical columns, unlike his first table, which grouped them horizontally. These historic documents still exist.

That Mendeleev realized that he had discovered, rather than designed, the periodic table is shown by his attitude towards it.

Mendeleev’s discovery

First, he left gaps in it for missing elements. Leaving such gaps in tables of elements was not in itself new, but Mendeleev was so sure of himself that he was prepared to predict the physical and chemical properties of these undiscovered elements. His most notable successes were with eka aluminium (= Gallium) and eka-silicon (= germanium). Lecoq de Boisbaudran discovered gallium in 1875 and reported its density as 4.7g cm -3 , which did not agree with Mendeleev’s prediction of 5.9 g cm -3 .

Mendeleev’s discovery

When he was told that his new element was Mendeleev’s eka-aluminium, and had most of its properties foretold accurately, Boisbaudran redetermined its density more accurately and found it to be as predicted, 5.956 g cm -3 . There could be no doubt now that Mendeleev had discovered a fundamental pattern of Nature.

Mendeleev’s discovery

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

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.

WWII.

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

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

These elements are also called the rare-earth elements.

InnerTransition Metals

Halogens

Noble Gases

The s and p block elements are called REPRESENTATIVE ELEMENTS.

Modern Periodic Table

The “Real” Periodic Table

Modern Periodic Table

Groups

•

Alkali Metals (Group 1)

– 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

•

Halogens (Group 17)

– most reactive nonmetals – never found alone in nature – gain or share e- to achieve stability Chlorine Bromine Iodine

•

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 Group 2 = +2 Group 18 = 0 Group 17 = -1 Aluminum = +3 Zinc = +2 Cadmium = +2 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 Z eff increases – increases down a group b/c more shells are added

I n r c e a s e se Increa

Atomic Radius

One half the distance between nuclei of identical atoms that are bonded together.

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 – decreases across a period – increases down a group

Ionic Radius

I n r c e a s e

Ionic Radius

One half the distance between nuclei of Identical ions that are bonded together .

Electronegativity

• Electronegativity – the ability of an atom to attract e- when in a compound (p.168) – Increases across a period. Why?

• Increased Z eff attracts electrons closer to nucleus. – decreases down a group

Electronegativity

The measure of the ability of an atom in a chemical compound to attract electrons

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.

I n s e c r e a

Ionization Energy

The amount of energy required to remove one electron from a neutral atom

Shielding

• Trends affecting ionization energy –

Shielding

– inner e- block Z eff nucleus for outer e attraction of (+) • constant across a period • increases down a group • e- feel less charge • How does this affect ionization energy?

– It causes IE to increase

Z eff

• Z eff = effective nuclear charge • This is how much charge the negative valence electrons “feel” from the positive protons • Z eff 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.

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