Development of the Periodic Table Antoine Lavosier In

Development of the Periodic Table • Antoine Lavosier • In the late 1790’s, compiled a list of the 23 elements known at that time. • J. W. Dobereiner • In 1829, published a classification system in which he grouped elements in triads – a set of 3 elements with similar properties. • By this time, almost 70 elements were known. • Unfortunately, all known elements could not be grouped in triads.

Development of the Periodic Table • John Newlands • In 1864, proposed an organization scheme with elements arranged by increasing atomic mass. • He observed that properties repeated every eighth element, calling this relationship the “law of octaves” (an octave is a group of musical notes that repeats every eighth tone). • Acceptance of the law of octaves was very little because it did not work with all of the known elements and scientists associated the word octave with music, not science. • Dmitri Mendeleev • In 1869, developed a periodic table in which elements were arranged into groups based on a set of repeating properties. • Left spaces in his table for elements that he predicted would be discovered in the future based on gaps he noticed in the properties of the elements that had been discovered.

Development of the Periodic Table • When new elements were discovered, the actual properties of those elements were very close to the predicted properties. • Because of this table’s ability to predict properties of undiscovered elements, it was widely accepted by the scientific community. • This arrangement of the periodic table was based on increasing atomic mass since little was known about atomic structure at this time. • Henry Moseley • In 1913, determined an atomic number for each known element. • He re-arranged the elements by their atomic number – the order observed in the modern periodic table. • The patterns observed became known as periodic law: When elements are arranged in order of increasing atomic number, there is periodic repetition of their physical and chemical properties.

The Periodic Table

Metals • Make up about 80% of the elements • Form cations (positively charged atoms) • Good conductors of heat and electricity • Malleable – can be hammered into thin sheets • Ductile – can be drawn into wire • High luster, or sheen

Nonmetals • Most are gases at room temperature • Forms anions (negatively charged atoms) • Poor conductors of heat and electricity • Solid nonmetals tend to be brittle and dulllooking

Metalloids • Staircase arrangement of elements • Has properties similar to those of metals and nonmetals • Many are semi-conductors of heat and electricity

Alkali Metals – Group 1 (except H) • Silvery, soft solid • 1 valence electron • Forms +1 charge • Very reactive with water • Very reactive, especially with Group 17 elements

Alkaline Earth Metals – Group 2 • Soft, shiny solids (harder than alkali metals) • 2 valence electrons • Forms +2 charge • Less reactive than alkali metals • Reactive with oxygen • Can be very reactive with water

Transition Elements - d and f blocks • Two types: transition metals (d- block) and inner transition metals (f-block) • Number of valence electrons varies

Halogens – Group 17 • Called “salt formers” – they can form compounds with almost all metals • Exist as solids (I and At), liquids (Br), and gases (F and Cl) at room temperature • 7 valence electrons • Forms -1 charge • Very reactive with metals

Noble Gases – Group 18 • Colorless gases • Unreactive, or inert • 8 valence electrons – a full outer energy level, except He which has only 2 valence electrons • Very stable

Nuclear Charge and Electron Shielding • Nuclear Charge – directly proportional to the number of protons in the nucleus. • Increases across a period due to increases in the number of protons in the nucleus without adding energy levels. • Decreases down a group due to additional energy levels • Electron Shielding – refers to the repulsion between negative core electrons and valence electrons and will change the effective nuclear charge. • Increases down a group due to the increases in energy levels.

Atomic Radius • Atomic Radius – the size of a neutral atom • Increases as you go down a group because of the addition of energy levels. • Decreases as you go across a period as a result of the increasing nuclear charge that occurs by the addition of one more proton and one more electron. • The electrons in the higher energy levels are pulled closer to the nucleus making the atom smaller.

Atomic Radius

Ionic Radius • Ionic radius – the size of an ion, a charged atom • A cation (positively charged atom) becomes smaller when it loses electrons. • An anion (negatively charged atom) becomes larger when it gains electrons.

Ionic Radius

Ionization Energy • Ionization energy (first ionization energy) – the energy required to remove the first valence electron from an atom. • First ionization energy decreases down a group because as the atom becomes larger the nuclear charge has a smaller effect on the outer electrons. • First ionization energy increases across a period because the nuclear charge increases.

Ionization Energy

Electronegativity • Electronegativity – the ability of an atom’s nucleus to attract electrons from another atom when the atom is in a compound. • Electronegativity decreases down a group because as the atom becomes larger the nuclear charge has a smaller effect on the outer electrons. • Electronegativity increases as you move across a period as a result of increasing nuclear charge.

Electronegativity INCREASES DECREASES

Reactivity • Reactivity– the chemical reactivity of an element depends upon the relative attraction between valence electrons of one element and the nucleus of another. Metals Nonmetals • increases down a group • decreases across a period • decreases down a group • Increases across a period