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 propertiesPredicted 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 levelDetermine 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 & electricityShiny surface appearanceMalleable – able to be hammered into sheetsDuctile – able to be drawn into a wire

VC – Properties of Metals: Malleability & Ductility

Nonmetals (in green)

Poor conductors of heat and electricityDull surface appearanceVery brittle

Metalloids

Characteristics of metals & nonmetalsSemiconductors (conduct electricity only at high temperatures)Surface can appear shiny or dullMore brittle than metals, but more malleable than nonmetals

VC – Comparing Metals, Nonmetals, & Metalloids

Main Group ElementsGroups 1,2, 13-18Very wide range of physical and chemical propertiesS and P blocksAKA: Representative Elements

Alkali MetalsGroup 1Extremely reactive

(1 valence electron)

React with water to make alkaline (basic) solutionsSoft, low density

Alkaline-Earth Metals

Group 2Highly reactive (2 valence electrons)

Harder than alkali metals

Halogens

Group 17NonmetalsHighly reactive

(7 valence electrons)

React with most metals to produce saltsGreek – “salt maker”

Noble GasesGroup 18Unreactive (8 valence electrons)

Helium used in balloonsOthers used in lamps

Transition Metals

Groups 3 – 12 D and F blocksRelatively unreactive

Lanthanide & Actinide series

F blockLanthanides – 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

HydrogenMost common element in the universeone proton + one electronBehaves unlike all other elementsFound in all living things

DiatomicThe 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 groupCauses ionization energy to decrease down a groupNo 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

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Rf Db Sg Bh Hs Mt

Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Atomic Radius Decreases

Ionization Energy Increases

Electronegativity Increases

Increases

Decreases

Decreases

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)

The Break Down

# of electron

in

energy level 1s2 that orbital

type of orbital

The RulesEvery box on the P.T. represents 1

electron Each row represents an energy level Each block represents an orbitalStart 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

4s23s22s21s2 5s2 6s2 7s2

4p6

4d10

4f14

3p62p6 5p6 6p6 7p6

3d10 5d10 6d10

5f14

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