Do not copy any notes in green lettering for this unit!

Dmitri Mendeleev (1869) First organize the elements in groups according to their physical and chemical properties Predicted undiscovered elements and their properties

Henry Moseley Reorganized Mendeleev’s table in order of increasing atomic number Just like today’s modern Periodic Table

Periodic Law When elements are arranged according to atomic numbers, elements with similar properties appear at regular intervals

Groups – vertical column of elements; share chemical properties Aka: Families Periods – horizontal row of elements VC – Periodic Table Overview

Valence Electrons Electrons in outer most energy level Determine the chemical properties of an element Show students how to count valence electrons using periodic table. Explain elements in the same group have the same chemical properties b/c they have the same number of valence electrons (This also explains the periodic law!)

Some elements more reactive than others! The closer an atom is to having 8 valence electrons, the more reactive it is Atoms with 8 valence electrons are unreactive (aka stable) Use Bohr models of sodium vs magnesium & sulfur vs chlorine to explain why group 1 is more reactive than group 2 & why group 17 is more reactive than group 16

Metals (in blue) Good conductors of heat & electricity Shiny surface appearance Malleable – able to be hammered into sheets Ductile – able to be drawn into a wire VC – Properties of Metals: Malleability & Ductility

Nonmetals (in green) Poor conductors of heat and electricity Dull surface appearance Very brittle

Metalloids Characteristics of metals & nonmetals Semiconductors (conduct electricity only at high temperatures) Surface can appear shiny or dull More brittle than metals, but more malleable than nonmetals VC – Comparing Metals, Nonmetals, & Metalloids

Main Group Elements Groups 1,2, 13-18 Very wide range of physical and chemical properties S and P blocks AKA: Representative Elements

Alkali Metals Group 1 Extremely reactive (1 valence electron) React with water to make alkaline (basic) solutions Soft, low density

Alkaline-Earth Metals Group 2 Highly reactive (2 valence electrons) Harder than alkali metals

Halogens Group 17 Nonmetals Highly reactive (7 valence electrons) React with most metals to produce salts Greek – “salt maker”

Noble Gases Group 18 Unreactive (8 valence electrons) Helium used in balloons Others used in lamps

Transition Metals Groups 3 – 12 D and F blocks Relatively unreactive

Lanthanide & Actinide series F block Lanthanides – aka rare-earth series Actinides are radioactive –unstable nucleus spontaneously breaks apart Uranium used in nuclear power plants & bombs

Alloy – homogeneous mixture of two or more metals Gold & silver in jewelry Copper & zinc make brass Iron mixed with a variety of elements to produce different types of steel

Hydrogen Most common element in the universe one proton + one electron Behaves unlike all other elements Found in all living things

Diatomic The seven elements that exist in nature as two atoms of the same element bonded together BrINClHOF or Br2 I2 N2 Cl2 H2 O2 F2 7-up

Periodic Trends

Atomic Radius – half the distance between the nuclei of two identical atoms that are not bonded together Use Bohr Models to explain trend

Ionization Energy – amount of energy required to remove an electron from an atom Electron Shielding Use Bohr Models and magnets to explain why

Electron Shielding Electrons on inner energy levels reduce the attraction between the nucleus and valence electrons (valence electrons held loosely to the atom) Causes atoms to get bigger down a group Causes ionization energy to decrease down a group No affect across periods Use Bohr Models to explain why

Electronegativity – ability of an atom in a chemical compound to attract electrons Use Bohr Models to explain why

Periodic Trends Decreases Decreases Increases Atomic Radius Decreases Ionization Energy Increases Electronegativity Increases

Electron Configuration The arrangement of electrons in atoms

Electron Cloud made of Orbitals Orbital – three-dimensional region around the nucleus that indicates the probable location of an electron 2 electrons per orbital 4 types of orbitals: s, p, d, f

Orbital blocks on periodic Table (pg 119)

energy level 1s2 that orbital The Break Down # of electron in energy level 1s2 that orbital type of orbital

The Rules Every box on the P.T. represents 1 electron Each row represents an energy level Each block represents an orbital Start with H, read left to right across the rows until the superscripts add up to the atomic # of the element you want Be careful when you get to the D and F blocks

Try These! Write the electron Configuration for: Boron, Phosphorus, Calcium

Follow the arrows 4f14 5f14 3d10 4d10 5d10 6d10 2p6 3p6 4p6 5p6 6p6

DO NOT DO LIKE THE TEXT BOOK DOES for elements after calcium: s-d-p is correct d-s-p is wrong

Try These! Write the electron Configuration for: Nickel, Strontium, Iodine

Noble Gas Notation Same rules as electron configuration, except don’t start with H, start with the last noble gas before the element you want

Try These! Write the Noble Gas Configuration for: Boron, Phosphorus, Calcium, Nickel, Strontium, Iodine