Mr Shields Regents Chemistry U 08 L 04

Mr. Shields Regents Chemistry U 08 L 04 1

Size Trends Atomic Radii follows two trends: 1) Radii increases going down a group 1) Radii decreases going across a period But how do we measure Atomic Radii? 2

Atomic Radii Atomic Radius is measured as ½ the distance between Adjacent nuclei in a molecule. Why not just measure ½ the diameter of the atom? Hint: What’s the definition of an atomic orbital? 3

Atomic Radii Trends Atomic Radii increasing Decreasing radii How can we account for this trend? 4

Atomic Radii The trend down a group may be easier to explain first As we move down a group what happens to the principal Energy level ? As the principal energy level increases, electrons move further and further from the nucleus. Nucleus n=1 2 3 4 5

Atomic Radii in Groups As we move down group 1 each successive S orbital electron Is further from the nucleus and thus Atomic radii increases n Element Radii (pm) change 2 Li 155 - 3 Na 190 35 4 K 235 45 5 Rb 248 13 6 Cs 267 19 7 Fr 270 3 6

Atomic Radii in groups Since the s orbital is further from the nucleus the radii of the Atom increases. But … The nuclear charge is also increasing since Atomic Number is increasing. Increasing nuclear charge Diminishes the rate of Change of increasing Atomic Radii Down a Group. Electrons are pulled Toward the nucleus more strongly Nuclear charge Radii (pm) change Li (3) 155 - Na (11) 190 35 K (19) 235 45 Rb (37) 248 13 Cs (55) 267 19 Fr (87) 270 3 7

Atomic Radii across periods We’ve seen how atomic radii increases going down a group But what happens when we go across a period? We’ll, in fact atomic radii decreases. But why? We can begin to understand what is happening if we look at Both the Atomic number and what principle energy level electrons are being added to. 8

Atomic Radii across periods As we move across a period Atomic numbers increase - Pos. Nuclear charge also increases so would expect the electrons to be pulled closer to the nucleus. So this could explain decreasing Atomic radius - BUT … this same thing happens as we move down a group. And for groups Atomic Radii increases as we add more electrons? So why does radius increase in groups but not across a period? 9

Atomic radii across periods The difference is that when we go across a period electrons do Not fill higher energy levels. They either occupy lower energy Levels or the same energy level Whereas when going down a group electrons occupy successively Higher principle energy levels. For example … n=3 n=4 n=5 1 2 -8 -8 -1 2 -8 -18 -8 -1 group 2 2 -8 -8 -2 - 3 2 -8 -9 -2 - 13 2 -8 -18 -3 - But why doesn’t atomic radii remain about the same across a group? 10

Atomic radii across periods As atomic number increases across a row additional electrons are added to the same (or lower) energy level. The effective Nuclear charge may also be increasing and electrons are pulled in More strongly towards the larger more positive nucleus. +11 2 -8 -1 e- Na: Effective nuclear charge = +1 +16 2 -8 -6 e- S: Eff. Nuclear Charge = +6 Unlike when moving down a group, there are no new principal energy levels being added to counteract the effect of increasing nuclear charge and increasing effective nuclear charge. 11

Nuclear charge (K) = 19 Effective Nuc. Chg. (Grp I) = 1 Nuclear charge (Br) = 35 Effective Nuc. Chg. (Grp VII)= +7 12

Ionic Radii We’ve now seen how atomic radii changes in Periods & Groups But what happens when Atoms either gain or lose electrons to form ions? How does Ionic Radii vary down Groups & across Rows? Remember … The representative elements of groups 1, 2, 13 and 14 give up electrons to form +1, +2, +3 and +4 ions respectively 13

Positive Ions When atoms lose all their valence electrons they lose the outermost quantum level (n). Consider Aluminums electron configuration. What is it? 2 -8 -3 (principle energy levels 1, 2 and 3 are occupied) What is the electron config after Al loses its 3 valence electrons? 2 -8 (only principle energy levels 1 and 2 are occupied) What is the charge on Aluminum? 14

Positive Ions The loss of the outermost valence shell has two effects: 1) The atoms radius shrinks because it loses it’s outermost principle quantum level AND … 2) The Nucleus now has more positive charge than the total negative charge from electrons. The larger effective nuclear charge will now pull electrons in closer to the nucleus 15

Positive Ions Notice that even though the ionic electron config is the same ionic radius gets progressively smaller moving across the period. Na Mg Al Atomic No. 11 12 13 Ionic charge +1 +2 +3 Atomic Radius (pm) 190 160 143 Ionic Radius (pm) 99 65 50 Ionic Elec. Config 2 -8 2 -8 This happens because the positive eff. nuclear charge seen by the same number of electrons increases as we move across a the row 16

Variation in atomic and Ionic Radii What’s going on Here? 17

Negative Ions Let’s next look at the non-metals, for example Chlorine Non-metals form ions by gaining electrons Cl 2 -8 -7 2 -8 -8 Cl- (negative ion) When we add electrons the effective nuclear charge per electron decreases AND there is increases electron repulsion So … you would expect the ionic radius to increase and it does Cl atomic radius = Cl- ionic radius = 99 nm 181 nm 18

Negative Ions Moving down groups the principal energy level increases - This is true for all atoms, Anions (-) & Cations (+) Li 2 -1 Na 2 -8 -1 K 2 -8 -8 -1 Li+ 2 Na+ 2 -8 K+ 2 -8 -8 F 2 -7 Cl 2 -8 -7 Br 2 -8 -18 -7 FCl. Br- 2 -8 -8 2 -8 -18 -8 So atom & ionic size increases going down Groups Going across periods Ionic size first decreases then jumps up When Oxidation states change from positive to negative. - after the jump up the downward trend in size continues 19

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The ionic compound Mg. O Mg atom O atom Electron Config 2 -8 If we were to look at individual atoms Mg would Actually be larger than Oxygen! 21