2. Friday 12/04/15

3. Objectives Understand the general trends in atomic properties in the periodic table Understand the nature of bonds and their relationship to electronegativity

4. Reading the periodic table Groups or families – vertical columns Periods – horizontal rows

5. Term Effective Nuclear Charge (ENC) 1) The net charge that pulls on the valence electrons in an atom. The greater the effective core charge, the greater the pull. It is determined by subtracting the number of core electrons from the number of protons in the nucleus  For example: Magnesium

6. Periodic Trends Atomic radius The distance from the center of an atoms nucleus to it’s outermost electron Measure of atomic size

7. Periodic Trends Graph the first 20 elements. What is the trend down a group? Across a Period? Atomic radius

8. Periodic Trends

9. Atomic Radius Group Trend  Increases down a group  More energy levels or quantum levels (or “shell”) as you go down a group – atomic radius increases Period Trend  Decreases across a period  All electrons in the same energy level. Increased # of protons holds them closer to nucleus.  Increase in Effective Core (Nuclear) Charge (ECC) Calculate ECC for elements in period 2

10. Table of Atomic Radii

11. Period Trend: Atomic Radius

12. Periodic Trends Worksheet Work with a table partner to answer questions:  1,5a, 6, 8,

13. Periodic Trends Ionic Size Size of an atom when electrons are added or removed. Electrons removed atom becomes smaller. Electrons added atoms become larger Why?  Electron-Electron Repulsion

14. Ionic Size Cations  Positively charged ions formed when an atom of a metal loses one or more electrons  Smaller than the corresponding atom Anions  Negatively charged ions formed when nonmetallic atoms gain one or more electrons  Larger than the corresponding atom

15. Periodic Trends Graph the first 20 elements. What is the trend down a group? Across a Period? Ionic Size (label P.T.)

16. Table of Ion Sizes

17. Ionic Size Group Trend  Increases down a group  More energy levels as you go down a group – ionic size increases Period Trend  Decreases as atoms lose more electrons  Increases dramatically as atoms start gaining electrons, decreases as atoms gain fewer electrons.

18. Periodic Trends Ionization Energy Energy needed to remove one of the electrons on an atom’s outer shell. How strongly does an atom hold it’s outermost electron.

19. Periodic Trends Graph the first 20 elements. What is the trend down a group? Across a Period? Ionization Energy

21. Group Trends  Increases up a group.  The closer outer shell electrons are to the nucleus the harder they are to remove. Period Trend  Increases across a period.  The more electrons in the outer shell the harder it is to remove one.  Increase in Effective Core Charge (ECC)

22. Periodic Trend: Ionization Energy

23. Periodic Trends Worksheet Work with a table partner to answer questions:  3, 5b, 7, 9

24. Periodic Trends Electronegativity Is a measure of the level of attraction (pull) an atom exerts on the electrons of another atom. Ability of an atom to attract electrons Which elements want to gain electrons the most?

25. Periodic Trends Graph the first 20 elements. What is the trend down a group? Across a Period? Electronegativity

27. Periodic Table of Electronegativities

28. Electronegativity Group Trend  Increases up a group  As radius decreases, electrons are closer to the nucleus (decrease in number of electron shells) Period Trend  Increases across a period  The more electrons in the outer shell (up to 7) the more the atom wants to attract electrons Exception: Trend does not apply to Noble Gases  Increase in Effective Core Charge (ECC)

29. Periodic Trend: Electronegativity

30. Periodic Trends Worksheet Work with a table partner to answer questions:  2, 4, 5c, 10, 11

31. Summary of Periodic Trends

32. Practice 1. Se and Br 1. Smallest atom 2. Lowest Ionization Energy 2. P, S, Se 1. Largest atom 2. Highest Ionization Energy 3. Cl, Cl 1-, Br, Br 1- 1. Largest ionic size 4. Mg, Mg 2+, Na, Na 1+ 1. Smallest ionic size

33. Atomic Properties Definitions For Quiz – Monday Effective Nuclear Charge:  It is the net charge that pulls on the valence electrons in an atom.  The greater the effective nuclear charge, the greater the pull.  It is determined by subtracting the number of core electrons from the number of protons in the nucleus Valence Electrons  Are found in the outermost, valence, electron shell (Bohr model) of the atom Core electrons  occupy all of the inner electron, core, shells

34. Atomic Properties Definitions Ionization Energy:  Energy needed to remove an electron from an atom or molecule.  The higher the effective core charge and lower the number of electrons shells, the greater the ionization energy Atomic size  How big (e.g., radius) an atom is  Atomic radius is measured from the center of the nucleus to the valence electron shell.  The higher the effective core charge and lower the number of electron shells, the smaller the atom. Electronegativity  Measure of the level of attraction (pull) an atom exerts on the electrons of another atom.  The higher the effective core charge and lower the number of electron shells, the greater the electronegativity

35. Homework Read pages:  327-331 Answer questions: Pg 336 69-78  343-345 Answer questions: Pgs 374-375, 7-20 Due 12/09

36. Periodic Table Objective: Students know how to use the periodic table to identify alkali metals, alkaline earth metals, transition metals, metals, semimetals (metalloids), nonmetals, halogens and noble gases.

38. Alkali Metals  All alkali metals have 1 valence electron  They are very reactive  Reactivity of these elements increases down the group  Alkali metals: http://video.google.com/videopl ay?docid=- 2134266654801392897# http://video.google.com/videopl ay?docid=- 2134266654801392897# Potassium, K reacts with water and must be stored in kerosene

39. Alkaline Earth Metals All alkaline earth metals have 2 valence electrons Alkaline earth metals are less reactive than alkali metals The word “alkaline” means “basic”  common bases include salts of the metals Ca(OH) 2 Mg(OH) 2

40. Properties of Metals  Metals are good conductors of heat and electricity  Metals are malleable  Metals are ductile  Metals have high tensile strength  Metals have luster

41. Transition Metals Copper, Cu, is a relatively soft metal, and a very good electrical conductor. Mercury, Hg, is the only metal that exists as a liquid at room temperature

42. Properties of Metalloids  They have properties of both metals and nonmetals.  Metalloids are more brittle than metals, less brittle than most nonmetallic solids  Metalloids are semiconductors of electricity  Some metalloids possess metallic luster

43. Silicon, Si – A Metalloid  Silicon has metallic luster  Silicon is brittle like a nonmetal  Silicon is a semiconductor of electricity Other metalloids include:  Boron, B  Germanium, Ge  Arsenic, As  Antimony, Sb  Tellurium, Te

44. Nonmetals  Nonmetals are poor conductors of heat and electricity  Nonmetals tend to be brittle  Many nonmetals are gases at room temperature Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.

45. Examples of Nonmetals Sulfur, S, was once known as “brimstone” Microspheres of phosphorus, P, a reactive nonmetal Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

46. Halogens  Halogens all have 7 valence electrons  Halogens in their pure form are diatomic molecules (F 2, Cl 2, Br 2, and I 2 ) Chlorine is a yellow-green poisonous gas

47. Noble Gases Noble gases have 8 valence electrons (except helium, which has only 2) they are chemically unreactive Colorless, odorless and unreactive; they were among the last of the natural elements to be discovered