1. PERIODIC PATTERNS Unit 3 – Periodic Table

2. What patterns exist on the periodic table? Lesson Essential Question:

3. METALLIC TREND INCREASES

4. ATOMIC RADIUS  Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.  Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x 10 -10 m.

5. ATOMIC RADIUS BROMINE = Br 2  Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms. 2.86 Å 1.43 Å

7. ATOMIC RADII TRENDS  As you go down a family the number of energy levels increases making the radius larger. DOWN A FAMILY OR GROUP WHY? INCREASES

8. ATOMIC RADII TRENDS  As you go across a period the number of protons increases, (nuclear charge) pulling the electrons in tighter making the radius smaller. ACROSS A PERIOD WHY? DECREASES

9. IONS - remember Metals  Lose electrons becoming positive. Calcium (Ca) Loses 2 electrons becoming Ca +2 and [Ar] Noble gas Configuration. (Octet Rule) Nonmetals  Gain electrons becoming negative. Chlorine (Cl) Gains one e - becoming Cl -1 and [Ar] Noble gas configuration. (Octet Rule)

10. IONS – How can I remember? Metals Nonmetals This is Ann ion - ANION She is unhappy and negative. This is Cat-ion - CATION He is a “plussy” cat!

11. IONIC RADII TRENDS  As you go down a family the number of electron shells increases making the radius larger. DOWN A FAMILY OR GROUP WHY? INCREASES

12. IONIC RADII TRENDS  For the metals the nuclear charge is greater than then number of electrons pulling them in tighter making the radius smaller.  At the nonmetals the radius gets larger because the ion has gained electrons. ACROSS A PERIOD WHY? DECREASES then INCREASE

13. METALLIC ATOM AND ION COMPARISON

14. NONMETALLIC ATOM AND ION COMPARISON Why do the Noble Gases not have an ionic Radius?

15. ATOM AND ION COMPARISON Why does Hydrogen not have an ionic Radius?

16. As more electrons are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held Shielding Effect

17. Example of Shielding Effect

18. The energy required to remove an electron from an atom. (measured in kilojoules, kJ) Ionization Energy

20. IONIZATION TREND INCREASES Why? Closer to nucleus (more +) Electrons less likely to be removed Requires more energy to form ion Less shielding

21. IONIZATION ENERGY The larger the atom is, the easier its electrons are to remove. (Why?) Ionization energy and atomic radius are inversely proportional. Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

22. IONIZATION TREND INCREASES NCREASES Why? Elements in alkali metals have 1 valence electron so what to remove that electron, they therefore take the least amount of energy to remove an electron

23. is a measure of the tendency of an atom to attract a bonding pair of electrons. Electronegativity

24. Why? Closer to nucleus (more +) so electrons are more attracted INCREASES http://www.thecatalyst.org/electabl.html

25. Electronegativity Why? Elements in halogens only need 1 more electron to have a full valence shell so are MOST likely to attract electrons. As you move to left elements are more likely to LOSE electrons. INCREASES http://www.thecatalyst.org/electabl.html

27. In Summary…. Electronegativity