The Periodic Table

The Periodic Table Understand the rationale behind the periodic table; view the table as an ordered database of element properties. Explain how the periodic table reflects the quantum mechanical structure of the atom. Explain and use periodic trends in: atomic radius ionic radius ionization energy Explain the connection between ionization energy and metallic character.

Quick Review  Quantum Numbers (4)  Represent or describe the space orbital in which the e - moves or may be found. 1. Distance from nucleus 2. Shape of the orbital 3. Position or location with respect to the 3 axes in space 4. Direction of spin of the e - in the orbital

Quantum Numbers 1st Quantum Number: Principle Quantum # (n) – avg. distance of e- from nucleus. Positive whole # 1,2,3, etc. The main energy level designation equal to K,L,M, etc. shells

Quantum # 2 Orbital Quantum # (l) – shape of orbital in which the e - moves. # of possible shapes = the value of the principal quantum #. 1 st energy level = 1 shape 1 st energy level = 1 shape 2 nd energy level = 2 shapes 2 nd energy level = 2 shapes 3 rd energy level = 3 shapes, etc 3 rd energy level = 3 shapes, etc Letter designations for shapes are s,p,d,f. These are also sublevels or sub orbitals S – Sphere, P – 8 ( 3 - one on each axis x,y,z)

3 rd Quantum # - Magnetic (m) – indicates the position about the three axes in space of the orbital. 3 rd Quantum # - Magnetic (m) – indicates the position about the three axes in space of the orbital. 1 position for s orbital 1 position for s orbital 3 positions for p orbitals (p x p y p z ) 3 positions for p orbitals (p x p y p z ) 5 positions for d orbitals 5 positions for d orbitals 7 positions for f orbitals 7 positions for f orbitals

4th Quantum Number Spin (s) – right hand or left hand, clockwise or counter-clockwise Therefore 2 e- can occupy any space designated by the first 3 quantum #’s but they will have opposite spins. No 2 e- have the same quantum #’s No 2 e- have the same energy

 In 1859 Mendeleyev devised the periodic table of elements.  “Properties of the elements are in periodic dependence on their atomic weights.”  Ramsay discovered the noble gases: Ne, Ar, Kr, Xe. Noble or inert: will not react with others.  Periodic Law: The physical & chemical properties of the elements are periodic functions of their atomic #’s. 1. Period/Series – horizontal rows 1. Period/Series – horizontal rows 2. Group/Family – vertical columns 2. Group/Family – vertical columns a. elements with similar properties have a similar a. elements with similar properties have a similar arrangement of outer-shell e - (same group in the arrangement of outer-shell e - (same group in the periodic table) or valence e -. periodic table) or valence e -.

 Most active metal – lower left  Most active non-metal – upper right  Transition elements are metallic & have either 1 or 2 e - in their outer shell (generally)  Rare Earth elements are basically identical. The outer shell contains 2 e -  Zigzag line: separates metals from the non- metals. Those touching are metalloids, have both or either characteristics of metals and non- metals.

Atomic Radii The radius of an atom does not increase with atomic #. 1. Atomic radius increases with atomic # in a particular group/family. (increased shells) 2. From group 1 to 18 (I to VIII) in a period, there is a general decrease in the atom radii. (increase in p+, increase attraction of the e-)

Atomic radius

Ionization Energy  Energy required to remove an e - from an atom.  normally, ionization removes valence electrons first  valence electrons are farthest from nucleus on average, so they feel the least attraction for the nucleus and are easiest to remove  end of valence electrons is marked by a big jump in ionization energies

 A o (nuetral) + Energy = A + (positive ion) + e -  kJ – kilo joules or kcal – kilo calories may be the energy unit of measure used per mole of substance.  Ex: It takes 124 kcal to remove the outer shell e - from 6.02 X atoms of Li.  Mole: amount of the substance equal to a Avogadro Number: 6.02 X atoms Avogadro Number: 6.02 X atoms  Mole: may also equal that substances atomic weight in grams.

Quick Check for understanding  normally, ionization removes valence electrons first  valence electrons are farthest from nucleus on average, so they feel the least attraction for the nucleus and are easiest to remove  end of valence electrons is marked by a big jump in ionization energies  Low ionization energy is characteristic of a metal.  High ionization energy is characteristic of a non- metal  Intermediate ionization energy for metalloids

 Within a group/family ionization energy generally decreases with increasing atomic #  Ionization energy does not vary uniformly within a period/series as a result of filled or unfilled sublevels.  To remove successive e - from the same atom becomes progressively more difficult due to the closeness of the remaining e - to the nucleus.  Na + Energy = Na + + e kJ 1 st ionization energy +496 kJ 1 st ionization energy  Na + Energy = Na e-  kJ 2 nd ionization energy

Electron Affinity  Energy released when an e - is added to a neutral atom.  A + e - = A - + energy (kcal) Note: E on Product side of the equation  Low electron affinity – weak bonding  High electron affinity – strong bonding  Electron affinities decrease in groups/families due to additional shells  Electronegativity: ability to attract e -

Additional Terms  Ion – an atom which has become charged by either gaining or losing electrons.  Anion – negatively charged atom, characteristic of the non-metals.  Cation – positively charged atom, characteristic of the metals.

Group or Family Names: 1 Alkali metals, 2 Alkaline earth metals, 3-12 transition metals, 13 Boron-Aluminum, 14 Carbon, 15 Nitrogen, 16 Chalcogen, 17 Halogens (salt formers), 18 Noble (Inert) or Group O Rare Earth Elements

Ionic Bonding Covalent Bonding Electron donor, lender Electron acceptor, borrower

CHEMICAL FAMILIES: ALKALI METALS: GROUP IA (1) All alkali metals react with H 2 O to form an alkaline (Basic) Solution. ** Hydrogen is NOT considered an alkaline metal in this case. ALKALINE EARTH METALS: GROUP IIA (2) All alkali earth metals also react with H 2 O to form a basic solution. HALOGENS: GROUP VIIA (17) Salt Formers NOBLE GASES: GROUP VIIIA (18) Noble Gases have all their orbitals Filled. The other groups are identified by the element at the top of the column. Example, GROUP IVA is the called the Carbon Family (or Group). Lanthanide Series: Fourteen elements beginning with lanthanum in which the highest energy electrons to be in the 4f sublevel. Actinide Series: Fourteen elements beginning with actinium in which the highest energy electrons to be in the 5f sublevel. Rare Earth Elements – Inner Transition Elements