Unit 4: The Periodic Table History and Trends Chapters 6 & 7 Test - October 31, 2008

Periodic Table A systematic arrangement of the elements by atomic number (protons) Similar properties fall into vertical columns

Dmitri Mendeleev (1869) developed the first periodic table Arranged by increasing atomic mass Aligned properties in vertical columns Some elements fell into the wrong column Examples: Te & I ; Co & Ni History of the Periodic Table

Mendeleev’s Periodic Table

Henry Moseley (1913) rearranged the modern table by atomic number Used X-ray spectrometer to find the atomic numbers History of the Periodic Table

Arrangement of the Periodic Table

States on the Periodic Table Red Gases Black Solids Blue Liquids

Types of Elements Metals, Nonmetals, & Metalloids

Metals Found on the left side of table Have 1, 2 or 3 electrons in valence Lose electrons to form positive ions (cations) Most are silver, shiny, solid, malleable, ductile & good heat/electrical conductors

Metals

Nonmetals Found on the right side of table Have 5, 6, or 7 electrons in valence Gain electrons to form negative ions (anions) Brittle, dull, non-conductors, and exist in all three states (solids, liquids, gases)

Nonmetals

Metalloids Elements found along the stairstep between metals and nonmetals, excluding Al B, Si, Ge, As, Sb, Te, Po Properties are intermediate between metals & nonmetals

Metalloids

Oxidation Numbers Metals + ; lose valence electrons Nonmetals - ; gain valence electrons Metalloids oxidation numbers vary Transition elements multiple oxidation numbers Number of electrons lost or gained when reacting to achieve full valence (ex: Na + ; O 2- )

Oxidation Numbers /2+ 0 Multiple oxidation numbers

Groups/Families A group/family is a vertical column They have similar properties ("A" families only!!) Same # of valence electrons (roman numerals at the top)

"A" Group Families Electron configuration ends in "s" or “p” sublevel Referred to as the representative elements

“A” Groups/Families

"B" Group Families a.k.a. Transition Metals Electron configuration ends in “d” sublevel

“B” Groups/Families

Inner-Transition Metals 2 bottom rows Electron configuration ends in “f” sublevels Really belong in rows 6 & 7 before transition metals Elements beyond #92 are man- made (too unstable to be found in nature)

Inner-Transition Metals

Group/Family Names IA – Alkali Metals Most active metals Stored under oil (water reactive) 1+ oxidation number Ex: Li, K

IA - Alkali Metals

Group/Family Names IIA – Alkaline Earth Metals Less reactive than IA 2+ oxidation number Ex: Mg, Ca, Sr

IIA - Alkaline Earth Metals

Group/Family Names VIIA- Halogens Most active nonmetals 1- oxidation number Ex: F, Cl

VIIA - Halogens

Group/Family Names VIIIA – Noble Gases Relatively inactive (inert) Contain 8 valence electrons All gases Ex: Ne, Xe

VIIIA - Noble Gases

Hydrogen Unique element, most abundant in the universe (75%) Nonmetal on the left side of the periodic table (found in metal column) Gas state One valence electron

Hydrogen

Nitrogen As the chemically unreactive molecule N 2, makes up 78 percent by volume of Earth’s atmosphere. Plants and animals can’t use nitrogen in this form. Lichens, soil bacteria, and bacteria in the root nodules of beans, clover and other similar plants convert nitrogen to ammonia and nitrate compounds. Lightning also converts atmospheric nitrogen to nitrogen monoxide (NO). Plants use these simple nitrogen compounds to make proteins and other complex nitrogen compounds that become part of the food chain.

Period/Series Horizontal rows Share the same number of energy levels

Periodic Table Trends Atomic Radius Electronegativity Metal Activity Nonmetal Activity

Atomic Radius Half the distance between the centers of two atoms that are just touching each other Influenced by two factors 1. The number of energy levels in an atom 2. The pull of the positively charged nucleus on its electrons The more energy levels, the ________ the atomic radius. (larger/smaller) The more protons in the nucleus, the ________ the atomic radius. (larger/smaller)

Atomic Radius & Ions Metals lose electrons to form cations LiLi + F F-F- Nonmetals gain electrons to form anions Atomic radius decreases - energy level is lost or “shed” Atomic radius increases - energy level expands because it is more “crowded” and electrons exert greater forces on each other

Ionic Radius

Atomic Radius - Trend Atomic radius increases as you move down a group Atomic radius decreases as you move from left to right in a period

Electronegativity How easily an atom attracts the valence electrons of another atom Assigned number “rates” the electronegativity (from 0.7 to 4.0) Low electronegativity = does not want to attract valence electrons High electronegativity = really wants to attract valence electrons

Electronegativity & Stability Because of electronegativity… Stable atoms find it harder to bond with other atoms and have lower (re)activity Example: noble gases Unstable atoms find it easier to bond with other atoms and have greater (re)activity Example: alkali metals and halogens

Electronegativity - Trend Electronegativity decreases as you move down a group Electronegativity increases as you move from left to right in a period

Metal Activity How easily metals lose e- to form cations Since electronegativity is how easily atoms attract valence electrons, the most active metals are the least electronegative (they want to lose e-)! Trend Increases as you move down a group Decreases as you move from left to right in a period *The most active metal is francium decreasing metal activity increasing metal activity

Nonmetal Activity How easily nonmetals gain e- to form anions Since electronegativity is how easily atoms attract valence electrons, the most active nonmetals are the most electronegative (they want to gain e-)! Trend Decreases as you move down a group Increases as you move from left to right in a period *The most active nonmetal is fluorine increasing nonmetal activity decreasing nonmetal activity

Question 1 Match each element in the left column with the best matching description in the right column. Descriptions may be used more than once. 1. Strontium a. halogen 2. Chromium b. alkaline earth metal 3. Iodine c. representative element d. transition element

Question 2 For each of the following pairs, predict which atom is larger. A. Mg, Sr B. Sr, Sn C. Ge, Sn D. Ge, Br E. Cr, W

Question 3 Compare the alkali and alkaline earth metals in terms of position in the periodic table, number of valence electrons, and overall properties.

Question 4 Compare the metallic character of the elements carbon, silicon, and lead. What do these elements have in common in terms of valance electrons and placement in the periodic table?