1. 3.1 The periodic table 3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number 3.1.2 Distinguish between the terms group and period 3.1.3 Apply the relationship between the electron arrangement of elements and their position in the periodic table up to z=20. 3.1.4 Apply the relationship between the highest occupied energy level for an element and its position in the periodic table.

2. Groups: vertical columns (18) Have similar properties because have same number of electrons in outer shell Periods: horizontal row (7) Family Names: Group 1: alkali metals Group 2: alkaline earth metals Group 17: halogens Group 18: noble gases Group 3-12: Transition metals Groups 1,2, 13-18: representative elements

5. Atomic Size The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time. Summary: it is the volume that an atom takes up http://www.mhhe.com/physsci/chemistry/essentialch emistry/flash/atomic4.swf http://www.mhhe.com/physsci/chemistry/essentialch emistry/flash/atomic4.swf

6. Group trends As we go down a group (each atom has another energy level) the atoms get bigger, because more protons and neutrons in the nucleus H Li Na K Rb

7. Trends Within Groups (Families) Increase as you move down a group. Even though nuclear charge increases as you go down a group, the orbital sizes increase so much that the atom becomes larger. The outer electrons are farther from the nucleus and are shielded from the positive charge of the nucleus by the other electrons.

8. Periodic Trends atomic radius decreases as you go from left to right across a period. Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter. Remember filling up same energy level, little shielding occurring. NaMgAlSiPSClAr

9. Ionization Energy Ionization Energy: The energy needed to overcome the attraction between the positive charge in the nucleus and the negative charge of the electron. How much energy is needed to remove an electron from an atom. Energy is measured in Joules.

10. Ionization Energy High ionization energy values indicate a strong hold on electrons. Unlikely to become a positive ion. Low ionization energy values indicate a weak hold on electrons. Likely to become positive ions.

11. Ionic Size Cations form by losing electrons. Cations are smaller than the atom they come from. Metals form cations. Cations of representative elements have noble gas configuration.

12. Ionic Size Anions form by gaining electrons. Anions are bigger than the atom they come from. Nonmetals form anions. Anions of representative elements have noble gas configuration.

13. Periodic Trends Metals losing from outer energy level, more protons than electrons so more pull, causing it to be a smaller species. Non metals gaining electrons in its outer energy level, but there are less protons than electrons in the nucleus, so there is less pull on the protons, so found further out making it larger. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

14. Why do positive ions become smaller? Two Reasons: The electron lost from the atom will always be a valence electron = smaller radius. The lost electron no longer shields the other electrons from the positive nucleus, so they are pulled closer to the nucleus.

15. Ionic Radii Atoms that gain electrons always become larger. Why? Additional electron causes other orbitals to be filled. Increased shielding causes other electrons to be farther away from the nucleus.

16. Ionic Radii http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/medialib/media_portfolio/text_images/CH08/FG08_13.JPG

17. Ionic Radii: Trends What is the pattern in Periods? Smaller until 5A. What is the pattern in Groups? Gradual increase in size.

18. Size of Isoelectronic ions Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

19. Ionization Energy http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/images/ionization-energy.jpg

20. Ionization Energy http://images.google.com/imgres?imgurl=http://www.webelements.com/webelements/properties/media/tables/cityscape-x/ionization-e nergy-1.jpg&imgrefurl=http://www.webelements.com/webelements/properties/text/image-cityscape/ionization-energy-1.html&h=1365& w=2048&sz=808&tbnid=lq6myNPljopuCM:&tbnh=100&tbnw=150&prev=/images%3Fq%3Dionization%2Benergy&start=1&sa=X&oi=im ages&ct=image&cd=1

21. Can I Remove More Than One Electron? A second, third, etc, electron can be removed from an atom. The ionization energies are termed accordingly: 2 nd Ionization energy to remove the 2 nd electron. 3 rd Ionization energy to remove the 3 rd electron.

22. 2 nd and 3 rd Ionization Energies Do you think they are higher values or lower values than the 1 st Ionization energy? Usually the values are higher since the atom holds onto the remaining electrons even tighter.

23. 12345678 H1312 He23725250 Li520729711810 Be89917571484521000 B800242636592502032820 C10862352461962213782047260 N140228554576747394425325064340 O131433885296746710987133207132084070 F168033756045840811020151601786092010 Ne20803963613093611218015240 Na49645636913954113350166002011325666 Mg737145077311054513627179952170025662 http://www.shodor.org/chemviz/ionization/students/background.html Ionization Energies in kJ/mol

24. Trends of Ionization Energy Within Periods: Increase as you move left-to-right. Due to increase in nuclear charge and a tight hold on electrons. Within Groups: Generally decreases as you move down a group. Electrons are farther away from nucleus.

25. Ionization Energy What happens when sodium loses an electron? What is its electron configuration? Na 1s 2 2s 2 2p 6 3s 1 Na +1 1s 2 2s 2 2p 6 The octet rule states that atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons.

27. Yet Another Trend! Electronegativity: ability of an element to attract an electron in a chemical bond. How badly does it want another electron? http://college.hmco.com/chemistry/intro/zumdahl/intro_chemistry/5e/students/protected/periodictables/pt/pt/table/t_e2.gif Why no Values?

28. Electronegativity The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How fair it shares. Big electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity.

30. Group Trend The further down a group the farther the electron is away and the more electrons an atom has. So as you go from fluorine to chlorine to bromine and so on down the periodic table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus. For that reason the electronegativity decreases as you go down the periodic table.

31. Period Trend Electronegativity increases from left to right across a period When the nuclear charge increases, so will the attraction that the atom has for electrons in its outermost energy level and that means the electronegativity will increase

32. Period trend Electronegativity increases as you go from left to right across a period. Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

33. Group Trend electronegativity decreases as you go down a group. Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

34. Shielding Shielded slightly from the pull of the nucleus by the electrons that are in the closer orbitals. Look at this analogy to help understandanalogy

35. Electronegativity Trends

36. Overall Trends! http://campus.ru.ac.za/full_images/img05206111510.jpg

37. Melting Points of Group 1 ElementMelting Point (K) Li453 Na370 K336 Rb312 Cs301 Fr295

38. Metallic bonding Collective bond, not a single bond Strong force of electromagnetic attraction between delocalized electrons (move freely). This is sometimes described as "an array of positive ions in a sea of electrons

39. Why does the melting point decrease going down the alkali metals family? Atoms are larger and their outer electrons are held farther away from the positive nucleus. The force of attraction between the metal ions and the sea of electrons thus gets weaker down the group. Melting points decrease as less heat energy is needed to overcome this weakening force of attraction.

40. Melting Points for halogens ElementMelting Point (K) Fluorine85 Chlorine238 Bromine332 Iodine457 Astatine610

42. Why does melting point increase going down the halogens? The halogens are diatomic molecules, so F 2, Cl 2, Br 2, I 2 As the molecules get bigger there are more electrons that can cause more influential intermolecular attractions between molecules. The stronger the I.A, the more difficult it will be to melt. (more energy needed to break the I.A)

43. What are these I.A? van der Waals forces (London dispersion): Electrons are mobile, and although in a diatomic molecule they should be shared equally, it is found that they temporarily move and form slightly positive end and negative end. Now that one end is + and the other -, there can be intermolecular attractions between the opposite charges of the molecules

44. van der Waals forces

45. IB requires knowledge specifically for halogens. Check out this site for more detail. http://www.chemguide.co.uk/inorganic/group7/propert ies.html

46. Period 3 melting point trends

47. Explanation M.P rise across the 3 metals because of the increasing strength of the metallic bonds. Silicon has a giant covalent structure just like diamond which makes its structure remarkably strong and therefore takes more energy to break apart.

48. The atoms in each of these molecules are held together by covalent bonds (except Ar) They would have weak I.A affecting the amount of energy needed to melt them. Ar has extremely weak forces of attraction between its atoms, so its easiest to melt.

49. 3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group. 3.3.2 Discuss the changes in nature from ionic to covalent and from basic to acidic of the oxides across period 3

50. Reactivity of alkali metals Generally group 1 metals become more reactive as you go down a group. The valence electron of group 1 are found further from the nucleus as you go down the group. It is easier to remove an electron from francium than from lithium

51. Alkali metal + water Li(s) + H 2 O (l)  LiOH(aq) + H 2 (g) (Li + and OH - in solution) The metal reacts with water to form the hydroxide of the metal (strong base) and bubbles off hydrogen gas. The larger the alkali metal, the more vigorous the reaction. Sometimes the H 2 gas actually lights itself (exothermic reaction, releases heat) causing the H 2 to burn.

52. MUST KNOW! Na (s) + H 2 O (l)  NaOH (aq)+ H 2 (g) K (s) + H 2 O (l)  KOH (aq)+ H 2 (g)

53. Alkali metals + halogens 2Na (s) + Cl 2 (g)  2NaCl (s) Halogens are good oxidizing agents, which means they cause electrons to be lost from another atom (the reducing agent) Halogens are 1 electron from stable octet and will try to remove electrons from valence electrons of other metallic atoms.

54. MUST KNOW! 2K (s) + Br 2 (l)  2KBr (s) 2Li (s) + I 2 (g)  2LiI (s)

55. Halogens reacting with halides Halogens want an electron and even will remove electrons from other soluble salts, we refer to as halides. When a salt dissolves it forms both of its ions in solution. Ex: NaCl (aq)  Na + (aq) and Cl - (aq) So halides are easily available for reactions

56. Done in aqueous systems Chlorine is stronger OA (oxidizing agent) than bromine because its found higher on the periodic table, so Cl 2 will remove the electron from Br -, making Cl - and Br 2 Cl 2 (aq) + 2Br -  2Cl - + Br 2 (aq) Cl 2 (aq) + 2I -  2Cl - + I 2 (aq) Br 2 (aq) + 2I -  2Br - + I 2 (aq)

57. Properties of Metals Shiny (lustre) Good conductors of heat and electricity Malleable and ductile (change shape and make wires) Tend to lose electrons Metal oxides form basic solutions in water (pH greater than 7)

58. Properties of non-metals Brittle Poor conductors of heat and electricity Tend to gain electrons Non-metal oxides tend to be basic when dissolved in water (pH less than 7)

59. Across Period 3: metallic to non-metallic oxides Basic solution from metallic oxide. Na 2 O(s) + H 2 O (l)  2 NaOH (aq) MgO (s) +H 2 O (l)  Mg(OH) 2 (aq) Hydroxides of group 1 and 2 generally considered strong. Acidic solution from non-metallic oxide. SO 3 (g) + H 2 O (l)  H 2 SO 4 (aq) P 4 O 10 (s) + 6H 2 O (l)  4 H 3 PO 4 (aq) Aqueous hydrogen involved with acidity

60. Properties of metalloids Based on chemical and physical properties Tend to have semi-conductive properties and form amphoteric oxides. Considered metalloids are: Boron (B) Boron Silicon (Si) Silicon Germanium (Ge) Germanium Arsenic (As) Arsenic Antimony (Sb) Antimony Tellurium (Te) Tellurium Polonium (Po Polonium

61. Amphoteric Behave as an acid or a base depending upon the reaction it is involved with. Also called amphiprotic (donate or accept a proton, H + ) Aluminum’s oxide is amphoteric. Al 2 O 3 (s) + 3HCl (aq) → AlCl 3 (aq)+ 3H 2 O (l) Reacts with a strong acid to make a to make a salt with water. Al 2 O 3 (s)+ NaOH (aq) → NaAl(OH) 4 (aq) Reacts with a strong base to form sodium aluminate