1. Periodic Table Trends

2. Periods  Rows (There are 7!) Families/Groups Columns Have similar chemical and physical properties

3. Most elements are metals that conduct heat and electricity quite well With the exception of Hg (Mercury) they all are solid at room temp. Gallium will melt in your hand! Most are malleable (able to be hammered into sheets) Most are Ductile  able to be pulled into wire Metals (Outlined in Black from left to right) **Gold (Au) is soft and is mixed with other metals such as copper and zinc to create an alloy for jewelry **White gold is comprised of “yellow “ gold and a whiter metal, such as Rhodium

4. Most are gasses, others are brittle solids. Bromine is a liquid at room temperature. Poor conductors of electricity Carbon and Hydrogen are the backbone of organic molecules Non-Metals (Dark Blue Elements on the Right Plus Hydrogen)

5. Have both metal and non-metal properties, depending on conditions Such as semi- conductors in electronics (silicon) Metalloids (Stair-step In Green)

6. Alkali Metals Group 1 (not including hydrogen) S-block (all have s 1 ; 1s 1, 2s 1, 3s 1 ….. Valence electron configuration) Lower densities than other metals One loosely bound valence electron and will readily form +1 ions, called cations, to bond with non- metals that “need” an octect Largest atomic radii in their periods Low ionization energies Low electronegativity Highly reactive with air and water, Will form solutions with high pH in H2O (bases), Cs will catch on fire when it is exposed to air!

7. Alkaline Earth Metals Group 2 S-block (all have valence electron configuration of s 2 ) Two electrons in the outer shell Readily form +2 cations. Will give up 2 electrons to nonmetals that need to fill their valence shell (octect rule!) low electron affinities Low electronegativity

8. Transition Metals D-block – Contains ALL s and p orbitals; stable elements half filled or half-filled d orbitals Form metallic bonds with other metals (a sea of electrons) All are solid at room temp. except for Hg (liquid) Will form cations that vary in number of valence electron’s (+1, +2, +3 )

9. Halogens Very electronegative (Fluorine the most) Readily gains an electron to form a -1 Anion Diatomic in Nature  F 2 Cl 2 Br 2 I 2 At 2 Called Halogens as they will bond with metals to produce salts, such NaCl or CaCl 2 Synthetic (man-made)

10. Noble Gasses Valence is completely filled with 2s2 for He and the rest are 3s 2 3p 6, 4s 2 4p 6, etc. satisfying the octet rule All of the other atoms are jealous All of the other atoms are jealous Inert, meaning not naturally reactive

11. Periodic Trends – Atomic Mass Tends to Increase as Atomic Number Increases

12. Periodic Trends – Atomic Radii Decreases Increases

14. Periodic Trends – Ionization Energy Energy required to “pluck” high energy electrons from the valence shell of an atom in its gaseous stage Increases across the table BECAUSE…. the atomic radius decreases across, and atoms are more tightly compact. It’s harder to remove an electron that is so close to the nucleus. Great attraction there!! Opposite in larger atoms, so ionization energy decreases as you go down (vertically) the table Halogens already have 7 valence electrons…..rather gain than lose one, so it takes LOTS of energy to get them to “let go” of their electrons!

15. Periodic Trends – Electron Affinity Typically applies to groups A6 and A7 since they need electrons to gain an octet Energy released when an electron is gained by a neutral atom to produce an anion Elements are in a gaseous phase

17. Periodic Trends - Electronegativity Electronegativity increases across the table (Fluorine is the MOST electronegative) measure of the tendency of an atom to attract a bonding pair of electrons. Those atoms in groups 6A and 7A are the most electronegative because they only need 2 or 1 electrons to complete their octect. Fluorine is assigned a value of 4.0, and values range down to cesium and francium which are the least electronegative at 0.7 (From the Pauling Scale)

19. Ionic Radii  Radius of the atom in it’s ionic state; after it either gained or lost or gained an electron or electrons Some Before and After Examples

20. Atomic and Ionic Radii Comparison (excluding transition metals). Atoms on the left and their ions on the right. Notice the size difference!!! Why does this occur???

21. Atomic Radii Versus Ionic Radii of Main Group Elements