1. PERIODIC TRENDS Unit 3 – Periodic Table

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

3. The Basics

5. Metals, Nonmetals, Metalloids  There is a zig-zag or staircase line that divides the table.  Metals are on the left of the line, in blue.  Nonmetals are on the right of the line, in orange.

6.  Elements that border the stair case, shown in purple are the metalloids or semi- metals.  There is one important exception.  Aluminum is more metallic than not.

7.  Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity.  They are mostly solids at room temp.  Get oxidized (lose electrons), giving a positive oxidation state  Nonmetals are the opposite.  They are dull, brittle, nonconductors (insulators).  Some are solid, but many are gases, and Bromine is a liquid.  Get reduced (gain electrons), giving a negative oxidation state

8.  Metalloids, aka semi-metals are just that.  They have characteristics of both metals and nonmetals.  They are shiny but brittle.  And they are semiconductors.

9. METALLIC TREND INCREASES

10.  Metallic characteristics are greatest to the left of the periodic table due to fact that the attraction between valence electron and the nucleus is weaker, enabling an easier loss of electrons.  Metallic characteristics increases as you move down a group because the atomic size is increasing so valence electrons are farther from the nucleus and easier to remove.

11. 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.

12. 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

14. 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

15. 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

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

17. 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

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

19. 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

20. Example of Shielding Effect

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

22. Trends in Ionization Energy

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

24. 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.

25. 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

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

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

28. 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

30. In Summary…. Electronegativity