1. Periodic Trends Mrs.Kay

2. Groups: vertical columns (18) Groups: vertical columns (18) Have similar properties because have same number of electrons in outer shell Have similar properties because have same number of electrons in outer shell Periods: horizontal columns (7) Periods: horizontal columns (7)

3. Valence electrons: electrons found in the outer most shell or valence shell Valence electrons: electrons found in the outer most shell or valence shell Each energy level/shell holds on a certain number of electrons Each energy level/shell holds on a certain number of electrons

5. Alkali metals all end in s 1 Alkali metals all end in s 1 Alkaline earth metals all end in s 2 Alkaline earth metals all end in s 2 really have to include He but it fits better later. really have to include He but it fits better later. He has the properties of the noble gases. He has the properties of the noble gases. s2s2 s1s1 S- block

6. Transition Metals -d block d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10

7. The P-block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

8. F - block inner transition elements inner transition elements

9. Metals: To the left of the staircase line Physical Properties: Physical Properties:  Luster (shiny)  Good conductors  High density  High melting point  Malleable Chemical Properties: Chemical Properties:  Easily lose electrons  Corrode easily (ex: rusting or tarnishing)  Low electronegativity

11. Non metals: to the right of the staircase line Physical Properties: Physical Properties:  Dull  Poor conductor  Brittle  Not malleable  Low density and melting point Chemical Property: Chemical Property:  Tend to gain electrons  High electronegativity

12. Metalloids: along the staircase line Solids Solids Shiny or dull Shiny or dull Malleable Malleable Conduct heat and electricity better than non metal but not as well as metals Conduct heat and electricity better than non metal but not as well as metals

13. Atomic Size

14. The electron cloud doesn’t have a definite edge. The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time. They get around this by measuring more than 1 atom at a time. Summary: it is the volume that an atom takes up Summary: it is the volume that an atom takes up

15. Atomic Size Atomic Radius = half the distance between two nuclei of a diatomic molecule. Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

16. 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 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

17. 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. 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. NaMgAlSiPSClAr

18. Reactivity

19. Reactivity Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons

20. For Metals: Period - reactivity decreases as you go from left to right across a period. Group - reactivity increases as you go down a group Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity

21. For Non-metals Period - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group. Period - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group. Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron. Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

22. Ionization Energy

23. The amount of energy required to completely remove an electron from a gaseous atom. The amount of energy required to completely remove an electron from a gaseous atom. An atom's 'desire' to grab another atom's electrons. An atom's 'desire' to grab another atom's electrons. Removing one electron makes a +1 ion. Removing one electron makes a +1 ion. The energy required is called the first ionization energy. The energy required is called the first ionization energy. X (g) + energy → X + + e- X (g) + energy → X + + e-

24. Ionization Energy The second and third ionization energies can be represented as follows: X + (g) + energy  X 2+ (g) + e- X + (g) + energy  X 2+ (g) + e- X 2+ (g) + energy  X 3+ (g) + e- X 2+ (g) + energy  X 3+ (g) + e- More energy required to remove 2 nd electron, and still more energy required to remove 3 rd electron More energy required to remove 2 nd electron, and still more energy required to remove 3 rd electron

25. Group trends Ionization energy decreases down the group. Ionization energy decreases down the group. Going from Mg to Be, IE decreases because: Be outer electron is in the 3s sub-shell rather than the 2s. This is higher in energy Be outer electron is in the 3s sub-shell rather than the 2s. This is higher in energy The 3s electron is further from the nucleus and shielded by the inner electrons The 3s electron is further from the nucleus and shielded by the inner electrons So the 3s electron is more easily removed So the 3s electron is more easily removed A similar decrease occurs in every group in the periodic table. A similar decrease occurs in every group in the periodic table.

27. Period trends IE generally increases from left to right. Why? Why? From Na to Ar (11 protons to 18 protons), the nuclear charge in each element increases. From Na to Ar (11 protons to 18 protons), the nuclear charge in each element increases. The electrons are attracted more strongly to the nucleus – so it takes more energy to remove one from the atom. The electrons are attracted more strongly to the nucleus – so it takes more energy to remove one from the atom.

29. Why is there a fall from Mg to Al? Al has configuration 1s2 2s2 2p6 3s2 3p1, its outer electron is in a p sublevel Al has configuration 1s2 2s2 2p6 3s2 3p1, its outer electron is in a p sublevel Mg has electronic configuration 1s2 2s2 2p6 3s2. Mg has electronic configuration 1s2 2s2 2p6 3s2. The p level is higher in energy and with Mg the s sub level is full – this gives it a slight stability advantage The p level is higher in energy and with Mg the s sub level is full – this gives it a slight stability advantage

30. Why is there a fall from P to S? This can be explained in terms of electron pairing. This can be explained in terms of electron pairing. As the p sublevel fills up, electrons fill up the vacant sub levels and are unpaired. As the p sublevel fills up, electrons fill up the vacant sub levels and are unpaired. This configuration is more energetically stable than S as all the electrons are unpaired. It requires more energy to pair up the electrons in S so it has a lower Ionisation energy. This configuration is more energetically stable than S as all the electrons are unpaired. It requires more energy to pair up the electrons in S so it has a lower Ionisation energy. There is some repulsion between the paired electrons which lessens their attraction to the nucleus. There is some repulsion between the paired electrons which lessens their attraction to the nucleus. It becomes easier to remove! It becomes easier to remove!

31. Driving Force Full Energy Levels are very low energy. Full Energy Levels are very low energy. Noble Gases have full energy levels. Noble Gases have full energy levels. Atoms behave in ways to achieve noble gas configuration. Atoms behave in ways to achieve noble gas configuration.

32. 2 nd Ionization Energy For elements that reach a filled or half filled sublevel by removing 2 electrons 2 nd IE is lower than expected. For elements that reach a filled or half filled sublevel by removing 2 electrons 2 nd IE is lower than expected. Makes it easier to achieve a full outer shell Makes it easier to achieve a full outer shell True for s 2 True for s 2 Alkaline earth metals form +2 ions. Alkaline earth metals form +2 ions.

33. 3 rd IE Using the same logic s 2 p 1 atoms have an low 3 rd IE. Using the same logic s 2 p 1 atoms have an low 3 rd IE. Atoms in the aluminum family form +3 ions. Atoms in the aluminum family form +3 ions. 2 nd IE and 3 rd IE are always higher than 1 st IE!!! 2 nd IE and 3 rd IE are always higher than 1 st IE!!!

34. Electron Affinity

35. The energy change associated with adding an electron to a gaseous atom. The energy change associated with adding an electron to a gaseous atom. Easiest to add to group 7A. Easiest to add to group 7A. Gets them to full energy level. Gets them to full energy level. Increase from left to right atoms become smaller, with greater nuclear charge. Increase from left to right atoms become smaller, with greater nuclear charge. Decrease as we go down a group. Decrease as we go down a group.

36. Ionization energy, electronegativity Electron affinity INCREASE

37. Atomic size increases, shielding constant Ionic size increases

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

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

40. Configuration of Ions Ions always have noble gas configuration. Ions always have noble gas configuration. Na is 1s 2 2s 2 2p 6 3s 1 Na is 1s 2 2s 2 2p 6 3s 1 Forms a +1 ion : 1s 2 2s 2 2p 6 Forms a +1 ion : 1s 2 2s 2 2p 6 Same configuration as neon. Same configuration as neon. Metals form ions with the configuration of the noble gas before them - they lose electrons. Metals form ions with the configuration of the noble gas before them - they lose electrons.

41. Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them. They end up with the configuration of the noble gas after them.

42. Periodic Trends Across the period nuclear charge increases so they get smaller. Across the period nuclear charge increases so they get smaller. Energy level changes between anions and cations. Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

43. Size of Isoelectronic ions Iso - same Iso - same Iso electronic ions have the same # of electrons Iso electronic ions have the same # of electrons Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 Al +3 Mg +2 Na +1 Ne F -1 O -2 and N -3 all have 10 electrons all have 10 electrons all have the configuration 1s 2 2s 2 2p 6 all have the configuration 1s 2 2s 2 2p 6

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

45. Electronegativity (optional coverage)

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

47. Group Trend The further down a group the farther the electron is away and the more electrons an atom has. 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. 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.

48. Period Trend Electronegativity increases from left to right across a period 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 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

49. 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. 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.

50. 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. 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. 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.

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