Chapter 5- The Periodic Law 5.1-History of the Periodic Table 5.2-The Periodic Table 5.3-Periodic Trends

5.1-History of the Periodic Table Pages

Mendeleev  Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass. Elements with similar properties were grouped together. There were some discrepancies.

Mendeleev  Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered elements.

Moseley  Henry Moseley (1913, British) Organized elements by increasing atomic number. Resolved discrepancies in Mendeleev’s arrangement. Periodic Law-the physical and chemical properties of the elements are periodic functions of the atomic numbers.

Organization of the Elements  Periodic table-an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

Additions to Mendeleev’s Periodic Table  Noble gases Group 18 Argon discovered in 1894 Took so long to discover because very unreactive  Lanthanides 14 elements with atomic numbers from Placed below the periodic table to conserve space  Actinides 14 elements with atomic numbers Also placed below periodic table

5.2-The Periodic Table Pages

 Metals  Nonmetals  Metalloids Metallic Character

 Main Group Elements  Transition Metals  Inner Transition Metals Areas of the Periodic Table

Alkali metals  Group 1 metals  ONE VALENCE ELECTRON  Silvery appearance and very soft (can be cut with a butter knife)  Not found pure naturally because so reactive  Because of extreme reactivity with moisture, usually stored under kerosene  Video: Disposal of Surplus SodiumDisposal of Surplus Sodium  Video: Alkali Metals in WaterAlkali Metals in Water

Alkaline-Earth metals  Group 2 metals  TWO VALENCE ELECTRONS  Harder, denser, & stronger than alkali metals  Also too reactive to be found free in nature (but less reactive than Gp. 1)  Video: Magnesium/silver nitrate mixture reacting with waterMagnesium/silver nitrate mixture reacting with water

TRANSITION METALS  d-Block Elements: Groups 3-12  Metals with typical metallic properties  Typically less reactive than Gps. 1&2, & some are extremely unreactive

METALLOIDS Metalloids Fall on both sides of a “stair-step” line separating metals and nonmetals Semi-conductors The elements in the upper right of the line show increasing non- metallic behavior and the elements at the lower left of the line show increasing metallic behavior. The list of metalloids in the periodic table are as follows: Boron (B) Silicon (Si) Germanium (Ge) Arsenic (As) Antimony (Sb) Tellurium (Te) Polonium (Po)

GROUP 17~HALOGENS React vigorously with most metals to form salts. “Salt formers” -SEVEN VALENCE ELECTRONS -Fluorine and chlorine are gases at room temp., bromine is a reddish liquid, and iodine is a dark purple solid. -Form diatomic molecules

NOBLE GASES ~ GROUP 18 Unreactive gases EIGHT valence electrons. Exception is He (2 valence electrons) Outershell is “FULL”, which is what makes them stable and nonreactive

Lanthanides – Atomic # added to the periodic table in Similar in reactivity to group 2. Actinides – Atomic # 90 – 103; All radioactive; 1st three are naturally occurring, the rest are lab made Inner Transition Metals

5.3-Electron Configuration & Periodic Properties Pages

Remember the Periodic Law  When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

 ½ the distance between the nuclei of identical atoms that are bonded together  Increases to the LEFT and DOWN Atomic Radius

Li Ar Ne K Na

 Why larger going down? Higher energy levels have larger orbitals Shielding - core e - block the attraction between the nucleus and the valence e -  Why smaller to the right? Increased nuclear charge without additional shielding pulls e - in tighter Atomic Radius

 First Ionization Energy-energy required to remove one electron from a neutral atom  Increases UP and to the RIGHT Ionization Energy

 First Ionization Energy Ionization Energy K Na Li Ar Ne He

 Why opposite of atomic radius? In small atoms, e - are close to the nucleus where the attraction is stronger  Why small jumps within each group? Stable e - configurations don’t want to lose e - Ionization Energy

 Successive Ionization Energies Mg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ Large jump in I.E. occurs when a CORE e - is removed. Ionization Energy

Al1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ  Successive Ionization Energies Large jump in I.E. occurs when a CORE e - is removed. Ionization Energy

 Energy change that occurs when an electron is acquired by a neutral atom  Tends to become less negative (less energy released) DOWN and to the LEFT Electron Affinity

 Ionic Radius Cations (+)  lose e -  smaller © 2002 Prentice-Hall, Inc. Anions (–)  gain e -  larger Ionic Radius

Electronegativity  A measure of the ability of an atom in a chemical compound to attract electrons  Most electronegative element is fluorine Given arbitrary value of 4; all others relative

 Which atom has the larger radius? BeorBa CaorBr Ba Ca Examples

 Which atom has the higher 1st I.E.? NorBi BaorNe N Ne Examples

 Which has the greater electonegativity? KorLi AlorCl Li Cl Examples

 Which particle has the larger radius? SorS 2- AlorAl 3+ S 2- Al Examples