Chemical Periodicity Trends in the periodic table

Atomic Size

 How do you measure the size of an atom?  The electron cloud doesn’t have a definite edge.  Can get around this by measuring covalent atomic radius.

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

Atomic size is influenced by two factors:  Energy Level –more occupied levels = bigger atom  Charge on nucleus –More charge pulls electrons in closer

Group trends  As we go down a group electrons are added to higher energy levels so the atoms get bigger. H Li Na K Rb

Periodic Trends  As you go across a period the radius gets smaller.  Same energy level, but protons pull electrons closer to nucleus. NaMgAlSiPSClAr

Trends in Atomic Radius

Questions: Of the following elements, which has the largest atomic radius? Why? a) Si, Mg, S b) Al, Na, Cl c) Li, Cs Mg – same energy level, smallest nuclear charge Na – same energy level, smallest nuclear charge Cs – higher occupied energy levels

Ionic Size

Cations  Positive ions - form by losing electrons.  Metals form cations  Cations of representative elements have noble gas configuration.  Smaller than the atom they come from because of increased attraction by nucleus for fewer remaining electrons +

Ionic size Anions  Negative ions - form by gaining electrons.  Nonmetals form anions.  Anions of representative elements have noble gas configuration.  Larger than the atom they come from, because nuclear attraction is less for an increased number of electrons. -

Group trends  Ions get bigger as you go down (adding energy levels) Li +1 Na +1 K +1 Rb +1 Cs +1

Periodic Trends  Across the period nuclear charge increases so both cations and anions get smaller from left to right. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

Questions: 1. Of the following ions, which ones should have the larger radius? Why? a) Na + or Cs + b) Br - or K + 2. The Mg 2+ and Na + ions have ten electrons surrounding the nucleus. Which ion would you expect to have the smaller radius? Why? Cs + It has more occupied energy levels Br - Anions are larger than cations Mg 2+ Greater nuclear charge

Ionization Energy

 The amount of energy required to completely remove an electron from a gaseous atom (how hard it is to pull an e - off an atom)  1 st IE = removing 1 e -, 2 nd IE=removing 2 e - Na (g) Na + + e -

Shielding  The electron on the outside energy level is shielded from the nucleus by the inner electrons

Group trends u As you go down a group first IE decreases because the electron is further away (more shielding)

Periodic trends  All the atoms in the same period have the same energy level (same shielding).  As you go from left to right, nuclear charge increases so IE generally increases.

Questions: 1. Which element in the following sets has the lowest ionization energy and why? a)B, C, F b)K, Na, Li B – same energy level, smallest nuclear charge K – electron farther away, more shielding

Electron Affinity

 The energy given off when an electron is added to an atom  how much an atom ‘wants’ an electron F (g) + e - F - (g)

Electron Affinity Group trends  Generally decreases as we go down a group because shielding increases Periodic trends  Increases from left to right as atoms become smaller with greater nuclear charge

Questions: 1. Of the following elements, which ones should have the higher electron affinity? Why? a)Se or Te b)Calcium or Chromium Se – smaller atom Chromium – greater nuclear charge

Electronegativity

 The tendency for an atom to attract electrons to itself when it is chemically combined (BONDED) with another element.  Big electronegativity means it pulls the electron towards itself.

Group Trends  The further down a group the farther the electron is away from the nucleus and the more electrons an atom has.  More willing to share = low electronegativity  So as you go down a group electronegativity decreases

Periodic Trends  As we go from left to right across the table, electronegativity increases, because nuclear charge is increasing and electrons are held in more strongly  Metals have low electronegativity  Non-metals have high electronegativities (they win the electron tug-of-war)

Questions: 1. Which element would you expect to have the highest electronegativity? Why? 2. Put the following elements in order of increasing electronegativity: Na, P, Cl F smallest nonmetal Na, P, Cl