Chapter Periodic Table Lecture

Do members of the same family, generally behave the same? Yes

The Periodic Table The Alkali Metals Lithium, Sodium, Potassium, Rubidium, Cesium, and francium very reactive 1 valence electron s 1 sublevel is filled Alkali Earth Metals Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium 2 valence electrons s 2 sublevel is filled The Transition Metals metals with atomic numbers highest s & d sublevels have electrons Metalloids Like metals & nonmetals Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium

Nonmetals Consists of Carbon, Nitrogen, Oxygen, Phosphors, Sulfur, Selenium poor conductors of heat and electricity compared to metals dull and brittle Halogens Consists of Fluorine, Chlorine, Bromine, Iodine, Astatine nonmetals have 7 valence electrons very reactive want one more electron (octet rule) Noble Gases consists of Helium, neon, Argon, Krypton, Xenon, Radon unreactive stable inert because they already have 8 valence electrons Inner Transition metals consists of elements with atomic numbers 58 through 71 and 90 through 103 F sublevels partially filled the Lanthanide Series has atomic numbers and the Actinide Series has atomic numbers Other Metals

Define the term inert gas?terminert noble gas –unreactive & stable

Group 1A Group 2A Group 3A Group 4A Group 5A Group 6A Group 7A Representative Elements #1 – Group IA-VIIA outer s & p orb partially filled Alkali Metals Alkaline Earth Nonmetals/Metalloids Halogens ns1ns1 ns2ns2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 Group 0 8 or 18 Noble Gasesns 2 np 6

Representative Elements #1 Lewis dot structure 1s 2 2s 2 2p 6 1s 2 Na

Group BTransition Metals Filling the “d” orbital Group LanthanidesFilling the “4f” orbital Group Actinides Filling the “5f” orbital

A. Ionic Size metals (group 1A-3A) lose electrons to become stable cation non-meta l (group 5A-7A) gain electrons to become stable anions. 1A = 2A = 3A = 5A = 6A = 7A = Loses 1 e - Loses 2 e - Loses 3 e - Gains 3 e - Gains 2 e - Gains 1 e -

7 P E R I O D S ! v A Family is a Group living between Columns

Periodic Table Song by Tom Lehrer above Periodic Table Song by Tom Lehrer above End of Lecture 6.1 Next Lecture h/elements.html h/elements.html

Who designed the 1 st periodic table in 1869? Dmitri Mendeleev grouped w/ similar chemical and physical properties & ordered by atomic mass.Ex: Co Ni Ar K Te I

Lecture 6.3 Periodic Trends

I. Periodic Trends - Atomic Size Atomic Radii: Nucleus Distance between nuclei Atomic Radius Measured as 1/2 distance between nuclei 2 atoms

Atomic Size generally INCREASES as you move down a group on the periodic table. Why? down a group increases # of energy levels Example: Ca atom larger than a Mg atom. Why? An energy level is added!

Atomic Size generally DECREASES across a row on the periodic table. Why? adding more p + pulls in extra electrons

Na < ionization energy than O because less protons pull. RELATIVE ELECTRONEGATIVITY, IONIZATION ENERGY, RADII, SHIELDING ETC… Hydrogen 2.1 Oxygen 3.5 Carbon 2.5 Sodium 0.9 Electro negativities: Hydrogen has the smallest atomic radius

B. Ionization Energy energy needed to pull an electron away from an atom.

B. Ionization Energy Example : Na Na +1 + e -

Ionization energy decreases as you move down a group. increased distance from protons reduces attractive force Why?

Period Trend: Ionization energy generally increases as you move across a period. nuclear charge increases (more protons) which increases attractive forces Why?

energy required to remove the 1st outermost electron is 1st ionization energy. What is the second ionization energy? Which is harder to remove? Why?

What happens to the shielding of the nucleus as you move across a period? ONLY adding electrons, NOT a new energy level. Remains constant Why?

What happens to the shielding of the nucleus as you move down a group? another energy level that shields those valence electrons. Increases Why?

Ca + ions – smaller than the original atom When electrons lost, a whole energy level lost decreases radius. Why?

Negative anions grow larger there are more e - than p + (increased electron repulsion), Why? N atom N -3 anion

from group 5A to the right, anions gradually decrease in size groups 6A &7A only gain 1 or 2 e - Have Same # of e-, but increased # of p + Why? N -3 O -2 F -1 anion anion anion.

B. Electronegativity Noble gases no electronegative # Why? inert / don’t form compounds. Can’t force a noble gas to take an electron – they have s 2 p 6

3. Period Trends left to right electronegativity increases. Why? High ionization energy = high electronegativity Resists electron loss Attracts electrons Fluorine is the most electronegative!

4. Group Trends Electronegativity decreases down a group. Why? Increased energy levels and shielding Cs has the lowest electronegativity