The Periodic Table Chemistry

Origin of Periodic Table Triads - groups of 3 elements with similar properties (Dobereiner – 1817) Law of Octaves – properties of elements repeat every 8 elements (Newlands – 1863)

Origin of Periodic Table Dimitri Mendeleev – 1869 Properties of elements are periodic functions of their atomic masses. Developed 8 column table Left spaces for undiscovered elements Columns contained elements with similar properties Problems with Ni, I, K (Why?)

Origin of Periodic Table X-ray experiments by Mosley (1913) led to discovery of atomic numbers. Modern Periodic Law – properties of elements are periodic functions of their atomic numbers.

Organization of Periodic Table Rows are called periods. Columns are called families or groups. All elements in a family have similar properties. Octet Rule – elements with 8 valence electrons are unreactive

Periodic Table Rows = Periods Columns = Families or Groups

Chemistry Chapter 5 The Periodic Law

Mendeleev’s Periodic Table Dmitri Mendeleev

Periodic Table with Group Names

Alkali Metals 1st column in blue

The Properties of a Group: the Alkali Metals Easily lose valence electron (Reducing agents) React violently with water Large hydration energy React with halogens to form salts H- Hydrogen Rb- Rubidium Li-Lithium Cs- Cesium Na-Sodium Fr-Francium K-Potassium

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

Alkaline Earth Metals 2nd column - green

Properties of alkaline metals Be, Mg, Ca, Sr, Ba, Ra The alkaline earth metals, or alkaline earths, are beryllium, magnesium, calcium, strontium, barium, and radium. Possess many properties of metals Low electro negativities Low electron affinities They have smaller atomic radii than the alkali metals

Transition Metals all have similar properties Middle – purple and can give up different amounts of electrons at different times

Properties of Metalloids B- Boron As- Arsenic At- Astatine Ge- Germanium Te-Tellurium Sb- Antimony Si- Silicon Al-Aluminum Po- Polonium 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

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

Right of the stair-step line Non-Metals Right of the stair-step line C-Carbon, N- Nitrogen, O-Oxygen, P-Phosphorus, Cl-Chlorine, Se-Selenium

7th tall column (light orange) Halogens 7th tall column (light orange) F-fluorine Cl-Chlorine Br-Bromine I- Iodine

Properties of Halogens a particular class of nonmetals. Very high electronegativities Seven valence electrons (one short of a stable octet) Highly reactive, especially with alkali metals and alkaline earths Halogens range from solid (I2) to liquid (Br2) to gaseous (F2 and Cl2) at room temperature.

last tall column (yellow) Noble Gases last tall column (yellow) He-Helium Ne-Neon Ar-Argon Kr-Krypton Xe-Xenon

Properties of Noble gases Unreactive gases Odorless Colorless All produce light when an electric current is applied

Lanthanides – Ce row Actinides – Th row

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

Electron Affinity - the energy change associated with the addition of an electron Affinity tends to increase across a period Affinity tends to decrease as you go down in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals

Table of Electron Affinities

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

Summation of Periodic Trends

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 Electronegativities tend to decrease down a group or remain the same

Periodic Table of Electronegativities