1. Drill pd 311/17/14 Rank in order of increasing ionization energy.  Si, C, S, F, Ca Rank in order of decreasing atomic radii.  F, S, N Rank in order of increasing atomic radius  As, Te, P, Al Rank in order of increasing electronegativity.  Ge, P, As, Cs HINT: Look at group trends first, then period trends

2. Answers  Ca, Si, S, C, F  F, N, S  P, Al, As, Te  Cs, Ge, As, P

3. Drill 4A/4B11/17& 18/2014  Differentiate between ionization energy and electronegativity.  What types of periodic trends are you noticing for atomic radius, ionization energy and electronegativity?  Ionization energy is the energy needed to REMOVE one of the atom’s outermost electrons while an atom’s electronegativity is its ability to ATTRACT electrons to itself in a chemical bond.

4. Announcement  Conclusions have been graded. You have one week to turn back into me for ½ credit.  PBIS stuff

5. SWBAT: Analyze the trends of ionization energy, electronegativity and atomic radius by completing a graphing activity.

6. Homework  -2 Review and Reinforcement – pd 4A and 4B  Graph

7. PERIODIC TRENDS GRAPHING ACTIVITY

8. GRAPH DATA Construct 1 graph for ionization energy, electronegativity and atomic radius. Answer questions on the Worksheet.

9. Octet Rule  Atoms tend to gain, lose, or share electrons so that each atom has a full outermost energy level, which typically consists of 8 electrons (called an octet)

10. Periodic Table Patterns

11. Nuclear Charge  Nuclear charge-the charge in the nucleus or the number of protons  Across – increases  Down – increases

12. Atomic Mass  Weighted average (based on mass and percent abundance of each naturally occurring isotope)  Across – increases  Down – increases  Remember there are some exceptions to this – like Te and I

13. Atomic Radius  One-half the distance between the nuclei of identical atoms that are bonded together.

14. Ionization Energy (IE)  Ionization energy is the energy needed to REMOVE one of the atom’s outermost electrons

15. Electronegativity  An atom’s electronegativity is its ability to attract electrons to itself in a chemical bond

16. Trend for Atomic Radius Atomic radius decreases as you go from the left to the right across a period. This is because there is an increased nuclear charge (protons) but no additional “shielding” electrons come between valence electrons and nucleus. Atomic radius increases as you go down a group. Increased distance of electrons and additional electron energy levels shield valence electrons.

17. Shielding Effect  The shielding effect is the reduction of attractive force between the nucleus (+) and its outer electrons (-) due to the blocking affect of the inner electrons Nucleus Shielding electrons Valence electrons ‘shielded’ by inner electrons

18. Snowman analogy

19. Trend for Ionization Energy Ionization energy increases as you go from the left to the right across a period  this means it’s MORE DIFFICULT to remove an electron because of increased positive nuclear charge. Ionization energy decreases as you go down a group  this means it’s EASIER to remove an electron because as atomic radius increases, outermost electrons are further away from nucleus.

21. Trend for Electronegativity Electronegativity increases as you go from the left to the right across a period Electronegativity generally decreases as you go down a group in the periodic table ** ** Because the noble gases form very few compounds, they do not have an electronegativity value.

22. Trend for Electronegativity (cont)  The highest electronegativity values are located in the upper-right hand corner of the periodic table  Fluorine has the highest electronegativity value of 4.0  The lowest electronegativity values are located in the lower-left hand corner of the periodic table  It is NOT an amount of energy

24. Ionization Energy (IE) - review  Ionization energy – this is the energy needed to remove one of the atom’s outermost electrons  Metals have low IE  Nonmetals have high IE

25. Trend for IE  1) IE’s decrease as you move down a group  why?

26. Why?  The electrons in larger atoms are held less strongly by the nucleus, therefore as you move down a group, the IE decreases

27. Trend for IE  2) IE’s increase as you move from the left to the right across a period  why?

28. WHY?  As you move across a period, there is a decrease in atomic radius, therefore causing the IE to increase

29. Removing electrons from positive ions  With sufficient energy, electrons can be removed from positive ions.  The energy required to remove a second electron from an atom is called its second IE, and so on…  The second IE will be greater than the first IE, the third IE greater than the second, etc. because as electrons are removed, fewer electrons remain to shield the attractive nuclear force.

30. Electron Affinity (EA)  Neutral atoms can acquire electrons.  An atom’s electron affinity is the energy change that occurs when an electron is acquired by a neutral atom.

31. Trend for EA  They change in irregular ways across a period and down a group  General rules:  nonmetals have higher electron affinities than do metals (this means it’s usually easier to add an electron to a nonmetal)  The halogens gain electrons most readily – higher electron affinities.

32. Trends Across the Periodic Table Shielding Effect Stays the Same Atomic Radius decreases Ionization Energy Increases Electronegativity Increases

33. Trends Down the Periodic Table Shielding Effect Increases Atomic Radius Increases Ionization Energy Decreases Electronegativity Decreases

34. Drill 11/19/14 Comparison of Na and Cl atoms 1. Both atoms are in the same period. Explain the differences in their sizes. 2. Predict the charge that an ion of each would have. Explain your answer. 3. Compare the ionization energy required to remove the first electron from each of these atoms. 4. Draw the ion for each atom. Make sure to represent their sizes relative to their original sizes. Ignore sizes of atoms:

35. Answers to Drill 1. Cl is smaller than Na because the increased positive nuclear charge exerts a stronger pull on the e- shrinking the cloud. 2. A Na ion would have a charge of +1 and Cl ion, a charge of -1 because Na must lose an e- to have a full octet and Cl must gain an e-. 3. The IE of Cl would be higher because nonmetals hold on to their valence e- more tightly. 4. The Cl ion is now larger than the Na ion. Sodium lost its outermost orbital. The positive nuclear charge is felt more strongly.