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