Reactivity series

1. Reactivity series http://www.youtube.com/watch?v=m55kgyApYrY BrainiacRb, Cs http://www.youtube.com/watch?v=SjowQJMS-W4&NR=1 F Long sort of boring but useful 10 minute clip for alkali earth metals up to strontium http://www.youtube.com/watch?v=SjowQJMS-W4&NR=1 Show mixing of Ba and Ca with water (same size pieces for rate and strength of reaction + show at 7:50 to show the product Ba(OH)2 and its difference in properties from the water http://www.youtube.com/watch?v=hiBVj8rT_hA&NR=1 likely not true—imagined impact http://www.webelements.com/francium/chemistry.html Commentary on Fr bomb video Commentary—you tube video likely not real http://www.youtube.com/watch?v=PyFLvSg6ZDw

2. Subatomic structure basics, 3.1—3.7Quantum mechanical model, chapter 3-7 Electron configuration of an atom The probable (>90% of the time) locations of all electrons in a stable atom of an element. Ground state Stable atoms have the lowest possible energy associated with organization and movement of their electron clouds—the ground state.

3. When an atom’s electrons are all in their ground state locations of the electron cloud, the atom is stable. The electrons are located as close as possible to the nucleus whose proton’s positive charges attract them. The electrons are located as far apart as possible from other electrons whose like charges repel them. The attraction & repulsion is balanced.

4. If an electron absorbs energy (e.g., when an element is heated or when electricity passes through it), it is destabilized—called excited. The electrons absorbing energy—the excited electrons-- move into a new, higher energy location of the electron cloud. They are repelled by the electrons now too near them, and the attraction to the nucleus versus repulsion by electrons is no longer balanced. The atom is made unstable.

5. Electrons are organized into the cloud in a systematic way according to energy Electrons are added to the electron cloud in order of their energy, from lowest to highest energy. The distance of an electron from the nucleus as it travels through the electron cloud is defined as its ENERGY LEVEL. In each energy level exist pathways of electron travel having characteristic energy and shape are called ORBITALS. The ground state is stable (lowest energy) If electrons absorb energy, they are excited and unstable; they quickly release the energy to return to ground state.

6. Electron configuration • The probable (>90% of the time) locations of all electrons in a stable atom of an element. • The address of each electron is specified: • Which energy level • Which orbital type • How many electrons in the orbital

7. Energy Levels are defined by their average distance from the nucleus Energy levels • are assigned numbers n = 1, 2, 3, 4, and so on. • increase in energy as the value of n increases. • are like the rungs of a ladder with the lower energy levels nearer the nucleus.

8. Energy Levels Energy levels have a maximum number of electrons equal to 2n2. Why? As the number of the energy level increases, the volume of the “shell” also increases. More electrons can travel in the shell without coming so close that like charge repulsion destabilizes the atoms. Energy level Maximum number of electrons n = 1 2(1)2 = 2(1) = 2 n = 2 2(2)2 = 2(4) = 8 n = 3 2(3)2 = 2(9) = 18

9. Orbitals An orbital • is a three-dimensional space around a nucleus, where an electron is most likely to be found. • has a shape that represents electron density (not a path the electron follows). • can hold up to 2 electrons.

10. Orbitals

11. Energy levels are sub-divided into orbitals—s, p, d, and f which have differing energies, even when in the same energy level. Not all energy levels contain all types of orbitals. Some orbitals are present in several copies in the same shell.

12. Electron configuration • The probable (>90% of the time) locations of all electrons in a stable atom of an element. • The address of each electron is specified: • Which energy level • Which orbital type • How many electrons in the orbital

13. Energy Levels are defined by their average distance from the nucleus Energy levels • are assigned numbers n = 1, 2, 3, 4, and so on. • increase in energy as the value of n increases. • are like the rungs of a ladder with the lower energy levels nearer the nucleus.

14. Energy Levels Energy levels have a maximum number of electrons equal to 2n2. Why? As the number of the energy level increases, the volume of the “shell” also increases. More electrons can travel in the shell without coming so close that like charge repulsion destabilizes the atoms. Energy level Max #of electrons n = 1 2 n = 2 8 n = 3 18

15. Orbitals An orbital • is a three-dimensional space around a nucleus, where an electron is most likely to be found. • has a shape that represents electron density (not a path the electron follows). • can hold up to 2 electrons.

16. Orbitals

17. Energy levels are sub-divided into orbitals—s, p, d, and f which have differing energies, even when in the same energy level. Not all energy levels contain all types of orbitals. Some orbitals (p, d, f) are present in several copies in the same shell.

18. d orbitals are even more complex and high energy than p orbitals d orbitals are found on the 3rd and higher energy levels

19. f orbitals are even more complex and high energy than p orbitals f orbitals are found on the 4th and higher energy levels

20. Orbitals A p orbital • has a two-lobed shape. • is one of three p orbitals in each energy level from n = 2. • An s orbital • has a spherical shape around the nucleus. • is found in each energy level.

21. Electron Level Arrangement In the electron level arrangement for the first 18 elements • electrons are placed in energy levels (1, 2, 3, etc.), beginning with the lowest energy level • there is a maximum number in each energy level. Energy level Number of electrons 1 2 (up to He) 2 8 (up to Ne) 3 18 (up to Ar) 4 32 (up to Ca)

22. Writing electron configurationsLesson for Friday 9/24 – Tuesday 9/28 Be able to use either the arrow diagram, the block version of the periodic table, or the Aufbau (orbital) diagrams to write electron configurations of elements. Write shorthand and complete electron configurations. Quiz on Wednesday, 9/29

23. Electrons enter the electron cloud in a predictable order, starting with the lowest energy orbital, 1s, then sequentially occupying the new higher available orbital. Order in which the energy levels (#) and orbitals (letter) fill Number electrons added to each orbital when it is filled. https://teach.lanecc.edu/gaudias/scheme.gif

24. The position of an element in the electron cloud reveals its electron configuration. Elements add electrons to the energy levels and orbitals in this order, one at a time, as the atomic number increases. www.mikeblaber.org/.../EPeriod/IMG00011.GIF

25. Each orbital contains only 2 electrons, but some orbitals have multiple copies in the same energy level3 equal energy p orbitals5 equal energy d* orbitals**7 equal energy f* orbitals*** SOMETIMES THESE WILL VIOLATE RULES OF FILLING (SKIP ONE OR TWO ELECTRONS IN THE LAST ENERGY S ORBITAL BEFORE MOVING INTO THE D ORBITAL—IT’S TO GET A STABLE HALF FILLED OR FULLY FILLED D OR F ORBITAL—DON’T WORRY ABOUT THIS AS IT’S BEYOND THE LEVEL OF THE COURSE…JUST FOLLOW THE GENERAL RULES FOR ELECTRON CONFIGURATION)**some periodic tableversions will move elementsout to the f block after thenext higher energy d blockadds 1 electron; some will fill all the forbitals before any electronsare added to the d block. Follow the atomic number on whatever chart your professor or standardized test provides!

26. Another way to visualize the order of electron addition is an Aufbau diagram. The orbitals are shown in order of increasing energy, the order in which they are filled. Aufbau’s principle 2 electrons per orbital, added in order of increasing energy Hund’s rule When orbitals in the same energy level have the same energy, then 1 electron is added to each before either is filled Pauli’s exclusion principle The electrons in the same orbital have opposite spin

27. Take out your Aufbau diagram, then fill it out as we use this animation that shows the order electrons are added to the electron cloud http://intro.chem.okstate.edu/WorkshopFolder/Electronconfnew.html

28. Now add the electron configurations for each Aufbau diagram you drew Homework • review slides from today. • review text 3-7 • Complete Aufbau diagrams using the website posted on the last slide • Check to see that the order follows the diagram chart and periodic table block charts↗

29. Stop—end of lesson for 10/29/09 http://faculty.washington.edu/dwoodman/aufbau/dswmedia/aufbauW.html AMAZING WEBSITE TO LINK THE AUFBAU DIAGRAMS AND WRITING ELECTRON CONFIGURATIONS

30. Electron configurations show the locations of each electron, without the need of showing the Aufbau diagram. 1s22s22p6… left hand number shows the energy level letter shows the orbital superscript shows the # electrons in the orbital type on that energy level

31. Shorthand electron configuration • Write the symbol of the previous noble gas (family 18) in brackets, followed by electron configuration details for every electron added to the element after the number added in that noble gas. • These shorthand notations ease identification of valence electrons (highest energy s and p electrons) Te 1s22s22p63s23p64s23d94p65s24d105p4= [Kr]5s24d105p4 Cl 1s22s22p63s23p5 = [Ne]3s23p5

32. Look at the Aufbau diagrams and confirm that you get these results! Homework: 11/4/09 Look at the aufbau diagrams Confirm results at this website http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html#aufbau

33. Class work 11/5 & 11/6 and Homework 11/6/09 Review to date (chapter 3-7, 3-8, notes, & power point lessons online at course webpage), work towards completing the webquest posted online at the course webpage this week The webquest provides important review (comprehensive for the topic of electron configuration and quantum mechanical model of the atom) for the unit test to be administered 11/23/09.

34. The position of the elemental symbol shows the location of its highest energy electrons S and p orbital electrons of the highest occupied ground state energy level are called valence electrons Another way to write an element’s electron configuration is shorthand (abbreviated) notation, showing only the highest s p d and f electrons and the symbol of the previous noble gas

35. Lesson for class on 11/9 and 11/10/09 and reinforced (Mr. Lis present) on 11/11/09 -Valence electron number and importance -Correlation between valence electron number and family/group # -Correlation between the valence energy level and the period # -Lewis dot structures of elements -Oxidation #s and correlations to valence electron number, electronegativity, & group/family # -Link between how each element reacts and its valence electron number

36. A Lewis dot structure (electron dot structure) shows only the valence electrons of an atom. In the s and p blocks, all members of the same group/family (columns) have the same number valence electrons, but these are in different energy levels (shells). In each period (row), all elements have valence electrons located in the same energy level (shell).

37. Lewis dot structures of the “representative elements”

38. Transition metals families may show variable # s orbital valence electrons Most members of d & f block families(called transition metals) also have same #s valence electrons. Exceptions occur in cases where skipping one or both valence s orbital electron positions allows for greater stability: • some d & f partial filled configurations are more stable than others • half & fully filled d & f orbitals are particularly stable. Mo [Kr] 5s1 4d5Ag [Kr] 5s1 4d10 Au [Xe] 6s1 4f14 5d10 Mo° Ag ° Au ° Not all tr metals violate filling rules Os [Xe] 6s24f145d6 Os:

39. Valence electrons control chemical properties The octet rule explains why: Atoms are most stable when they possess a filled valence (outermost) shell octet. (i.e., when their outermost energy level—highest energy level—contains a total of 8 electrons in s and p orbitals. Members of group 18 (the noble gases—aka inert gases) have highest stability and are the only atoms commonly found unbonded. Inert means nonreactive.

40. The behavior of atoms as they satisfy the octet rule is controlled by electronegativity. Metals have low electronegativity (property indicating how tightly an atom would pull electrons toward itself in a bond with another atom). Nonmetals have high electronegativity. So, metals form + ions by giving away valence shell electrons, exposing filled inner shell s & p orbitals. So, nonmetals form – ions by taking valence shell electrons to fill their existing valence shell.

41. The oxidation number of an element is the charge its atoms attain when forming stable ions that satisfy the octet rule.Metalloids alter strategy according to bonding partners. If bonding to an element with much lower electronegativity, the nonmetal will adopt the – oxidation number. If bonding to an element with much higher electronegativity, they will adopt the positive oxidation number. If bonding to a partner of similar electronegativity, they will form covalent bonds, sharing electrons.

42. Write the electron configuration, draw Aufbau diagrams, draw the lewis structure, and predict the most common oxidation # for the middle 3 members of families 1, 2, and 13—18. Highlight orbitals that will be used in bonding to fulfill the octet rule—if the orbitals will lose electrons to form + ions, highlight them blue, if they will gain electrons by ionic or covalent bonding, shade them yellow.On the second page for each element, For these same elements, write the electron configuration, draw aufbau diagrams, draw the Lewis structure, and predict the charge of a stable ion. (finish this for homework due 11/12, work on the webquest for homework—all of the concepts covered in it will be important for unit test before Thx break) Tomorrow Mr. Lis can help you with this assignment OR with the webquest.

43. Homework due 11/10/09 Please review the slides dated for chapter 3 Practice identifying/drawing the Lewis dot structure for an atom based on its element’s location on the periodic table; practice identifying the oxidation number based on periodic table location. By 11/12, Be able to write electron configurations & draw Aufbau (orbital) diagrams & Lewis dot structures for elements and their most stable ions. This material can also be reviewed on pages 107-108 and sample problem 3.12. By 11/13, complete the webquest, also.

44. Except for the noble gases, atoms need to either lose, gain, or share electrons to have a stable valence shell (highest energy level s & p orbital octet)

45. Higher electronegativity nonmetal Cl takes an electron from lower electronegativity metal Na

46. Electron configuration changes when ions form • + ions lose electrons from the s and p valence orbitals to “shed” the outer shell and reveal a stable shell underneath • - ions gain electrons to fill the remaining s and p orbitals of the incomplete valence shell • Why + ions? More p+ than e- • Why – ions? More e- than p+ • Why the same oxidation # per family/group? • Same # valence electrons

47. Electron Dot Structure

48. Ionic Bond: exchange of electrons • Transfer of electrons creates oppositely charged ions, which are attracted to each other • BOTH a gain and a loss of electrons occur in the formation of an ionic bond • Metals with low electronegativity donate electrons (all of their valence s and p electrons—don’t worry about the highest energy d and f electrons, as they’ll always be in a shell at least 1 energy lower than the valence shell!) to the nometals • Metal ions are + charged • The nonmetals complete the empty s and p orbitals of their valence shell. • Nonmetals are – charged

49. Lab—show only the protons (red) and electrons (gold, yellow) • Follow the rules of orbital filling • Aufbaus—lowest energy orbital fills first • Hunds—equal energy orbitals add one electron each before each gains a second electron • Pauli’s exclusion rule—electrons in the same orbital have opposite spins (gold vs yellow)

50. Concepts to check/reinforce with the atom boards By examining the atom boards, • confirm your count of valence electrons and their locations (highest energy s and p orbitals) • Check your Aufbau diagrams • Confirm your # lewis dots • Confirm your oxidation number (+ or – and how many charges) • Fill in the table and the orbital drawings! Note where valence electrons are located!