1. Periodic Properties of the Elements Chapter 7 Sam White Nathan Chan Keyana Porter ***Bronstar***

2. We will cover… Scientists Development of the Sizes of Atoms Ionization Energy Electron Affinities Metals, Nonmetals, and Metalloids Group Trends for Active Metals and Selected Nonmetals

3. Scientists to Remember Jar Newlands- arranged elements in order of atomic mass Dmitri Medeleev- Credited with creating the first periodic table, he sorted elements by increasing mass. Saw many patterns in properties, and put elements with similar properties in the same column. He was even able to predict the properties of undiscovered elements (such as Germanium)

4. More Scientists Henry Mosely- Developed concept of atomic number, basing it off of the number of protons in the atom. Also discovered that elements emitted X-rays when bombarded with high energy electrons, and the X-ray frequency increased with atomic mass.

5. Electron Shells in the Atom Rows on the Periodic Table = Periods Columns on the Periodic Table = Families/Groups –These elements have similar properties and similar configurations

6. Electron Shells in an Atom As a period increases, quantum number n increases. n represents an electron shell / principle energy level Gilbert N. Lewis suggested that electrons are arranged in spherical shells around the nucleus http://library.thinkquest.org/18406/media/Bohr.jpg

7. Electron Distribution You can calculate the distribution of electrons in an atom by using quantitative aspects of quantum mechanics (electron configurations) Radial Electron Density- Probability of finding an electron at a particular distance from the nucleus

8. Electron Clouds Even though distribution has a shell structure, there are no “hard” boundaries. Thus, the edges of an atom are “fuzzy” http://members.tripod.com/craigjm /Atom2-2.jpg

9. Atomic Radius Non-bonding radius of an atom is the closest distance seperating the nuclei of a collision between 2 atoms Bonding atomic radius is the radius of 2 atoms attractively bonded (this is smaller than non-bonding radius)

10. Periodic Trends Atomic radius increases down a group and across a period http://www.district87.org/staff/sutterm/web questpt/images/ptable.gif

11. Ionization Energy Ionization energy- Minimum Energy required to remove an electron from the ground state of an isolated gaseous atom or ion (measured in J/mole) The greater the Ionization Energy the more difficult it is to remove an electron. –Because the positive nuclear charge, which is the attractive force, remains the same, while the number of electrons, which produces repulsive interaction, decreases

12. Ionization Energy Energy needed to remove an electron from the outer shell depends on both effective nuclear charge and average distance from the nucleus When inner shell electrons are removed, there is a sharp increase of ionization energy –They are closer to the nucleus, so they experience a much greater nuclear attraction than valence electrons

13. Stages of Ionization Energy 1 st Ionization Energy (I 1 ): Energy needed to remove the first electron from a neutral atom Na (g) -> Na + (g) + e - 2 nd Ionization Energy (I 2 ): Energy needed to remove the second electron Na + (g) -> Na 2+ (g) + e - This continues for successive removal of additional electrons

14. Periodic Trends Within each row, ionization energy increases with increasing atomic number (with few irregularities) Within each column, ionization energy decreases with increasing atomic number Elements with highest sub-level of s or p show a wider variation of I 1 values than elements with highest sub-level d or f

15. Ionization Energy Trends grandinetti.org

16. Ionization Energy Trends http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/ionize.html

17. Electron Affinities Electron affinity: the energy change that occurs when an electron is added to a gaseous atom The less the attraction between a given atom and added electron, the more negative the affinity will be *Cl is the most electronegative of all the elements

18. Electron Affinities As you go down a group, the distance that the incoming electron is from the nucleus increases This decreases the attraction and decrease the energy that is released as electron affinity Becomes increasingly negative as we proceed toward the halogens in each row

19. Electron Affinity vs Ionization Energy Ionization Energy measures the ease with which an atom LOSES an electron Electron Affinity measures ease with which an atom GAINS an electron

20. Electronegativity Electronegativity: measure of the tendency of an atom to attract a bonding pair of electrons Pauling scale: 0.7 (cesium and francium which are the least electronegative) to 4.0 (fluorine- most electronegative element) If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons (they would need to be the same atom)

21. Metals, Nonmetals, & Metalloids Metals –Great heat & electricity conductors –Able to be pounded into thin sheets (malleable) & drawn into a wire (ductile) –Bright luster –Solid at room temp. –Lose electrons during chemical reactions –Transition metals have more than one oxidation state

22. Metals, Nonmetals, & Metalloids Non-metals –Brittle & soft –Dull colors –Poor conductors of heat & electricity –Tend to gain electrons –2 bonded non-metals are molecular substances Metalloids/Semi-metals –Possesses traits of both metals & non-metals

23. Group Trends for Active Metals ands Selected Nonmetals Metallic character: the more an element exhibits physical and chemical properties of a metal