Chemistry Chapter 5 The Periodic Law

Mendeleev’s Periodic Table Dmitri Mendeleev

Mendeleev – organized periodic table Vertical columns in atomic mass order Made some exceptions to place elements in rows with similar properties (Te and I) Horizontal rows have similar chemical properties Gaps for “yet to be discovered” elements Left questions: why didn’t some elements fit in order of increasing mass? Why did some elements exhibit periodic behavior?

Moseley Discovered that periodic table was in atomic number order, not atomic mass order Explained the Te-I anomaly

Periodic Law Physical and chemical properties of the elements are periodic functions of their atomic numbers

Modern Periodic Table Discovery of noble gases yields new family (Group 18 – aka inert gases) Lanthanides (#58 - #71) Actinides (#90 – #103)

Periods and Blocks of the Periodic Table Periods – horizontal rows Groups/Families – vertical columns; these elements share similar chemical properties (they have the same number of valence electrons) Blocks – periodic table can be broken into blocks corresponding to s, p, d, f sublevels

Orbital filling table

Blocks and Groups – s block Group1 – “The alkali metals” One s electron in outer shell Soft, silvery metals of low density and low melting points Highly reactive, never found pure in nature

Blocks and Groups – s block Group 2 – “Alkaline Earth Metals” 2 s electrons in outer shell Denser, harder, stronger, less reactive than Group 1 Too reactive to be found pure in nature

Periodic Table with Group Names

The Properties of a Group: the Alkali Metals Easily lose valence electron (Reducing agents) React violently with water React with halogens to form salts

Blocks and Groups – d block Metals with typical metallic properties Referred to as transition metals Group number = sum of outermost s and d electrons

Properties of Metals Metals are good conductors of heat and electricity Metals are malleable Metals are ductile Metals have high tensile strength Metals have luster

Examples of Metals Potassium, K reacts with water and must be stored in kerosene Copper, Cu, is a relatively soft metal, and a very good electrical conductor. Zinc, Zn, is more stable than potassium Mercury, Hg, is the only metal that exists as a liquid at room temperature

Blocks and Groups – p block Properties vary greatly – metals, metalloids, and nonmetals Group 17 – halogens are most reactive of non metals Group 18 – noble gases are NOT reactive

Properties of Nonmetals Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element. Nonmetals are poor conductors of heat and electricity Nonmetals tend to be brittle Many nonmetals are gases at room temperature

Examples of Nonmetals Microspheres of phosphorus, P, a reactive nonmetal Sulfur, S, was once known as “brimstone” Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

Properties of Metalloids Metalloids straddle the border between metals and nonmetals on the periodic table. They have properties of both metals and nonmetals. Metalloids are more brittle than metals, less brittle than most nonmetallic solids Metalloids are semiconductors of electricity Some metalloids possess metallic luster

Silicon, Si – A Metalloid Silicon has metallic luster Silicon is brittle like a nonmetal Silicon is a semiconductor of electricity Other metalloids include: Boron, B Germanium, Ge Arsenic, As Antimony, Sb Tellurium, Te

Blocks and Groups – f block Lanthanides – shiny metals similar to group 2 Actindes – all are radioactive; plutonium – lawrencium are man-made

Determination of Atomic Radius: Half of the distance between nucli in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels

Table of Atomic Radii

Ionization Energy - the energy required to remove an electron from an atom Increases for successive electrons taken from the same atom Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels     Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group Outer electrons are farther from the nucleus

Ionization of Magnesium Mg + 738 kJ  Mg+ + e- Mg+ + 1451 kJ  Mg2+ + e- Mg2+ + 7733 kJ  Mg3+ + e-

Table of 1st Ionization Energies

Another Way to Look at Ionization Energy

Ionic Radii Cations Anions Positively charged ions Smaller than the corresponding atom Anions Negatively charged ions Larger than the corresponding atom

Table of Ion Sizes

Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons Electronegativities tend to increase across a period * more nuclear charge, more power to attract electrons Electronegativities tend to decrease down a group or remain the same * additional energy levels result in less attraction to the nucleus

Periodic Table of Electronegativities