2.  Periodic table: trends in chemical and physical properties that occur within the same groups and periods  First attempt (Mendeleev and Meyer) elements arranged in order of increasing atomic weight  Modern periodic table: elements arranged in order of increasing atomic number

3.  Effective nuclear charge (Z eff ): positive charge felt by a valence electron in a many-electron atom  Depends on its distance from the nucleus and the number of core electrons  Core electrons block, or screen, valence e - from the full attraction of the nucleus  As average number of screening/core electrons (S) increases, effective nuclear charge decreases  As distance from the nucleus increases, S increases and Z eff decreases

4.  Mg: [Ne]3s 2  Z = 12  S = core e - = 10  Z – S = 12 – 10 = 2

5.  Effective nuclear charge… › Increases as you move across a period  Number of core e-, S, remains the same  Atomic number increases; higher nuclear charge, Z  Valence e- added to counter charge of p+, but valence e- don’t shield one another › Decreases as you go down a group  Large e- cores (think atomic radius) aren’t able to screen valence e- from the nuclear charge

6.  Bounce two of the same atoms off each other (B)  The distance between two nuclei is the nonbonding atomic radius (aka: van der Waals radius) + - - - - -+ - - - - - B rad

7.  Homonuclear diatomic molecule  Distance between 2 nuclei is bond length  Half the bond length is the bonding atomic radius

8.  Decrease in atomic radius across a period › Effective nuclear charge increases, drawing the electrons in towards the nucleus very tightly  Increase in atomic radius down a group › As the principal quantum number increases, the probability of finding the electrons further from the nucleus increases

10.  Arrange the following atoms in order of increasing atomic size: › Mg, Ca, Sr › B, F, Ge, Pb

11.  Ionic radius: distance between ions in an ionic compound  When atoms form ions, they tend to achieve a full valence shell by either losing e- or gaining e-  Cations are smaller than parent atom › Outermost orbital is emptied - decreasing radius of the ion  Anions are larger than parent atom › Outermost orbital fills up, taking up more space

13.  For ions of the same charge, ion size increases down a group  Ion size decreases across a period  ASIDE: Transition metals lose electrons in outermost orbital first! › Ex.: Fe(3d 6 4s 2 )  Fe 3+ (3d 5 ) + 3e -

14.  Isoelectric series: group of ions containing the same number of electrons  Ex: Se 2-, Br -, Rb +, Sr 2+, Y 3+ › Se 2- = atomic # 34 (36 e - ) › Br - = atomic # 35 (36 e - ) › Rb + = atomic # 37 (36 e - ) › Sr 2+ = atomic # 38 (36 e - ) › Y 3+ = atomic # 39 (36 e - )  Ionic size in isoelectric series decreases with increasing atomic number due to increasing effective nuclear charge

15.  Ionization energy (I): minimum energy required to remove an e - from a gaseous atom or ion › The larger the ionization energy, the harder it is to remove the e -  First ionization energy (I 1 ): energy needed to remove 1 st e- from neutral atom › Na (g)  Na + (g) + e - (g) I 1 = +496 kJ/mol  Second ionization energy (I 2 ): energy needed to remove 2 nd e - › Na + (g)  Na 2+ (g) + e - (g) I 2 = +4562 kJ/mol

16.  Each successive ionization energy is higher than the previous  Sharp increase in ionization energy when a core e - is removed

17.  Decreases as you move down a group and atomic number gets larger › Outermost e - is easier to remove as you go down a group and atom gets bigger  Increases as you move across a period › Outermost e - is harder to remove as you go across a period and atom gets smaller

19.  Electron affinity (E): energy change associated with adding an e - to a gaseous atom or ion › I (g) + e - (g)  I - (g) ΔE = -295 kJ/mol  For most atoms, energy is usually released when an e - is added › Negative value

20.  Puts an e- in a new (higher energy) subshell (unfavorable)  Ne (g) = e - (g)  Ne - (g) ΔE > 0  To determine whether e - affinity is (+) or (-), look at e - configuration: › Do you have to add the e - to a higher energy orbital or subshell? › Ne = 1s 2 2s 2 2p 6 › Ne - = 1s 2 2s 2 2p 6 3s 1

22. Ionization Energy Electron Affinity  How easily an atom or ion loses an e -  Forms consecutively more (+) species  Measure of how tightly an e - is held on to  Usually requires energy [(+) value]  How easily an atom or ion gains an e -  Forms consecutively more (-) species  Measure of attraction for an outside e -  Usually releases energy [(-) value]

23.  What is the ionization energy for F - (g) if the electron affinity for F (g) is -328 kJ/mol?  Which member of each ion pair has the higher ionization energy? Why? › Sodium or rubidium › Silicon or phosphorous  Which element has the higher electron affinity, sulfur or chlorine?

24. Metals Nonmetals  Shiny luster  Various colors, but most are silvery  Malleable and ductile  Good conductors  Metal oxides form bases in H 2 O  Form cations in solution  All are solids at RT (except Hg)  No luster  Various colors  Brittle  Poor conductors  Nonmetal oxides form acids in H 2 O  Form anions in solution  Solids, gases, and liquid at RT

25.  Metals form cations  Low ionization energies  Metallic character: › Increases down a group › Decreases across a period

27.  When metals are oxidized they form characteristic cations: › Group 1A metals:  1+ ions › Group 2A metals:  2+ ions › Most transition metals have variable charges

29.  Form anions › When nonmetals react with metals, nonmetals gain e- and metals lose e-  Metal + Nonmetal  Salt  2Al (s) + 3Br 2(l)  2AlBr 3(s)  When nonmetals are reduced they form characteristic anions › Group 7A nonmetals  1- ions › Group 6A nonmetals  2- ions › Group 5A nonmetals  3- ions

30.  Soft  Chemistry dominated by loss of single e- › M  M + + e -  Reactivity increases down the group b/c easier to lose e-  Alkali metals react with water to form MOH and H 2 gas › 2M (s) + 2H 2 O (l)  2MOH (aq) + H 2(g)

31.  Alkali metal ions are reduced to metal gas atoms that emit characteristic colors when placed in a high temp flame  “s” e - is excited by flame and jumps to “p” sublevel and then emits light energy when it returns to ground state Na line (589 nm): 3p  3s transition Li line: 2p  2s transition K line: 4p  4s transition

32.  Harder and more dense  Chemistry dominated by loss of 2 e- › M  M 2+ + 2e -  Reactivity increases down the group b/c easier to lose e-

33.  Very unique element that forms › Colorless diatomic gas, H 2 › Metallic solid at high pressures › Can behave as a cation or anion: 2Na (s) + H 2(g)  2NaH (s) less common (H - ; hydride) H 2(g) + Cl 2(g)  2HCl (g) most common (H + ; proton)

34.  Metallic character increases down the group › O 2 colorless gas; others are solids › Te is a metalloid, Po is a metal, etc.  Gain e- to form 2- anions  2 important forms of oxygen: O 2 & O 3 3O 2(g)  2O 3(g)  H = +284.6 kJ  O 2 : potent oxidizing agent › Two oxidation states for oxygen: › 2- (Ex. H 2 O) › 1- (Ex. H 2 O 2 )

35.  All form diatomic molecules › F 2 and Cl 2  gases; Br 2  liquid; I 2  solid  Gain an e - to form an anion X 2 + 2e -  2X -  Fluorine is one of the most reactive substances known 2F 2(g) + 2H 2 O (l)  4HF (aq) + O 2(g)  H = -758.7 kJ  All react with metals to produce ionic halide salts Mg (s) + Cl 2(g)  MgCl 2(s)  All react with hydrogen gas to form gaseous hydrogen halide compounds Cl 2(g) + H 2(g)  2HCl (g)

36.  All monatomic  Notoriously unreactive  Have completely filled s and p subshells  In 1962 the first compound of the noble gases was prepared: XeF 2, XeF 4, and XeF 6 › Why these particular compounds?