The Periodic Table

Force of Attraction:

Valence Electrons (Outer-Shell Electrons)  Electrons that can participate in the formation of chemical bonds.  Electrons in the outermost “s” and “p” orbitals.  The number of valence electrons corresponds to the group number!  Transition & Rare Earth elements are tricky!  Sometimes their highest level “d” and “f” electrons behave like valence electrons and sometimes they behave like shielding electrons.

Shielding Electrons Core Electrons (Inner-Shell Electrons)  Electrons included in a noble gas “core” in the electron configuration.  They are between the nucleus and the valence electrons.  The are called shielding electrons because they shield the valence electrons from the attractive force of the nucleus.

Nuclear Charge  In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons.  The nuclear charge is equal to the number of protons in the nucleus or to the atomic number (Z)

Effective Nuclear Charge  Net positive charge felt by an electron.  Takes into account the core electrons which shield (or screen or cancel out) some of the nuclear charge.

Effective Nuclear Charge Z eff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons Example: Na: 1s 2 2s 2 2p 6 3s 1 Z=11, S=10, Z eff = +1 *This is an approximation – there are better methods beyond the scope of this course – this calculation is close enough for us!

Z eff Across a Series LiBeBCNOFNe Z S Z eff Z eff increases across a series The electrons are held closer to the nucleus are held more tightly by the nucleus as we move across a series.

Z eff Down a Group ZSZ eff Li32+1 Na K Rb Cs We calculate Z eff to be fairly constant going down a group, however it actually increases slightly because the core electrons are spread out over a larger space and therefore don’t screen the valence electrons as well.

Size of Atoms: Across a Series  Atoms decrease in size as electrons are being added to the same energy level.  Z eff increases so the force on the valence electrons gets stronger and they are pulled closer to the nucleus as we move across a series. ATOM Na Mg Al Si P S Cl Ar SIZE (radius in Å )

Size of Atoms: Down a Group  Atoms increase in size as the electrons go into a higher energy level (get farther away from the nucleus) ATOM SIZE (radius in Å ) H.32 Li 1.23 Na 1.54 K 2.03 Rb 2.16 Cs 2.35

Size of Ions  A positive ion (cation) is smaller than its atom – electrons leave from the outermost orbitals, electron-electron repulsion decreases.  A negative ion (anion) is larger than its atom – adding electrons causes more electron-electron repulsion causing the electrons to spread out

IONIZATION ENERGY (I.E.)  Low ionization energy is good for making ions!  The First Ionization Energy, I 1, is the energy required to remove the first electron from the atom  The Second Ionization Energy, I 2, is the energy required to remove the 2 nd electron  This continues for the successive removal of electrons. I 1 < I 2 < I 3, etc. because each electron removed is pulled from a more positive ion.

ACROSS A SERIES  I 1 increases because of the increased attraction between the nucleus and the electrons (higher Z eff ). ATOM Na Mg Al Si P S ClAr I.E (kcal/mole)

ACROSS A SERIES ATOM Na Mg Al Si P S ClAr I.E (kcal/mole) It’s not a perfect trend because of more stable filled & half filled orbitals….. The 3p sublevel has a higher energy than 3s so it’s easier to take an electron from the 3p (Al) than from the 3s (Mg) Sulfur’s electron is removed from an orbital with paired electrons which have some electron-electron repulsion. Phosphorus’s electron is removed from an orbital with a single electron.

DOWN A GROUP  Ionization energy decreases as the electrons go into higher energy levels which are farther away from the nucleus. ATOMI.E. (kcal/mole) H314 Li124 Na119 K100 Rb96 Cs90

Making Ions of Transition Elements  The electrons which are removed to make a positive ion are always removed from the orbitals with the highest principal quantum number (energy level) first! Example: Fe = [Ar]4s 2 3d 6 Fe +2 = [Ar]3d 6 Fe +3 = [Ar]3d 5

ELECTRON AFFINITY (E.A.)  The energy change when an electron is gained by an atom  The negative on these values means energy is released (it’s exothermic) when an electron is gained.  The more negative the electron affinity, the easier an atom will gain an electron.

ACROSS A SERIES  Electron affinity increases to a maximum in Group VII  It is a minimum in Group VIII because these elements already have stable electron structures  We see the same anomalies in this trend due to the filled & half filled sublevel stability

DOWN A GROUP  Electron affinity decreases slightly  Fluorine is an exception because it’s extra small size contributes to stronger electron-electron repulsion.

TRENDS - arrow points in the direction of the INCREASING trend I.E.E.A. → ↑ → ↑ I.E. and E.A. Summary Small IE is good for making Positive ions! (Metals) High EA is good for making Negative ions! (Non-Metals)

Metals vs. Non-metals  ACROSS A SERIES, elements become LESS metallic.  DOWN A GROUP, elements become MORE metallic. The big “stair-step” line separates metals from nonmetals. Elements to the left of the line are metals (except H!) Elements to the right of the line are nonmetals Some elements right along the line are metalloids.

Metal vs. Nonmetal

Test Note:  Be careful to distinguish trends from explanations!  If a question asks you to explain why an oxygen atom is smaller than a nitrogen atom, “because it’s farther right on the periodic table” is not the answer. That is a trend…the explanation is that there is a higher effective nuclear charge on oxygen which pulls the electrons closer to the nucleus.